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Chapter-1*********(BASIC CONCEPTS)***************
Short Question With Answer
Q.1 Calculate the grams atoms in 0.4 gm of potassium.
Ans.
Gram atoms of potassium =
= = 0.01 grams atoms
Q.2 23 grams of sodium and 238 gram of uranium have equal number of atoms in them.
Ans.
Mass of sodium = 23 gms= 1mole=6.02 x 1023 atoms
Mass of uranium = 238g=1 mole= 6.02x 1023 atoms
Both the substances have equal number of atoms because they have same no.of moles.
Q.3 Mg atom is twice heavier than that of carbon.
Ans.
The atomic mass of Mg is 24 which is to twice as mass as compared to the atomic mass of carbon i.e. 12. So Mg atom is twice heavier than that of carbon.
Q.4 180 grams of glucose and 342 gram of sucrose have the same number of molecules but different number of atoms present in them.
Ans.
180 grams of glucose (C6H12O6) and 342 grams of sucrose (C12H22O11) are their molar masses indicating one mole of each (glucose and sucrose) one mole of a substance contains equal number of molecules i.e. 6.02 x 1023.
Mass of glucose (C6H12O6)=180g=1mole=6.02x1023molecules
=24NAatoms
Mass of Sucrose (C12H22O11)=342g=1mole=6.02 x 1023 molecules
=45NAatoms
Q.5 4.9 g of H2SO4 when completely ionized in water have equal number of positive and negative ions, but the number of positively charged ions are twice the number of negatively charged ions.
Ans.
H2SO4 → 2H+ + SO4-2
When one mole of H2SO4 ionizes, it produces 2H+ and
SO4–2 ions. Hydrogen ions contains +1 charge while sulphate ions have – 2 charge. Hydrogen ions are twice in number than that of SO ion. Charges on both ions are equal (with opposite sign). Similarly ions produced by complete ionization of 4.9 grams of H2SO4 in water will have equal +ve and –ve charges but the number of H+ ions are twice than number of negatively charged sulphate ions.
Q.6 One mg of K2CrO4 has thrice the number of ions than the number of molecules when ionized in excess of water.
Ans. K2CrO4 → 2K+ + CrO4–2
When K2CrO4 ionizes in water, its one molecule gives three ions i.e. two K+ and one CrO4–2 (chromate) ions. The ratio between the number of molecules and number of ions than the number of molecules when ionized in water.
Q.7 Two grams of H2, 16g of CH4 and 44 gram of CO2 occupy separately the volumes of 22.414 dm3 at STP, although the sizes and masses of molecules of three gases are very different from each other.
Ans.
One mole of gas at STP occupies a volume of 22.4 dm3 sizes and masses of molecules of different gas do not affect the volume. Normally it is known that in the gaseous state, the distance between the molecules is 300 times greater that their diameter. Therefore two grams of H2, 16 grams of CH4 and 44 grams of CO2 (1 mole of each gas) separately occupy a volume of 22.4 dm3. This is called molar volume.
2gH2=1mole, 16gCH4=1 mole, 44gCO2=1 mole
1mole=22.414dm3
Q.8 Define Stoichiometry ?
Ans.
Stoichiometry is the branch of chemistry which gives a quantitative relationship between reactants and products in balanced chemical equation.
Q.9 What is limiting reactant? How does it control the quantity of the product formed? Explain with three examples. /Many chemical reactions taking place in our surroundings involve limiting reactants give examples?
Ans.
The reactant which controls (limits) the amount of product formed during a chemical reaction is called limiting reactant. In our surrounding many chemical reactions take place which involve limiting reactants some of these reactions are:
(i) Burning of coal to form CO2---Coal is limiting reactatnt C + O2 ® CO2
(ii) Burning of sui gas to form CO2 and H2O
CH4 + 2O2 ® CO2 + 2H2O
(iii) Rusting of iron----iron is limiting reactant
In above reactions oxygen is always in excess, while other reactants are consumed earlier. So other reactants are limiting reactants.
Q.10 One mole of H2O has two moles of bands, three moles of atoms, ten moles of electron and twenty–eight moles the total fundamental particles present in it.
Ans.
One molecule of H–O–H has two bounds between hydrogen and oxygen. There are three atoms i.e. two
H atoms and one O atom, therefore one mole of H2O contains two moles of bonds and three moles of atoms
(2 moles of H atoms and one mole of O atoms).
Similarly, there are eight elections in oxygen and one electron in each of the two, H atoms one molecule of H2O so has 10 electrons, so one mole of water contains 10 moles of electrons. There are 28 moles of all fundamental particles in one mole of water i.e.
10 moles of electrons.
10 moles of protons.
8 moles of neurons (8 neutrons in oxygen and there is no neutron in hydrogen) 28 moles of fundamental particles.
Q.11 One mole of H2SO4 should completely react with two moles of NaOH. How does Avogadro’s number help to explain it?
Ans.
The balanced chemical equation between H2SO4 and NaOH
H2SO4 + 2NaOH ® Na2SO4 + H2O
H2SO4® 2H+ + SO4-2
2NaOH®2Na+ + 2OH-
2H+ + 2OH- ® 2H2O
2NA 2NA
This is an acid base reaction, one mole of H2SO4 releases two moles of H+ ion in solution. It needs two moles of OH ions for complete neutralization. So two moles of NaOH which releases two moles of OH are required to react with one mole of H2SO4. One mole of H2SO4 releases twice the Avogadro’s number of H+ ions and it will need the Avogadro’s number of OH ions for complete neutralization.
Q.12 N2 and CO have same number of electrons, protons and neutrons.
Ans.
Both N2 and CO have same number of electrons, protons and neutrons as it is clear from the following explanation.
For N2 No. of electrons in N2 = 7 + 7 = 14
No. of protons in N2 = 7 + 7 = 14
No. of neutrons = 7 + 7 = 14
For CO number of electrons
in C = 6
No. of electrons in O = 8
Total no. of protons = 6 + 8 = 14
No. of neutrons in C = 6
No. of neutrons in O = 8
Total no. of neutrons = 6 + 8 = 14
Q.13 How many molecule, of water are in 12 gram of ice?
Ans.
Mass of ice (water) = 12.0 gm
Molar mass of water = 18 g/mol
No. of molecules of water
=
=
No. of molecules of water = 0.66 x 6.02 x 1023
= 3.97 x 1023
Q.14 Differentiate between limiting and non–limiting reactant ?
Ans. Limiting Reactant:
A limiting reactant is a reactant and that controls the amount of the product formed in a chemical reaction.
Non–Limiting Reactant:
The reactant which produces the excess amount of the product is called non–limiting reactant.
Q.15 Distinguish between actual yield and theoretical yield ?
Ans. Actual Yield:
The amount of the products obtained in a chemical reaction is called actual yield based on experiment.
Theoretical (Experiment) Yield:
The amount of the products calculated from the balanced chemical equation is called theoretical yield.
Q.16 What do you mean by percent yield? Give its significance ?
Ans.
The yield which is obtained by dividing actual yield with theoretical yield and multiplying by 100 is called percent yield.
% yield = x 100
Significance:
(i) % yield indicates the efficiency of reaction.
(ii) More is the percent yield higher will be the efficiency of reaction.
Q.17 Why actual yield is less than the theoretical yield?
Ans.
(a) Side reaction may takes place
(b) All the reactant may not be converted into products
(c) Mecahanical loss may occur like during
e.g Filtration, evaporation, crystallization, distillation etc.
Q.18 Calculate the mass of 10–3 moles of MgSO4.
Ans.
MgSO4 is an ionic compound. We will consider its formula mass instead of molecular mass.
Number of moles of substance
=
Formula mass of MgSO4 = 120 gm/ml
Number of moles of MgSO4 = 10–3 moles
Applying formula
10–3 =
Mass of MgSO4 = 120 x 10–3 = 0.12 moles
Q.19 Define Avogadro’s number ?
Ans.
Avogadro’s number is the number of atoms, molecules and ions in one gram atom of an element, one gram molecule of a compound and one gram ion of substance, respectively. It is equal to 6.02 x 1023.
Q.20 Define mole ?
Ans.
The molecular mass of a substance expressed in grams is called molecule or gram mole or simply the mole of a substance.
Moles of substance =
1 mole of water = 18.0 g
1 mole of H2SO4 = 98.0 g
Q.21 Define isotopes ?
Ans.
Atoms of the same element which have different masses but same atomic numbers are called isotopes. For example carbon has three isotopes.
12C6 13C6 14C6 and expressed as C–12, C–13 and C–14. Similarly hydrogen has three isotopes H H H called protium, deuterium and tritium.
Q.22 Define (i) ions (ii) Positive ion (iii) Negative ion.
Ans. Ion
As specie having positive or negative charges are called ions. For example Cl–1, NO, Na+, Ca++.
Positive Ion (Cation):
A specie has +ve charge is called positive ion and attracted towards Cathode . For example Na+, K++, Ca++.
Negative Ion (Anion)
A specie which has negative charge is called negative ion and attracted towards anode . For example F–1,Cl–1,Br–1andS–2P–3,C–4,SO, Cr2O, CO.
Q.23 Define and explain the molecular ion ?
Ans.
When a molecule loses or gains an electron, molecular ion is formed. For example CH4+, CO+, N2+. Cationic molecular, ions are more abundant than anionic ions.
The molecular ions find applications of in calculation of molecular mass of a compound. The molecular ions also help in the determination of structure of macro molecules.
The break down of molecular ions obtained from the natural products can give important information about their structure.
Q.24 What do understand by the relative atomic mass ?
Ans.
Relative atomic mass is the mass of an atom of element as compared to the mass of an atom of carbon taken as 12.
The unit used to express the relative atomic mass is called atomic mass unit (amu). It is th of the mass of one carbon atom. The relative atomic mass of 12C6 is 12.00 amu. The relative atomic mass of H is 1.0078 amu.
Q.25 Define Gram atom ?
Ans.
The atomic mass of an element expressed in grams is called gram atom of an element.
Number of gram atoms of a meter an element
=
For example 1 gram atom of hydrogen = 1.008 gm
1 gram atom of carbon = 12.00 gm
1 gram at of uranium = 238 gm
Q.26 Define gram ion ?
Ans.
The ionic mass of an ionic specie expressed in grams is called one gram ion or one mole of ions.
Number of gram ions =
1 gram ion of OH–1 = 17 grams
1 gram ion of SO = 96 gram
1 gram ion of CO = 60 gram
Q.27 Define gram formula and moles ?
Ans.
The formula mass of an ionic compound expressed in grams is called gram formula of the substance.
Number of gram formula or moles of a substance
=
1 gram formula of NaCl = 58.50 gms
1 gram formula of Na2CO3 = 106 gm
1 gram formula of AgNO3 = 170 gm
The atomic mass, molecular mass, formula mass or ionic mass of the substance expressed in grams is called moles of those substances.
Q.28 Define molar volume ?
Ans.
The volume occupied by one mole of an ideal gas at standard temperature and pressure (STP) is called molar volume. The volume is equal to 22.414 dm3.
Q.29 Define and explain atomicity ?
Ans.
The number of atoms present in a molecule is called the atomicity. The molecule can be monoatomic, diatomic and triatomic etc. If the molecule contains one atom it is monoatomic, if it contains two atoms it is diatomic, and if it contains three atoms it is triatomic. Molecules of elements may contain one two or more same type of atoms. For example He, Cl2, O3, P4, S8. The molecules of compounds consist of different kind of atoms. For example HCl, NH3, H2SO4, C6H12O6.
Q.30 Define an atom and molecule ?
Ans. Atom:
Atom is now defined as the smallest particle of an element, which may or may not have independent existence. For example He and Ne atoms have independent existence. While atoms of hydrogen, nitrogen and oxygen do not exist independently.
Molecule:
A molecule is the smallest particle of a pure substance( element or Compound ) which can exist independentally. For example N2, O2, Cl2, HCl, NH3 and H2SO4 are examples of molecules.
Q.31 What do you mean by empirical formula and molecular formula? How they are related to each other ?
Ans. Empirical Formula:
It is the simplest formula that gives information about the simple ratio of atoms present in a compound.
In an empirical formula of a compound Ax By, there are X atoms of an element A and y atoms of an element B.
Molecular Formula:
The formula of a substance which is based on the actual molecule is called molecular formula. It gives the usual number of atoms present in the molecule. For example molecular formula of benzene is C6H6, while that of glucose is C6H12O6. The molecular formula and empirical formula are related to each other by the following relationship.
Molecular formula = n x (Empirical formula)
Where “n” is simple integer.
Q.32 Is it true many compounds have same empirical and molecular formula ?
Ans.
There are many compounds, whose empirical formulas and molecular formulas are the same. For example H2O, CO2, NH3 and C12H22O11 have the same empirical and molecular formulas. Their simple multiple n is unity. Actually value of “n” is the ratio of molecular mass and empirical formula mass.
n =
Q.33 Ethylene glycol is used in automobile antifreeze. It has 38.7% carbon, 9.7% hydrogen and 51.6% oxygen. Its molar mass is 62 gms mole–1. Determine its empirical and molecular formula ?
Ans.
C = 38.7%, H = 9.7%, O = 51.6%
Dividing above %ages by atomic mass.
We get molar ratios
C = = 3.225
H = = 9.7
O = = 3.225
Dividing above molar ratio by least ratio we get atomic ratio.
C = = 1
H = = 3
O = = 1
Empirical formula is CH3O
Molar mass = 62
Empirical formula mass = 12 + 3 + 16 = 31
Now
n =
= = 2
Molar formula = n x Empirical formula
= 2 x CH3O
Molecular formula = C2H6O2
Hence molecular formula of Ethylene glycol = C2H6O2
Q.34 The combustion analysis of an organic compound shows it to contain 65.44% carbon 5.5% hydrogen and 29.06% of oxygen. What is the empirical formula of the compound? If the molecular mass of the compound is 110.15. Calculate the molecular formula of the compound.
Ans.
First of all divide the percentage of each element by its atomic mass to get the number of from atoms or moles.
No. of gram atoms of carbon =
= 5.45 gram atoms of C
No. of gram atoms of hydrogen =
= 5.45 gram atoms of H
No. of gram atoms of oxygen =
= 1.82 gram atoms of 0
Mole ratio C : H : O
4.45 5.45 1.82
Divide number of grams atoms by the smallest number
C : H : O
: :
3 : 3 : 1
Carbon, hydrogen and oxygen are present in the given organic compound in ratio of 3 : 3 : 1. So the empirical formula is C3H3O.
In order to calculate the molecular formula first calculate the empirical formula mass.
Empirical formula mass = 3 x 12 + 3 x 1 + 16
= 36 + 3 + 16 = 55.05
Molar mass of the compound = 110.15
h = = = 2
Molecular formula = n x empirical formula
= 2 x C3H3O
= C6H6O2
Q.35 Give relationships, between the amounts of substances and number of particles. There are three useful relationships ?
Ans.
1. Number of atoms of an element = x NA
2. Number of molecules of a compound
= x NA
3. Number of ions of ionic species = x NA
NA is the Avogadro’s number. The value is 6.02 x 1023.
Q.36 What are the types of relationships of stoichiometric calculations ?
Ans.
There are three types of relationships of stoichiometric calculations.
1. Mass–Mass Relationship
The relationship in which the mass of one substance is given and the mass of other substance is calculated.
2. Mass–mole or mole–mass relationship
The relationship in which mass of one substance is given and moles of other substance is to be calculated or vice versa.
3. Mass–volume or volume mass relationship
The relationship in which the mass of one substance is given and the volume of other substance is to be calculated or vice versa.
Q.37 Law of conservation of mass has to be obeyed during the stoichiometric calculations ?
Ans.
Stoichiometric calculations are based on balanced chemical equation and equation is balanced on the basis of Law of conservation of mass e.g
C+O2→ CO
In this equation stoichiometric calculations are not possible because it is not a balanced equation and it is not obeying Law of coseravtion.
Q.1 Calculate the grams atoms in 0.4 gm of potassium.
Ans.
Gram atoms of potassium =
= = 0.01 grams atoms
Q.2 23 grams of sodium and 238 gram of uranium have equal number of atoms in them.
Ans.
Mass of sodium = 23 gms= 1mole=6.02 x 1023 atoms
Mass of uranium = 238g=1 mole= 6.02x 1023 atoms
Both the substances have equal number of atoms because they have same no.of moles.
Q.3 Mg atom is twice heavier than that of carbon.
Ans.
The atomic mass of Mg is 24 which is to twice as mass as compared to the atomic mass of carbon i.e. 12. So Mg atom is twice heavier than that of carbon.
Q.4 180 grams of glucose and 342 gram of sucrose have the same number of molecules but different number of atoms present in them.
Ans.
180 grams of glucose (C6H12O6) and 342 grams of sucrose (C12H22O11) are their molar masses indicating one mole of each (glucose and sucrose) one mole of a substance contains equal number of molecules i.e. 6.02 x 1023.
Mass of glucose (C6H12O6)=180g=1mole=6.02x1023molecules
=24NAatoms
Mass of Sucrose (C12H22O11)=342g=1mole=6.02 x 1023 molecules
=45NAatoms
Q.5 4.9 g of H2SO4 when completely ionized in water have equal number of positive and negative ions, but the number of positively charged ions are twice the number of negatively charged ions.
Ans.
H2SO4 → 2H+ + SO4-2
When one mole of H2SO4 ionizes, it produces 2H+ and
SO4–2 ions. Hydrogen ions contains +1 charge while sulphate ions have – 2 charge. Hydrogen ions are twice in number than that of SO ion. Charges on both ions are equal (with opposite sign). Similarly ions produced by complete ionization of 4.9 grams of H2SO4 in water will have equal +ve and –ve charges but the number of H+ ions are twice than number of negatively charged sulphate ions.
Q.6 One mg of K2CrO4 has thrice the number of ions than the number of molecules when ionized in excess of water.
Ans. K2CrO4 → 2K+ + CrO4–2
When K2CrO4 ionizes in water, its one molecule gives three ions i.e. two K+ and one CrO4–2 (chromate) ions. The ratio between the number of molecules and number of ions than the number of molecules when ionized in water.
Q.7 Two grams of H2, 16g of CH4 and 44 gram of CO2 occupy separately the volumes of 22.414 dm3 at STP, although the sizes and masses of molecules of three gases are very different from each other.
Ans.
One mole of gas at STP occupies a volume of 22.4 dm3 sizes and masses of molecules of different gas do not affect the volume. Normally it is known that in the gaseous state, the distance between the molecules is 300 times greater that their diameter. Therefore two grams of H2, 16 grams of CH4 and 44 grams of CO2 (1 mole of each gas) separately occupy a volume of 22.4 dm3. This is called molar volume.
2gH2=1mole, 16gCH4=1 mole, 44gCO2=1 mole
1mole=22.414dm3
Q.8 Define Stoichiometry ?
Ans.
Stoichiometry is the branch of chemistry which gives a quantitative relationship between reactants and products in balanced chemical equation.
Q.9 What is limiting reactant? How does it control the quantity of the product formed? Explain with three examples. /Many chemical reactions taking place in our surroundings involve limiting reactants give examples?
Ans.
The reactant which controls (limits) the amount of product formed during a chemical reaction is called limiting reactant. In our surrounding many chemical reactions take place which involve limiting reactants some of these reactions are:
(i) Burning of coal to form CO2---Coal is limiting reactatnt C + O2 ® CO2
(ii) Burning of sui gas to form CO2 and H2O
CH4 + 2O2 ® CO2 + 2H2O
(iii) Rusting of iron----iron is limiting reactant
In above reactions oxygen is always in excess, while other reactants are consumed earlier. So other reactants are limiting reactants.
Q.10 One mole of H2O has two moles of bands, three moles of atoms, ten moles of electron and twenty–eight moles the total fundamental particles present in it.
Ans.
One molecule of H–O–H has two bounds between hydrogen and oxygen. There are three atoms i.e. two
H atoms and one O atom, therefore one mole of H2O contains two moles of bonds and three moles of atoms
(2 moles of H atoms and one mole of O atoms).
Similarly, there are eight elections in oxygen and one electron in each of the two, H atoms one molecule of H2O so has 10 electrons, so one mole of water contains 10 moles of electrons. There are 28 moles of all fundamental particles in one mole of water i.e.
10 moles of electrons.
10 moles of protons.
8 moles of neurons (8 neutrons in oxygen and there is no neutron in hydrogen) 28 moles of fundamental particles.
Q.11 One mole of H2SO4 should completely react with two moles of NaOH. How does Avogadro’s number help to explain it?
Ans.
The balanced chemical equation between H2SO4 and NaOH
H2SO4 + 2NaOH ® Na2SO4 + H2O
H2SO4® 2H+ + SO4-2
2NaOH®2Na+ + 2OH-
2H+ + 2OH- ® 2H2O
2NA 2NA
This is an acid base reaction, one mole of H2SO4 releases two moles of H+ ion in solution. It needs two moles of OH ions for complete neutralization. So two moles of NaOH which releases two moles of OH are required to react with one mole of H2SO4. One mole of H2SO4 releases twice the Avogadro’s number of H+ ions and it will need the Avogadro’s number of OH ions for complete neutralization.
Q.12 N2 and CO have same number of electrons, protons and neutrons.
Ans.
Both N2 and CO have same number of electrons, protons and neutrons as it is clear from the following explanation.
For N2 No. of electrons in N2 = 7 + 7 = 14
No. of protons in N2 = 7 + 7 = 14
No. of neutrons = 7 + 7 = 14
For CO number of electrons
in C = 6
No. of electrons in O = 8
Total no. of protons = 6 + 8 = 14
No. of neutrons in C = 6
No. of neutrons in O = 8
Total no. of neutrons = 6 + 8 = 14
Q.13 How many molecule, of water are in 12 gram of ice?
Ans.
Mass of ice (water) = 12.0 gm
Molar mass of water = 18 g/mol
No. of molecules of water
=
=
No. of molecules of water = 0.66 x 6.02 x 1023
= 3.97 x 1023
Q.14 Differentiate between limiting and non–limiting reactant ?
Ans. Limiting Reactant:
A limiting reactant is a reactant and that controls the amount of the product formed in a chemical reaction.
Non–Limiting Reactant:
The reactant which produces the excess amount of the product is called non–limiting reactant.
Q.15 Distinguish between actual yield and theoretical yield ?
Ans. Actual Yield:
The amount of the products obtained in a chemical reaction is called actual yield based on experiment.
Theoretical (Experiment) Yield:
The amount of the products calculated from the balanced chemical equation is called theoretical yield.
Q.16 What do you mean by percent yield? Give its significance ?
Ans.
The yield which is obtained by dividing actual yield with theoretical yield and multiplying by 100 is called percent yield.
% yield = x 100
Significance:
(i) % yield indicates the efficiency of reaction.
(ii) More is the percent yield higher will be the efficiency of reaction.
Q.17 Why actual yield is less than the theoretical yield?
Ans.
(a) Side reaction may takes place
(b) All the reactant may not be converted into products
(c) Mecahanical loss may occur like during
e.g Filtration, evaporation, crystallization, distillation etc.
Q.18 Calculate the mass of 10–3 moles of MgSO4.
Ans.
MgSO4 is an ionic compound. We will consider its formula mass instead of molecular mass.
Number of moles of substance
=
Formula mass of MgSO4 = 120 gm/ml
Number of moles of MgSO4 = 10–3 moles
Applying formula
10–3 =
Mass of MgSO4 = 120 x 10–3 = 0.12 moles
Q.19 Define Avogadro’s number ?
Ans.
Avogadro’s number is the number of atoms, molecules and ions in one gram atom of an element, one gram molecule of a compound and one gram ion of substance, respectively. It is equal to 6.02 x 1023.
Q.20 Define mole ?
Ans.
The molecular mass of a substance expressed in grams is called molecule or gram mole or simply the mole of a substance.
Moles of substance =
1 mole of water = 18.0 g
1 mole of H2SO4 = 98.0 g
Q.21 Define isotopes ?
Ans.
Atoms of the same element which have different masses but same atomic numbers are called isotopes. For example carbon has three isotopes.
12C6 13C6 14C6 and expressed as C–12, C–13 and C–14. Similarly hydrogen has three isotopes H H H called protium, deuterium and tritium.
Q.22 Define (i) ions (ii) Positive ion (iii) Negative ion.
Ans. Ion
As specie having positive or negative charges are called ions. For example Cl–1, NO, Na+, Ca++.
Positive Ion (Cation):
A specie has +ve charge is called positive ion and attracted towards Cathode . For example Na+, K++, Ca++.
Negative Ion (Anion)
A specie which has negative charge is called negative ion and attracted towards anode . For example F–1,Cl–1,Br–1andS–2P–3,C–4,SO, Cr2O, CO.
Q.23 Define and explain the molecular ion ?
Ans.
When a molecule loses or gains an electron, molecular ion is formed. For example CH4+, CO+, N2+. Cationic molecular, ions are more abundant than anionic ions.
The molecular ions find applications of in calculation of molecular mass of a compound. The molecular ions also help in the determination of structure of macro molecules.
The break down of molecular ions obtained from the natural products can give important information about their structure.
Q.24 What do understand by the relative atomic mass ?
Ans.
Relative atomic mass is the mass of an atom of element as compared to the mass of an atom of carbon taken as 12.
The unit used to express the relative atomic mass is called atomic mass unit (amu). It is th of the mass of one carbon atom. The relative atomic mass of 12C6 is 12.00 amu. The relative atomic mass of H is 1.0078 amu.
Q.25 Define Gram atom ?
Ans.
The atomic mass of an element expressed in grams is called gram atom of an element.
Number of gram atoms of a meter an element
=
For example 1 gram atom of hydrogen = 1.008 gm
1 gram atom of carbon = 12.00 gm
1 gram at of uranium = 238 gm
Q.26 Define gram ion ?
Ans.
The ionic mass of an ionic specie expressed in grams is called one gram ion or one mole of ions.
Number of gram ions =
1 gram ion of OH–1 = 17 grams
1 gram ion of SO = 96 gram
1 gram ion of CO = 60 gram
Q.27 Define gram formula and moles ?
Ans.
The formula mass of an ionic compound expressed in grams is called gram formula of the substance.
Number of gram formula or moles of a substance
=
1 gram formula of NaCl = 58.50 gms
1 gram formula of Na2CO3 = 106 gm
1 gram formula of AgNO3 = 170 gm
The atomic mass, molecular mass, formula mass or ionic mass of the substance expressed in grams is called moles of those substances.
Q.28 Define molar volume ?
Ans.
The volume occupied by one mole of an ideal gas at standard temperature and pressure (STP) is called molar volume. The volume is equal to 22.414 dm3.
Q.29 Define and explain atomicity ?
Ans.
The number of atoms present in a molecule is called the atomicity. The molecule can be monoatomic, diatomic and triatomic etc. If the molecule contains one atom it is monoatomic, if it contains two atoms it is diatomic, and if it contains three atoms it is triatomic. Molecules of elements may contain one two or more same type of atoms. For example He, Cl2, O3, P4, S8. The molecules of compounds consist of different kind of atoms. For example HCl, NH3, H2SO4, C6H12O6.
Q.30 Define an atom and molecule ?
Ans. Atom:
Atom is now defined as the smallest particle of an element, which may or may not have independent existence. For example He and Ne atoms have independent existence. While atoms of hydrogen, nitrogen and oxygen do not exist independently.
Molecule:
A molecule is the smallest particle of a pure substance( element or Compound ) which can exist independentally. For example N2, O2, Cl2, HCl, NH3 and H2SO4 are examples of molecules.
Q.31 What do you mean by empirical formula and molecular formula? How they are related to each other ?
Ans. Empirical Formula:
It is the simplest formula that gives information about the simple ratio of atoms present in a compound.
In an empirical formula of a compound Ax By, there are X atoms of an element A and y atoms of an element B.
Molecular Formula:
The formula of a substance which is based on the actual molecule is called molecular formula. It gives the usual number of atoms present in the molecule. For example molecular formula of benzene is C6H6, while that of glucose is C6H12O6. The molecular formula and empirical formula are related to each other by the following relationship.
Molecular formula = n x (Empirical formula)
Where “n” is simple integer.
Q.32 Is it true many compounds have same empirical and molecular formula ?
Ans.
There are many compounds, whose empirical formulas and molecular formulas are the same. For example H2O, CO2, NH3 and C12H22O11 have the same empirical and molecular formulas. Their simple multiple n is unity. Actually value of “n” is the ratio of molecular mass and empirical formula mass.
n =
Q.33 Ethylene glycol is used in automobile antifreeze. It has 38.7% carbon, 9.7% hydrogen and 51.6% oxygen. Its molar mass is 62 gms mole–1. Determine its empirical and molecular formula ?
Ans.
C = 38.7%, H = 9.7%, O = 51.6%
Dividing above %ages by atomic mass.
We get molar ratios
C = = 3.225
H = = 9.7
O = = 3.225
Dividing above molar ratio by least ratio we get atomic ratio.
C = = 1
H = = 3
O = = 1
Empirical formula is CH3O
Molar mass = 62
Empirical formula mass = 12 + 3 + 16 = 31
Now
n =
= = 2
Molar formula = n x Empirical formula
= 2 x CH3O
Molecular formula = C2H6O2
Hence molecular formula of Ethylene glycol = C2H6O2
Q.34 The combustion analysis of an organic compound shows it to contain 65.44% carbon 5.5% hydrogen and 29.06% of oxygen. What is the empirical formula of the compound? If the molecular mass of the compound is 110.15. Calculate the molecular formula of the compound.
Ans.
First of all divide the percentage of each element by its atomic mass to get the number of from atoms or moles.
No. of gram atoms of carbon =
= 5.45 gram atoms of C
No. of gram atoms of hydrogen =
= 5.45 gram atoms of H
No. of gram atoms of oxygen =
= 1.82 gram atoms of 0
Mole ratio C : H : O
4.45 5.45 1.82
Divide number of grams atoms by the smallest number
C : H : O
: :
3 : 3 : 1
Carbon, hydrogen and oxygen are present in the given organic compound in ratio of 3 : 3 : 1. So the empirical formula is C3H3O.
In order to calculate the molecular formula first calculate the empirical formula mass.
Empirical formula mass = 3 x 12 + 3 x 1 + 16
= 36 + 3 + 16 = 55.05
Molar mass of the compound = 110.15
h = = = 2
Molecular formula = n x empirical formula
= 2 x C3H3O
= C6H6O2
Q.35 Give relationships, between the amounts of substances and number of particles. There are three useful relationships ?
Ans.
1. Number of atoms of an element = x NA
2. Number of molecules of a compound
= x NA
3. Number of ions of ionic species = x NA
NA is the Avogadro’s number. The value is 6.02 x 1023.
Q.36 What are the types of relationships of stoichiometric calculations ?
Ans.
There are three types of relationships of stoichiometric calculations.
1. Mass–Mass Relationship
The relationship in which the mass of one substance is given and the mass of other substance is calculated.
2. Mass–mole or mole–mass relationship
The relationship in which mass of one substance is given and moles of other substance is to be calculated or vice versa.
3. Mass–volume or volume mass relationship
The relationship in which the mass of one substance is given and the volume of other substance is to be calculated or vice versa.
Q.37 Law of conservation of mass has to be obeyed during the stoichiometric calculations ?
Ans.
Stoichiometric calculations are based on balanced chemical equation and equation is balanced on the basis of Law of conservation of mass e.g
C+O2→ CO
In this equation stoichiometric calculations are not possible because it is not a balanced equation and it is not obeying Law of coseravtion.
CHAPTER#2 *Experimental Techniques in Chemistry******
SHORT QUESTIONS WITH ANSWERS
Q.1 Define analytical chemistry ?
Ans.
The branch of chemistry which deals with the qualitative and quantitative analyses of sample is called analytical chemistry.
Q.2 Define analysis and analyte ?
Ans. Analysis:
The determination of the composition of a sample is called analysis.
Analyte:
The sample being analyzed is called analyte.
Q.3 Differentiate between qualitative and quantitative analysis. What is analytical data ?
Ans. Qualitative analysis:
The analysis which deals with the detection or identification of the elements present in a compound is called qualitative analysis. It includes salt analysis and detection of functional groups.
Quantitative analysis:
The analysis in which the relative amounts of constituents are estimated is called quantitative analysis. For example combustion analysis.
Analytical data:
The results obtained by qualitative and quantitative analysis are called analytical data.
Q.4 Name the experimental techniques for purification of substances ?
Ans.
(1) Filtration.
(2) Crystalization.
(3) Sublimation.
(4) Solvent Extraction.
(5) Chromatography.
Q.5 Define filtration ?
Ans.
The process in which the insoluble particles are separated from the liquid by passing through several types of filter media is called filtration.
Q.6 Define the following.
(i) Filter (ii) Filtrate (iii) Residue
(iv) Filter Medium.
Ans. Filter:
Any water insoluble porous material having measurable degree of rigidity is called filtrate.
Filtrate:
The liquid which after passing the mixture through filter medium is collected is called filtrate.
Residue:
The solid left behind on the filter medium during filtration is called residue.
Filter medium:
The porous material used for filtration is called filter medium. It may be filter paper, cotton, woven wire, cloth etc.
Q.7 Give reasons that the funnel in filtration should be several inches long ?
Ans.
The stem of the funnel should be several inches long so that it can extend a few centimeters down into the receiving beaker and tip should touch the side of beaker in order to avoid splashing.
Q.8 Discus the points should be kept in mind during the folding of the filter paper ?
Ans.
1. Filter paper should be folded twice. The first fold should be along the diameter, and the second fold should be such that edges do not quite match.
2. Folded filter paper should be open slightly on the larger section. This provides a cone with three fold thickness half way and one thickness the other half way round.
3. The apex angle is greater than 60o.
4. The paper may be inserted in to 60 degree funnel moistened with water and firmly pressed down.
Q.9 Rate of filtration through funnel can be increased by using a fluted filter paper why ?
Ans.
A fluted filter paper is prepared by folding ordinary filter paper in such a way that fan like arrangement with alternate elevation and depression at various folds is formed.
Q.10 Differentiate b/w Gooch crucible and sintered glass crucible.
Ans.
(i) The filtering process can be done of the Gooch crucible is placed in suction filtering apparatus. But in filtration by sintered glass crucible no preparation is needed.
(ii) The reagents which react with paper e.g. HCl, KMnO4 etc. Cannot be filtered through Gooch crucible. But reactive solution like HCl, KMnO4 can be filtered with out any alteration in the sintered crucible.
(iii) Gooch Crucible for filtration of such types of precipitates which need to be heated at high temperature. But in sintered crucible while collecting the residue there is no contamination of filter paper.
Q .11 Define crystallization.
Ans.
The removal of a solid from solution by increasing its concentration above the saturation point in such a way that the excess solid separates out in the form of crystals is called crystallization.
Q.12 What is the principle of crystallization?
Ans.
The basic principle of crystallization is the fact that solute should be soluble in a suitable solvent at high temperature and the excess amount of the solute is thrown out as crystals when it is cooled.
Q.13 What is the objective of crystallization ?
Ans.
The preparation of chemical compound usually has a crude product and there is a need to purify it by crystallization from a suitable solvent
Q.14 Name the solvents commonly used for crystallization.
Ans.
The solvents which are mostly used for crystallization are, water, rectified spirit, (95% ethanol), absolute ethanol, diethylether acetone, chloroform, carbon tetrachloride, acetic acid and petroleum ether.
If none of the above solvents is found suitable for crystallization, a combination of two or more miscible solvents may be employed.
Q.15 Describe the preparation of saturated solution ?
Ans.
After selecting a suitable solvent the substance is then dissolved in a minimum amount of solvent and is heated directly or on a water bath with constant stirring. Add more solvent to the boiling solution, if necessary until all the solute has dissolved.
Q.16 What are the safe and reliable methods for drying the crystals ?
Ans.
1. Drying through vacuum desiccator
A safe and reliable method of drying crystals is through a vacuum desiccator. In this process the crystals are spread over water glass and kept in a vacuum desiccator for several hours. In the desiccator CaCl2, silica gel or P2O5 are used as drying agents
2. Drying in an oven
The crystals are dried in an oven provided the substance does not melt or decompose on heating at 1000c.
Q.17 How the undesirable colours are removed in crystals ?
Ans.
Sometimes during the preparation of crude substance, the colouring matter or resinous products affect the appearance of product and it may appear coloured. Such impurities are conveniently removed by boiling the substance in a solvent with the sufficient quantity of finely powdered animal charcoal and the pure decolourized substance crystallizes out from the filtrate on cooling.
Q.18 Define mother liquor? How the crystals can be obtained from mother liquor ?
Ans.
The remaining solution after the formation of crystals is called mother liquor.
1. The mixture of crystals and mother liquor is filtered through a Gooch crucible connected with a vacuum pump.
2. After full suction to drain the mother liquor as effectively as possible. When the filter is rigid enough it is pressed carefully but by firmly by means of a cork in order to drain the left over liquid.
3. The crystals are then washed will small portion of cold solvent repeating this process many times.
4. The crude mother liquor is concentrated by evaporation and it get good crops of crystals.
Q.19 Define sublimation.
Ans.
When substance is heated it goes directly in to vapours without passing through the liquid state and vapours thus formed are condensed back it form the solid on cooling once again with out passing though liquid state is called sublimation.
Examples of such solids are iodine, ammonium chloride, naphthalene , benzoic acid and camphor.
Q.20 What is sublimand ?
Ans.
The compound which is sublimed is called sublimand.e.g In the mixture of benzoic acid in sand, benzoic acid is sublimand.
Q.21 What is the main function and limitation of sublimation ?
Ans.
By this process of sublimation certain substance can be purified. It is only suitable for those substances which have high V.P than their melting point.
Q.23 What is solvent extraction? When it is applicable ?
Ans.
It is a technique, in which a solute can be separated from a solution by shaking the solution with a solvent in which the solute is more soluble and added solvent does not mix with the solution.
The technique of solvent extraction is mostly applied to separate organic compounds from water.
Q.24 What is the most common solvent in solvent extraction method? Why we choose ether in solvent extraction?
Ans.
The common solvent is ether in the solvent extraction we choose ether in the solvent extraction because ether layer is separated and organic product is obtained by evaporating ether repeated extractions using small portions of solvent ether are more efficient than using single but larger volume of solvent.
Q.25 State distribution law or partition law ?
Ans.
Distribution law:
This law states that a solute distribute itself between two immiscible liquids in a constant ratio irrespective of the amount of solute added.
The ration of the amounts of solute dissolved in two immiscible liquids at equilibrium is called distribution coefficient.
Distribution coefficient KD =
Q.26 Discuss the importance of solvent extraction method.
Ans.
Separation can be carried out on macro as well as micro level.
2. There is no need of any instrumentation except separating funnel.
3. It can be used for preparation, purification and analysis on all scales of working.
Q.27 Define chromatography? What is the principle of chromatography?
Ans.
Chromatography is a technique used for separating the components of a mixture. These components are separated due to the relative affinity for stationary phase and mobile phase.
Q.28 Define and explain, stationary phase and mobile phase.
Ans.
Stationary phase:
The phase over which mobile phase flows in chromatography is called stationary phase.
The stationary phase may be a solid or liquid supported on a solid. It adsorbs the mixture under separation.
Examples of stationary phase are silica gel, alumina and filter paper etc.
Mobile phase:
The solvent or mixture of solvents for this separation of components is called mobile phase.
The mobile phase may be liquid or gas and while passing one the stationary phase, competes with it for the constituents of mixture.
Examples of mobile phase are, water, ethanol, ethanoic acid and propanone (acetone) etc.
Q.29 What is the principle of chromatography ?
Ans.
The principle involved in the chromatography depends upon the relative solubilities of the components, between the two phases. The distribution of the components mixture between the two phases is governed by the distribution coefficient KD, which is ratio of component in mobile phase to the concentration of component in stationary phase.
KD =
Q.30 What is the importance of distribution coefficient ?
Ans.
(i) The component of a mixture with a small value of KD mostly remains in the stationary phase as moving phase flows over it.
The component with a greater value of KD remains largely dissolved in the mobile phase and passes over the stationary phase quickly.
Q.31 Differentiate between adsorption chromatography and partition chromatography ?
Ans. Adsorption chromatography:
Type of Chromatography in which the stationary phase is solid, is called adsorption chromatography. Example of this chromatography is Thin layer chromatography.
Q.32 What is partition chromatography ?
Ans. Partition chromatography:
Type of Chromatography in which the stationary phase is liquid is called partition chromatography. Example of this chromatography is paper chromatography.
Q.33 Define and explain paper chromatography ?
Ans.
It is a technique of partition chromatography in which the stationary phase is water adsorbed on a paper. The mobile phase is usually an organic liquid.
In paper chromatography the adsorbed water behaves as an immiscible liquid towards the mobile phase, which passes over the paper.
Q.34 Name the different ways of paper chromatography.
Ans.
There are three ways of carrying out paper chromatography.
(i) Ascending chromatography
(ii) Descending chromatography
(iii) Radial/Circular chromatography.
Q.35 What do you mean by Rf value?
Ans.
Each component has specific retardation factor called Rf value. Rf value is related to distribution coefficient and is given by
Rf =
Q.37 What is chromatogram ?
Ans. Finished or Devolped paper obtained after Chromatography is called Chromatogram:
Q.38 Name the types of chromatography on the basis of phase.
Ans.
There are four types of chromatography.
(i) Liquid–liquid chromatography
(ii) Liquid–solid chromatography
(iii) Gas–liquid chromatography
(iv) Gas–solid chromatography
Q.39 Give uses of chromatography.
Ans.
The techniques of chromatography is very useful in organic synthesis for separation, isolation, and purification of the products.
Mostly used for the separation of amino acids.
Q.40 Why is there a need to crystallize the crude products ?
Ans.
When a chemical compound is synthesised, it is crude product. Therefore, there is need to purify the compound. This is done by crystallizing the compound.
Q.41 A water soluble organic compound aspirin is prepared by the reaction of salicylic acid with a mixture of acetic acid and acetic anhydride. How will you separate the product from the reaction mixture ?
Ans.
During the preparation of aspirin, it is obtained as only liquid which can be separated by solvent extraction technique using a non–polar solvent like CCl4 and mixture is transferred to separating funnel where only layer is separated.
Q.42 A solid compound is soluble in water as well as in chloroform. During its preparation it remains in aqueous layer. Describe a method to obtain it from the layer.
Ans.
The compound can be extracted by solvent extraction technique. As it is mentioned that compound is soluble in polar solvents (water, chloroform). If a non–polar solvent is mixed and the mixture is transferred to a separating funnel, where two layers are formed. By separating water layer and evaporating it, organic compound is obtained.
Q.43 why repeated extractions using small portions of solvent are more efficient than using a single but larger volume of solvent?
Ans.
It has been observed that repeated extractions using small portions of solvent are more efficient than using a single but larger volume of solvent. Because more product is extracted with more extractions using small portions of solvent.
Q.44 Write down the main characteristics of a solvent selected for the crystallization of compound.
Ans.
A solvent should have the following characteristics
1. It should dissolve a large amount of solute in its boiling part.
2. It should have not reaction with the solute.
3. It should neither dissolve the impurities, nor crystallize them with the solute.
4. It should be perfectly safe to use.
5. It should be easily removable.
6. It should be inexpensive.
Q.1 Define analytical chemistry ?
Ans.
The branch of chemistry which deals with the qualitative and quantitative analyses of sample is called analytical chemistry.
Q.2 Define analysis and analyte ?
Ans. Analysis:
The determination of the composition of a sample is called analysis.
Analyte:
The sample being analyzed is called analyte.
Q.3 Differentiate between qualitative and quantitative analysis. What is analytical data ?
Ans. Qualitative analysis:
The analysis which deals with the detection or identification of the elements present in a compound is called qualitative analysis. It includes salt analysis and detection of functional groups.
Quantitative analysis:
The analysis in which the relative amounts of constituents are estimated is called quantitative analysis. For example combustion analysis.
Analytical data:
The results obtained by qualitative and quantitative analysis are called analytical data.
Q.4 Name the experimental techniques for purification of substances ?
Ans.
(1) Filtration.
(2) Crystalization.
(3) Sublimation.
(4) Solvent Extraction.
(5) Chromatography.
Q.5 Define filtration ?
Ans.
The process in which the insoluble particles are separated from the liquid by passing through several types of filter media is called filtration.
Q.6 Define the following.
(i) Filter (ii) Filtrate (iii) Residue
(iv) Filter Medium.
Ans. Filter:
Any water insoluble porous material having measurable degree of rigidity is called filtrate.
Filtrate:
The liquid which after passing the mixture through filter medium is collected is called filtrate.
Residue:
The solid left behind on the filter medium during filtration is called residue.
Filter medium:
The porous material used for filtration is called filter medium. It may be filter paper, cotton, woven wire, cloth etc.
Q.7 Give reasons that the funnel in filtration should be several inches long ?
Ans.
The stem of the funnel should be several inches long so that it can extend a few centimeters down into the receiving beaker and tip should touch the side of beaker in order to avoid splashing.
Q.8 Discus the points should be kept in mind during the folding of the filter paper ?
Ans.
1. Filter paper should be folded twice. The first fold should be along the diameter, and the second fold should be such that edges do not quite match.
2. Folded filter paper should be open slightly on the larger section. This provides a cone with three fold thickness half way and one thickness the other half way round.
3. The apex angle is greater than 60o.
4. The paper may be inserted in to 60 degree funnel moistened with water and firmly pressed down.
Q.9 Rate of filtration through funnel can be increased by using a fluted filter paper why ?
Ans.
A fluted filter paper is prepared by folding ordinary filter paper in such a way that fan like arrangement with alternate elevation and depression at various folds is formed.
Q.10 Differentiate b/w Gooch crucible and sintered glass crucible.
Ans.
(i) The filtering process can be done of the Gooch crucible is placed in suction filtering apparatus. But in filtration by sintered glass crucible no preparation is needed.
(ii) The reagents which react with paper e.g. HCl, KMnO4 etc. Cannot be filtered through Gooch crucible. But reactive solution like HCl, KMnO4 can be filtered with out any alteration in the sintered crucible.
(iii) Gooch Crucible for filtration of such types of precipitates which need to be heated at high temperature. But in sintered crucible while collecting the residue there is no contamination of filter paper.
Q .11 Define crystallization.
Ans.
The removal of a solid from solution by increasing its concentration above the saturation point in such a way that the excess solid separates out in the form of crystals is called crystallization.
Q.12 What is the principle of crystallization?
Ans.
The basic principle of crystallization is the fact that solute should be soluble in a suitable solvent at high temperature and the excess amount of the solute is thrown out as crystals when it is cooled.
Q.13 What is the objective of crystallization ?
Ans.
The preparation of chemical compound usually has a crude product and there is a need to purify it by crystallization from a suitable solvent
Q.14 Name the solvents commonly used for crystallization.
Ans.
The solvents which are mostly used for crystallization are, water, rectified spirit, (95% ethanol), absolute ethanol, diethylether acetone, chloroform, carbon tetrachloride, acetic acid and petroleum ether.
If none of the above solvents is found suitable for crystallization, a combination of two or more miscible solvents may be employed.
Q.15 Describe the preparation of saturated solution ?
Ans.
After selecting a suitable solvent the substance is then dissolved in a minimum amount of solvent and is heated directly or on a water bath with constant stirring. Add more solvent to the boiling solution, if necessary until all the solute has dissolved.
Q.16 What are the safe and reliable methods for drying the crystals ?
Ans.
1. Drying through vacuum desiccator
A safe and reliable method of drying crystals is through a vacuum desiccator. In this process the crystals are spread over water glass and kept in a vacuum desiccator for several hours. In the desiccator CaCl2, silica gel or P2O5 are used as drying agents
2. Drying in an oven
The crystals are dried in an oven provided the substance does not melt or decompose on heating at 1000c.
Q.17 How the undesirable colours are removed in crystals ?
Ans.
Sometimes during the preparation of crude substance, the colouring matter or resinous products affect the appearance of product and it may appear coloured. Such impurities are conveniently removed by boiling the substance in a solvent with the sufficient quantity of finely powdered animal charcoal and the pure decolourized substance crystallizes out from the filtrate on cooling.
Q.18 Define mother liquor? How the crystals can be obtained from mother liquor ?
Ans.
The remaining solution after the formation of crystals is called mother liquor.
1. The mixture of crystals and mother liquor is filtered through a Gooch crucible connected with a vacuum pump.
2. After full suction to drain the mother liquor as effectively as possible. When the filter is rigid enough it is pressed carefully but by firmly by means of a cork in order to drain the left over liquid.
3. The crystals are then washed will small portion of cold solvent repeating this process many times.
4. The crude mother liquor is concentrated by evaporation and it get good crops of crystals.
Q.19 Define sublimation.
Ans.
When substance is heated it goes directly in to vapours without passing through the liquid state and vapours thus formed are condensed back it form the solid on cooling once again with out passing though liquid state is called sublimation.
Examples of such solids are iodine, ammonium chloride, naphthalene , benzoic acid and camphor.
Q.20 What is sublimand ?
Ans.
The compound which is sublimed is called sublimand.e.g In the mixture of benzoic acid in sand, benzoic acid is sublimand.
Q.21 What is the main function and limitation of sublimation ?
Ans.
By this process of sublimation certain substance can be purified. It is only suitable for those substances which have high V.P than their melting point.
Q.23 What is solvent extraction? When it is applicable ?
Ans.
It is a technique, in which a solute can be separated from a solution by shaking the solution with a solvent in which the solute is more soluble and added solvent does not mix with the solution.
The technique of solvent extraction is mostly applied to separate organic compounds from water.
Q.24 What is the most common solvent in solvent extraction method? Why we choose ether in solvent extraction?
Ans.
The common solvent is ether in the solvent extraction we choose ether in the solvent extraction because ether layer is separated and organic product is obtained by evaporating ether repeated extractions using small portions of solvent ether are more efficient than using single but larger volume of solvent.
Q.25 State distribution law or partition law ?
Ans.
Distribution law:
This law states that a solute distribute itself between two immiscible liquids in a constant ratio irrespective of the amount of solute added.
The ration of the amounts of solute dissolved in two immiscible liquids at equilibrium is called distribution coefficient.
Distribution coefficient KD =
Q.26 Discuss the importance of solvent extraction method.
Ans.
Separation can be carried out on macro as well as micro level.
2. There is no need of any instrumentation except separating funnel.
3. It can be used for preparation, purification and analysis on all scales of working.
Q.27 Define chromatography? What is the principle of chromatography?
Ans.
Chromatography is a technique used for separating the components of a mixture. These components are separated due to the relative affinity for stationary phase and mobile phase.
Q.28 Define and explain, stationary phase and mobile phase.
Ans.
Stationary phase:
The phase over which mobile phase flows in chromatography is called stationary phase.
The stationary phase may be a solid or liquid supported on a solid. It adsorbs the mixture under separation.
Examples of stationary phase are silica gel, alumina and filter paper etc.
Mobile phase:
The solvent or mixture of solvents for this separation of components is called mobile phase.
The mobile phase may be liquid or gas and while passing one the stationary phase, competes with it for the constituents of mixture.
Examples of mobile phase are, water, ethanol, ethanoic acid and propanone (acetone) etc.
Q.29 What is the principle of chromatography ?
Ans.
The principle involved in the chromatography depends upon the relative solubilities of the components, between the two phases. The distribution of the components mixture between the two phases is governed by the distribution coefficient KD, which is ratio of component in mobile phase to the concentration of component in stationary phase.
KD =
Q.30 What is the importance of distribution coefficient ?
Ans.
(i) The component of a mixture with a small value of KD mostly remains in the stationary phase as moving phase flows over it.
The component with a greater value of KD remains largely dissolved in the mobile phase and passes over the stationary phase quickly.
Q.31 Differentiate between adsorption chromatography and partition chromatography ?
Ans. Adsorption chromatography:
Type of Chromatography in which the stationary phase is solid, is called adsorption chromatography. Example of this chromatography is Thin layer chromatography.
Q.32 What is partition chromatography ?
Ans. Partition chromatography:
Type of Chromatography in which the stationary phase is liquid is called partition chromatography. Example of this chromatography is paper chromatography.
Q.33 Define and explain paper chromatography ?
Ans.
It is a technique of partition chromatography in which the stationary phase is water adsorbed on a paper. The mobile phase is usually an organic liquid.
In paper chromatography the adsorbed water behaves as an immiscible liquid towards the mobile phase, which passes over the paper.
Q.34 Name the different ways of paper chromatography.
Ans.
There are three ways of carrying out paper chromatography.
(i) Ascending chromatography
(ii) Descending chromatography
(iii) Radial/Circular chromatography.
Q.35 What do you mean by Rf value?
Ans.
Each component has specific retardation factor called Rf value. Rf value is related to distribution coefficient and is given by
Rf =
Q.37 What is chromatogram ?
Ans. Finished or Devolped paper obtained after Chromatography is called Chromatogram:
Q.38 Name the types of chromatography on the basis of phase.
Ans.
There are four types of chromatography.
(i) Liquid–liquid chromatography
(ii) Liquid–solid chromatography
(iii) Gas–liquid chromatography
(iv) Gas–solid chromatography
Q.39 Give uses of chromatography.
Ans.
The techniques of chromatography is very useful in organic synthesis for separation, isolation, and purification of the products.
Mostly used for the separation of amino acids.
Q.40 Why is there a need to crystallize the crude products ?
Ans.
When a chemical compound is synthesised, it is crude product. Therefore, there is need to purify the compound. This is done by crystallizing the compound.
Q.41 A water soluble organic compound aspirin is prepared by the reaction of salicylic acid with a mixture of acetic acid and acetic anhydride. How will you separate the product from the reaction mixture ?
Ans.
During the preparation of aspirin, it is obtained as only liquid which can be separated by solvent extraction technique using a non–polar solvent like CCl4 and mixture is transferred to separating funnel where only layer is separated.
Q.42 A solid compound is soluble in water as well as in chloroform. During its preparation it remains in aqueous layer. Describe a method to obtain it from the layer.
Ans.
The compound can be extracted by solvent extraction technique. As it is mentioned that compound is soluble in polar solvents (water, chloroform). If a non–polar solvent is mixed and the mixture is transferred to a separating funnel, where two layers are formed. By separating water layer and evaporating it, organic compound is obtained.
Q.43 why repeated extractions using small portions of solvent are more efficient than using a single but larger volume of solvent?
Ans.
It has been observed that repeated extractions using small portions of solvent are more efficient than using a single but larger volume of solvent. Because more product is extracted with more extractions using small portions of solvent.
Q.44 Write down the main characteristics of a solvent selected for the crystallization of compound.
Ans.
A solvent should have the following characteristics
1. It should dissolve a large amount of solute in its boiling part.
2. It should have not reaction with the solute.
3. It should neither dissolve the impurities, nor crystallize them with the solute.
4. It should be perfectly safe to use.
5. It should be easily removable.
6. It should be inexpensive.
Chap 4 ****LIQUIDS AND SOLIDS**
SHORT QUESTION WITH ANSWERS Q.1 what is difference between (i) Intermolecular forces and intermolecular forces (ii) Polar molecules and non–polar molecules (iii) Induce and dipole and instantaneous dipole (iv) Dipole Ans. (i) Intermolecular Forces: The forces of attraction between two different atoms ions and molecules are called intermolecular forces. For example H—Cl……H--Cl…….H--Cl Intermolecular Forces: The forces of attraction between two atoms or group of atoms present with in the same molecule, are called intermolecular forces.e.g covalent bond ,ionic bond etc (ii) Polar molecules: A molecule which has partial +ve and partial –ve charges on it due to difference of electro negativity between bonded atoms is called polar molecules. For example H8+ – Cl8– Non–polar molecules: A molecule in which bonded atoms have zero or negligible electro negativity difference is called non–polar molecules. For example H | Cl–Cl, H–H, H–C–H | H Important point to remember: All molecules having same atoms(Homoatomic) are non polar (iii) Induced Dipole: A molecule in which polarity is created due to other polar molecule is called induced dipole. Instantaneous Dipole: The temporary dipole (polarity) produced in a non–polar molecule at a certain instant is called instantaneous dipole. (iv) Dipole A molecule which has two poles i.e. two charges partial +ve and partial –ve is known asdipole.e.g H8+ – Cl8– Q.2 Define intermolecular forces, and the types of intermolecular forces? Ans. Intermolecular Forces: The forces of attraction that exist between all kinds of atoms, molecules, when they are sufficiently close to each other are called intermolecular forces. Types of intermolecular forces: There are four types of intermolecular forces. (a) Dipole–dipole forces (b) Ion–dipole forces (c) Dipole–induced dipole forces (Debye forces) (d) Instantaneous dipole–induced dipole forces or (London dispersion forces). (a) Dipole–dipole forces: The forces of attraction between the positive end of one polar molecule and the negative end of other polar molecule are known as dipole–dipole forces. Example is of HCl. H8+ ® Cl8– ® H8+ ® Cl8– ® H8+ ® Cl8– (b) Ion–Dipole forces: The forces of attraction in which the negative ends of polar molecules are attracted towards the cation (+ve ion) and positive ends towards anion (– ion) are called ion–dipole forces. Ionic compounds like Mx are normally soluble in polar solvent like water. Water molecules break the crystal lattice and the ions are set free. These positive and negative ions are then surrounded by water molecules. The negative ends of the dipole of the water are attracted towards the cation (M+) while the positive ends are attracted towards the anion (X–). The dissolution of most of the ionic compounds in water is due to this reason. The forces of attraction between ions and water molecules are known as Ion–dipole forces. (c) The forces of attraction that exist between already polar molecules and the molecule having induced dipole forces. The forces are also called Debye forces. (d) Instantaneous Dipole–Induced Dipole forces. (London dispersion forces). The momentary forces of attraction that exist between instantaneous dipole and the induced dipole are called instantaneous dipole–induced dipole forces. The momentary force of attraction between instantaneous dipole and the induced dipole is known as instantaneous–induced dipole forces. Q.3 Explain the factors affecting the London forces. Ans. The strength of these forces depends upon the following two factors. 1. Size of electronic cloud: As the size of electronic cloud of atoms or molecules increases, dispersion becomes easier and these forces are more permanent. The elements of zero group are monoatomic gases due to their complete outermost shells, they do not form covalent bonds. Their boiling points increase from top to bottom in a group. 2. Polarizability: The quantitative measurement of the extent to which the electronic cloud can be polarized or distortedis called polarizability. The boiling points of halogens increase from top to bottom i.e. from fluorine to iodine. 3. Number of atoms: As the number of atoms in non–polar molecule increases polarizability of the molecule increases and London forces become stronger. The boiling points of saturated hydrocarbons increase as the number of atoms increases. Q.4 Define and explain hydrogen bonding./ What is the origin of intermolecular forces in water? Ans. Hydrogen bonding: “The electrostatic force of attraction between electronegative atom and partial positive hydrogen atom is called hydrogen bonding.” Explanation: Consider water molecules to understand hydrogen bonding oxygen is more electronegative than hydrogen. So water is polar molecule. There will be dipole–dipole forces of attraction between water molecules. The electrostatic force of attraction between electronegative oxygen of one molecule and partial positive hydrogen of other molecule is called hydrogen bonding. Strength of H–Bonding: Hydrogen bonding is stronger than simple dipole–dipole forces. This is due to the following reasons. 1. There are two lone pairs on oxygen atom. Oxygen forms coordinate covalent bond with hydrogen. 2. There is sufficient partial positive charge on hydrogen. Both positively charged hydrogen of water molecules produce strong electric field due to their small size. Hydrogen bonding in water molecules acts like a bridge between two electronegative oxygen atoms. Generally, the strength of H–bonding is twenty times less than that of covalent bond. Q.5 Give the properties of compounds containing hydrogen bonding. Ans. 1. There are dynamic properties of covalent compounds. 2. Solubility of Hydrogen bonded molecules. 3. Cleansing action. 4. Application of hydrogen bonding in biological compounds. 5. Surface tension. 6. Effect of hydrogen bonding on viscosity. 7. Hydrogen bonding in paints and dyes. 8. Clothing. 9. Food material. 10. Structure of ice. Q.6 Explain the following with reasons. (a) In the hydrogen bonded structure of H–F, which is stronger bond, the shorter covalent bond or the longer hydrogen between different molecules? (b) In a very cold winter fish in garden ponds owe their lives to hydrogen bonding. (c) Water and ethanol can mix easily and in all proportions. Ans. (a) There is sufficient hydrogen bonding in H–F molecules and it gives zig zag structure. Fluorine atom is present at the end while H atoms are entrapped between two strong electronegative atoms. The covalent bond between H and F is stronger because it is produced by the overlapping of orbital’s and two electrons have been shared to give sigma bond. The bond which is shown by the dotted line is the hydrogen bond due to electrostatic forces of attraction so, it is a weaker bond. (b) When water is frozen at 0oC, then it expands. This is due to the fact that due to H-bonding in ice the molecules become arranged density of ice is decreased. That’s why ice floats on water. (c) Water (H-OH) and ethanol (C2H5OH) have both are polar solvents and having OH groups. So, they can do the hydrogen bonding extensively. That is they can mix with each other in all proportions. Q.7 Why H2S is a gas while H2O is liquid at room temperature? Ans: This is due to high electro negativity of oxygen as compared to sulphur. Water has hydrogen bonding, but H2S does not have. Due to absence of hydrogen bonding in H2S at room temperature, it is a gas. Q.8 Earthen ware vessels keep water cool? Ans: Earthen were vessels are porous- water molecules come out from these pores and evaporate. Heat of the atmosphere can not enter into the liquid. So temperature of the liquid in earthenware's remains less. Q.9 one feels sense of cooling under the fan after bath? Ans: When one takes bath and sits in front of a fan, water on the surface of body evaporates with greater rate.The high energy molecules escape from surface of the body and one feels sense of cooling. Q.10 Why the heat of vapourization of water is greater than that of CH4? Ans: Water is a polar liquid and due to strong hydrogen bonding high energy is required to separate the molecules from each other at its boiling point. CH4 is a non-polar and has weak London dispersion forces. Q.11 Define and explain evaporation is a cooling process. Give reason. Ans. Evaporation: The spontaneous change of liquid into its vapours is called evaporation. It continues at all temperature. Evaporation increases with the increase of temperature. Explanation: The molecules of liquid are not motionless. The energy of the molecules is not equally distributed.The molecules which have low kinetic energy move slowly while others with high kinetic energy move faster. If one of the higher speed molecules reaches the surface, it may escape the attractions of its neighboring molecules and leaves the bulk of the liquid. This spontaneous change of liquid into its vapours is called evaporation Evaporation causes cooling: The reason is that when high energy molecules leave the liquid and low energy molecules are left behind, the temperature of the liquid falls and heat moves from the surrounding to the liquid and the temperature of the surrounding also falls. So evaporation is a cooling process. Q.12 H-bonding is present in chloroform and acetone-justify it? Ans: Chloroform is a polar compound. Acetone is also a polar compound. When chloroform and acetone are mixed with each other, than they create the forces of attractions due to hydrogen bonding. Q.13 Evaporation of a liquid takes place at all temperatures give reason? Ans: Evaporation takes place due to the K.E of the molecules since the K.E of the molecules can not be zero at any temperature therefore evaporation takes place at all temperatures. Q.14 What are the factors that affect the rate of evaporation? Ans. 1. Surface Area: Evaporation takes place from liquid surface. If area of the surface of liquid increases the rate of evaporation will also increase. 2. Temperature: Temperature also affects rate of evaporation Higher the temperature faster will be the rate of evaporation. 3. Intermolecular forces: Stronger the intermolecular attractive forces slower is the value of evaporation and vice versa. Q.15 Define and explain the vapour pressure. Ans. Vapour Pressure: The pressure exerted by the vapours on the surface of liquid at equilibrium state at a given temperature is called vapour pressure. Explanation: Consider a liquid closed in container at a certain temperature. High energy molecules leave the surface of liquid and gather above the surface in the empty space in the form of vapours. These molecules collide with the walls of container as well as with the surface of liquid. In this way they lose some their kinetic energy and there is a chance that these molecules are recaptured by the liquid surface. This process is known as condensation. Both the process i.e. condensation and evaporation continue, till rates of both processes become equal. This state is called dynamic equilibrium, and the pressure exerted by the vapours at this state on the liquid surface at particular temperature is called vapour pressure. Vapour pressure does not depend upon amount or volume of liquid and surface area. Q.16 What are the factors affecting vapour pressure. Ans. 1. Nature of liquid 2. Strength of intermolecular forces 3. Size of molecules 4. Temperature. Q.17 Define boiling point. Ans. The temperature at which the vapour pressure of liquid becomes equal to the external atmospheric pressure is called boiling point of liquid. Q.18 Give variation of vapour pressure and boiling point. Ans. Vapour pressure is closely related to boiling point. Variation in vapour pressure depends upon the following factors. 1. Temperature: vapour pressure of a liquid increases by increasing temperature. Higher the temperature more will be the vapour pressure and vice versa. Liquids boil at that temperature when their vapour pressures are equal to 760 torr at sea level. By increasing external pressure boiling point can be increased. 2. Strength of intermolecular forces: Stronger the intermolecular forces lower will be vapour pressure and higher will be the boiling point. Q.19 What is the effect of external pressure on boiling point? Ans. A liquid boils when its internal pressure becomes equal to external atmospheric pressure so, by changing external pressure, a liquid can be boiled at any temperature. If external pressure is greater, the liquid needs more heat to equalize the internal pressure to external atmospheric pressure. Similarly if external pressure is lower, liquid needs less amount of heat to equalize its vapour pressure, the external pressure. under 700 torr (at Murree hills) water boils at 98oC. Q.20 Why boiling point of water is 980C at Murree? Ans: At high altitudes the atmospheric pressure becomes low therefore B.P of water at Murree is 980C. Q.21 Why boiling point of water is 1200C at 1489 torr why? Ans: The normal B.P of H2O is 1000C at 760 torr since B.P increases by increasing pressure therefore B.P of H2O is 1200C at 1489 torr. Q.22 Why the boiling points of the hydrides of second period in group IV-A,V-A,VI-A and VII-A are greater than the B.P of hydrides of third period? Ans: The elements of second period are more electronegative than the respective element third period. So,the polarities of the bonds with hydrogen are greater than the third period elements. H2 O > H2 S; NH3 > P H3 ;HF > HCI; CH4 < SiH4 Q.23 Define molar heat of vapourization? Ans: The amount of heat required to vapourize one mole of liquid at its boiling point is called molar heat of vapourization. Q.24 What is vacuum distillation? Explain. Ans. Definition: The process in which liquid is heated under reduced pressure, to convert it into its vapours at low temperature and then to condense these vapours into liquid is known as vacuum distillation. Explanation: In vacuum distillation boiling point of liquid decreases by reducing the pressure. This is done by connecting the distillation apparatus to the vacuum pump. In this way liquids with high boiling points can be boiled at low temperature. Q.25 Define enthalpy change. Ans. If physical or chemical change occurs at constant pressure then it is known as enthalpy change. Q.26 What are types of enthalpy changes? Ans. There are three types of enthalpy changes. 1. Molar Heat of Fusion (DHf): The amount of heat absorbed by one mole of a solid to melt it into liquids at its melting point at atmospheric pressure is called molar heat of fusion. It is denoted by D Hf. 2. Molar Heat of vapourization (D Hv): The amount of heat absorbed by one mole of a liquid to convert it into one mole of vapours at its boiling point at 1 atmospheric pressure is called molar heat of vapourization. It is denoted by DHv. 3. Molar Heat of sublimation (D Hs): The amount of heat absorbed by one mole of a solid to convert it directly into one mole of its vapours at particular temperature at 1 atmospheric pressure is called molar heat of sublimation. It is denoted by D Hs. Q.27 What are liquid crystals? Give their types. Ans. The molecules which are large somewhat rigid and linear having some of structures of solids showing optical properties and some of the freedom of motion of liquids are called liquid crystals. Types of liquid crystals: (a) Smectic liquid crystals. (b) Nematic liquid crystals. (c) Cholesteric liquid crystals. Q.28 What are solids? Ans. Solids are those substances which are rigid, hard, have definite shape and definite volume. The atoms, ions, and molecules, that make up a solid are close packed. They are held together by strong cohesive forces. Q.29 Crystals have their own habits justify it? Ans: The shape of a crystal in which it usually grows called habit of a crystal. The shape of the crystal remains same if its conditions remain same. For example When 10%urea is added in NaCl then needle like crystals are formed instead of cubic crystals Q.30 Justify that solids are rigid? Ans: The solids are very rigid. This rigidity is due to the fixed positions of the particles. The presence of strong cohesive forces makes particles unable to change their positions. This rigidity of solids can be changed under stress Q.31 Give types of solids? Ans. There are two types of solids: (i) Crystalline solids (ii) Amorphous solids Crystalline solids: Those solids in which atoms, ions or molecules are arranged to a definite three dimensional pattern,are called crystalline solids. Amorphous solids: Those solids whose constituent atoms, ions or molecules do not possess a regular orderly arrangement are called amorphous solids.The best examples are glass, plaster and rubber, glue, etc. Q.32 Define the following: (i) Cleavage planes. (ii) Anisotropy (iii) Symmetry (iv) Habit of a crystal Ans. (i) Cleavage planes: whenever the crystalline solids are broken they do so along definite planes. These planes are called the cleavage planes. (ii) Anisotropy: Some of the crystals show variation in physical properties depending upon the direction; such properties are called anisotrophic properties and the phenomenon is called anisotropy. (iii) Symmetry: The repetition of faces angles or edges when a crystal is rotated by 360o along its axis is called symmetry. (v) Habit of a crystal: The shape of a crystal in which it usually grows is called habit of crystal. Q.32 Define the following: (i) Isomorphism (ii) Polymorphism (iii) Allotropy (iv) Transition temperature (v) Crystal lattice (vi) Unit cell Ans. (i) Isomorphism: Isomorphism is the phenomenon in which two different substances exist in the same crystalline form. These different substance are called isomorphs of each other. Examples of ismorphs are NaNO3, CaCO3,K2SO4, K2CrO4. (ii) Polymorphism: Polymorphism is a phenomenon in which a substance exists in more than one crystalline forms. The substance which exists in more than one crystalline forms is called polymorphic, and these forms are called polymorphs of each other. Polymorphs have same chemical properties but they differ in the physical propertiesAgNO3, CaCO3 are polymorphs. (iii) Allotropy: The existence of an element in more than one crystalline form is known as allotropy and these forms of the element are called allotropes or allotropic forms. Element Crystalline forms Carbon Cubic (Diamond) Hexagonal (Graphite). (iv) Transition Temperature: It is that temperature at which two crystalline forms of the same substance can coexist in equilibrium with each other. At this temperature one crystalline form of substance changes to one another.95.50 Sulphur S8 (rhombic) sulphur S8 (monoclinic) (v) Crystal lattice: A crystal lattice is defined as an array of points representing atoms, ions or molecules of a crystal arranged at different sites in three dimensional space. (vi) Unit cell: The smallest part of crystal lattice has all the characteristic features of the entire crystal is called unit cell. The simplest unit cell is a cubic unit cell. Q.33 Name the crystal systems. Ans. (Cu T Or T He Mo Tri) 1. Cubic system 2. Tetragonal system 3. Orthorhombic or Rhombic system 4. Monoclinic system 5. Hexagonal system 6. Trigonal system 7. Triclinic system Q.34 Define lattice energy. Ans. The energy released when one mole of the ionic crystal is formed from the gaseous ions.It is also defined as the energy required to break one mole of solid into isolated ions in the gas phase. It is expressed in kJ mol–1. Na+(g) + Cl–(g) ® NaCl(s) D H – 792 kJ mol–1 or NaCl(3) ® Na+(g) + Cl–(g). Q.35 Describe the types of crystalline solids. Ans. There are four types of crystalline solids, depending upon the type of bond present in them. 1. Ionic solids. 2. Covalent solids. 3. Metallic solids. 4. Molecular solids. 1. Ionic Solids: Crystalline solids in which the particles forming the crystals are positively and negatively charged ions are called ionic solids.These ions are held together by strong electrostatic forces of attraction. These attractive forces are also called ionic bonds. The crystals of NaCl, KBr etc. are ionic solids. 2. Covalent solids: The crystalline solids in which atoms of similar or different elements are held together by covalent bonds are known as covalent solids. They are also called atomic solids. There are two types of covalent solids. Type 1: When covalent bonds give joint molecules like diamond, silicon carbide or Aluminum nitride. Type 2: When atoms join to form the covalent bonds and separate layers are produced like that of graphite, cadmium iodide and boron nitride. 3. Molecular solids: The solid substance in which the particles forming the crystals are polar or non–polar molecules or atoms, are called molecular solids. In solidified noble gases, there are non–polar atoms. Two types of intermolecular forces hold them together. 1. Dipole–dipole interactions 2. Vander Waal’s forces These intermolecular forces are much weaker then the forces of attraction between the cations and the anions in ionic crystals and between the atoms in the covalent crystals. Ice and the sugar are the best example of crystals having polar molecules, whereas iodine sulphur and carbon dioxide form crystals containing non–polar molecules. 4. Metallic solids: The crystalline solids in which the metal atoms are held together by metallic bonds are known as metallic solids. Metallic Bond: The force of attraction that binds positive metal ion to the number of electrons with in its sphere of influence is called metallic bond. Theories of metallic bond: 1. Electron gas theory 2. Valence bond theory 3. Molecular orbital theory Q.36 Iodine dissolves readily in tetrachloromethane. Give reason. Ans. We know that “like dissolve like”. Iodine is a non–polar substance. So it becomes solvable in non–polar solvent CCl4. Q.37 Justify molecular solids are soft and compressible? Ans: The forces which hold the molecules together in molecular structure are weak so, they are soft and compressible Q.38 What is crystallite? Ans: The small regions in amorphous solids where particles have a regular arrangement are called crystallites. Q.39 Why diamond is bad conductor of electricity? Ans: In diamond each carbon is SP3 hybridized there is no free electron to conduct electricity therefore it is bad conductor. Q.40 Why metals have shiny surface? Ans: When light falls on the surface of metals then the electrons are excited after de-exictation they emit energy in the form of light therefore they show shiny surface. Q.41 Why Na is soft while Cu is hard? Ans: In sodium only one mobile electron is present while in copper two mobile electrons are present due to strong metallic bond in copper it is hard. Q.42 Why ionic crystals are brittle? Ans: Because ionic solids are composed of parallel layers which contain cations & anions in alternate positions, so that the opposite ions in the various parallel layers lie over each other. When an external force is applied one layer of the ions slide pass over other layer. In this way due to repulsion of similar ions the crystals show brittleness. Q.43 Electrical conductivity of metals decreases by increasing temperature? Ans: With the increases in temperature the positive ions of metals also vibrate which hinders the motion of mobile electrons due to this hindrance electrical conductivity also decreases. Q.44 What is coordination no. of an ion? What is the coordination no of the cation in (a)NaCl, and(b)CsCl? Ans: The no. of positive ions which surround the anion called coordination no. of anion (a) Coordination no. Na in NaCl is 6 (b) Coordination no of Cs in CsCl is 8 (due to the greater size of Cs) |
CHAPTER#5 ATOMIC STRUCTURE
SHORT QUESTIONS AND ANSWERS
Q.1 Why it is necessary to decrease the pressure in the discharge tube to get the cathode rays?
Ans.
The current does not flow through the gas at ordinary pressure even at high voltage about 500 volts. However when the pressure inside the tube is decreased, the gas in the tube begins to conduct electricity at low pressure. Therefore it is necessary to decrease the pressure in the discharge tube to get the cathode rays.
Q.2 Which ever gas is used in the discharge tube the nature of the cathode rays remains the same why?
Ans.
A cathode ray consists of beam of electrons and electrons are constituents of all matter so, cathode rays do not depend upon the nature of the gas. Therefore, whichever gas is used in the discharge tube, the nature of cathode rays remains the same.
Q.3 Why e/m value of cathode rays is just equal to that of electrons?
Ans.
A cathode ray consists of beam of electrons, so cathode rays are actually electrons. Therefore e/m value of cathode ray is just equal to that of electron.
Q.4 The bending of the cathode rays in the electric and magnetic field show that they are negatively charged.
Ans.
The cathode ray beam travels in a straight line from the cathode to anode. The beam bends toward the south pole of the magnet when it passes through the magnetic field, which shows the cathode rays are negatively charged.
Q.5 Why positive rays are also called canal rays?
Ans.
Since positive rays produced in the discharge tube passed through the canals or holes of cathode, therefore positive rays are also called canal rays.
Q.6 The e/m values of positive rays for different gases are different but those for cathode rays, the e/m value is the same.
Ans.
The e/m value of positive rays depends upon the nature of gas used in the discharge tube. The characteristic of the gas varies from gas to gas, but for cathode rays e/m value is independent of the nature of the gas. Therefore, e/m values of positive rays for different gases are different but those for cathode rays the e/m value is the same.
Q.7 The e/m value for positive rays obtained from hydrogen gas 1836 times more than that of an electron?
Ans.
The mass of hydrogen gas is 1836 times more than that of an electron. Cathode rays consist of beam of electrons. The e/m value for positive rays depends upon the gas used in the tube, and e/m value for cathode rays is independent of the nature of the gas. Therefore e/m value for positive rays obtained from H2 gas is 1836 times less than that of cathode rays. Heavier the gas, the smaller the e/m value for positive rays.
Q.8 Justify, that cathode rays are material particles.
Ans.
Cathode rays drive a small paddle, wheel which shows that these rays posses momentum. From this observation, it is inferred that cathode rays are not rays but particles having a definite mass and velocity. Therefore cathode rays are material particles.
Q.9 How neutrons are produced?
Ans.
When a stream of a–particles from a polonium source is directed at beryllium target, penetrating radiations are produced, which are called neutrons.
He + Be ® C + n
Q.10 Why the neutrons are used as projectile?
Ans:
The particles, which hit the nucleus and can change its nature are called projectile. A projectile must be chargeless otherwise it will be captured or repelled by the nucleus. The slow moving neutrons cause nuclear reactions like fission and are used in artificial radioactivity. They are chargeless; therefore they can be used as projectile in nuclear research.
n + Cu → Cu + hv (γ- radiations)
Cu → Zn +-1e (β-particle)
Q.10 How are x–rays produced?
Ans.
X–rays are produced when fast moving electrons collide with heavy metal anode in the discharge tube.
Q.11 why the potential energy of bounded electron is negative in Bohr’s model?
Ans.
The potential energy of bounded electron is negative, because the energy of separated nucleus and electron is taken to be zero. As electron is brought from infinity towards the nucleus to form a stable state of the atom, energy is released because of attractive forces and the energy becomes less than zero, or negative. Therefore, the energy of the bounded electron is negative.
Q.12 Why the total energy of bounded electron in negative in Bohr’s model?
Ans.
The total energy of bounded electron is negative because the electron is under the force of attraction of the nucleus to have a stable state of the atom. More over when we calculate the total energy of the bounded electron, which is the sum of K.E. and P. E comes which is also negative.
Q.13 Explain that energy of an electron is inversely proportional to n2, but energy of higher orbits are always greater than those of the lower orbits in Bohr’s model.
Ans.
The energy of an electron in the nth orbit is
En = –
where e, m, 00 and h are all constants, thus En µ
The more negative the energy is the more stable will be the atom. The energy becomes successively less negative, therefore the energy values of higher orbits are always greater than those of the lower orbits.
Q.14 Explain the energy difference between adjacent levels goes on decreasing sharply in Bohr’s model.
Ans.
The energy difference between adjacent levels goes on decreasing, because the distance between the adjacent orbits increases.
Q.15 why does cathode rays produce shadow of an opaque object placed in their path.
Ans.
Any object which is material in nature, produces its shadow. Since cathode rays are material in nature, therefore, they produce shadow of an opaque object placed in their path.
Q.16 Give the main points of quantum theory of radiation.
Ans.
1. Energy is emitted or absorbed by atoms only in the form of packets called quantum.
2. The amount of energy associated with a quantum of radiation is proportional to the frequency (u) of the radiation.
E µ u
or E = hu
3. A body can emit or absorb energy only in terms of integral multiples of quantum.
E = nhu (where n = 1, 2, 3, 4, 5, ……..)
Q.17 Define frequency, wavelength and wave number.
Ans. Frequency (u):
The number of waves passing through a point per second is called frequency (u). Its units are cycles s–1.
Wavelength (l):
The distance between two successive crests or troughs is called wavelength “l” and is expressed in Ao or nm.
Wave number:
The number of waves per unit length is called wave number and is reciprocal of wave length.
=
The wave number is expressed (m–1) or per meter.
Q.18 What is spectrum? Differentiate between continuous spectrum and line spectrum.
Ans.
The dispersion of the components of white light, when it is passed through prism is called spectrum. The distribution among various wavelengths of the radiant energy emitted or absorbed by an object is also called spectrum.
Continuous spectrum:
A spectrum containing light of all wavelengths is called continuous spectrum.
In this type of spectrum, the boundary line between the colours cannot be marked. The colours diffuse into each other. One colour merges into another without any dark space. The best example of continuous spectrum is rainbow.
Line spectrum:
When an element or its compound is volatilized on a flame and the light emitted is seen through, a spectrometer. We see distinct lines separated by dark spaces. This type of spectrum is called line spectrum. This is the characteristic of an atom.
Q.19 Describe briefly Rutherford’s atomic model.
Ans.
According to Rutherford’s model most of the mass of the atom (99.95%) is concentrated in a positively charged centre, called nucleus around which the negatively charged electrons move.
Q.20 On which experiment Rutherford’s atomic model is based on, describe it briefly?
Ans.
Rutherford’s atomic model is based on the scattering of a–particles emitted from radioactive substances pass through the metal atoms of the foil undeflected by the light weight electrons. When an a–particle does happen to hit a metal–atom nucleus. However, it is scattered at a wide angle because it is repelled by the massive positively charged nucleus.
Q.21 Define orbit and orbital.
Ans. Orbit:
A definite circular path at a definite distance from the nucleus in which the electrons revolve around the nucleus is called an orbit.
K, L, M, N are orbits.
Orbital:
A three dimensional region or space around the nucleus, within which the probability of finding an electron is maximum called an orbital, s, p, d and f are atomic orbitals.
Q.22 What do you understand by wave particle duality and what is the de Broglei relation?
Ans.
According to de Broglei, all matter particles in motion have a dual character. It means that electrons, protons, neutrons, atoms, and molecules, possess the characteristics of both the material particle and a wave. This is called wave particle duality in matter.
De Broglei derived a mathematical equation which relates the wavelength (l) of the electron to the momentum of electron (mv)
l =
Where l = wavelength v = velocity of electron
M = mass of electron and h is Planck’s constant.
This equation l = is called de Broglie relation.
Q.23 What is Heisenberg’s uncertainty principles?
Ans.
Heisenberg showed that it is impossible to determine simultaneously both the position and momentum of an electron. Suppose that Dx is the uncertainty in the measurement of the position and Dp is the uncertainty in the measurement of momentum of an electron.
Dx . Dp ³
This relationship is called uncertainty principle.
Q.24 What are quantum numbers?
Ans.
The dimensionless numbers, rise naturally when the Schrodinger wave equation is solved for electron wave patterns and their energies are called quantum numbers.
These numbers describe the behaviour of electron in an atom completely.
There are four quantum numbers.
1. Principal quantum number “n”
It describes the energy of an electron in an atom. The value of n represents the shell or energy level in which the electron revolves around the nucleus. These shells are named as K, L, M, N, O, P, having the values of n, 1, 2, 3, 4, 5 and 6 respectively. The greater the value of n, the greater will be the distance from the nucleus and greater will be the energy of electron in the shell.
2. Azimuthal quantum number “l”
It determines the shape of orbital, it can have any integer value from 0 to n–l. this quantum number is used to represent the sub–shells, and these value are l = 0, 1, 2, 3. These values represent different sub–shells which are designated as s, p, d, and f, with values of l = 0, 1, 2, 3 respectively.
3. Magnetic quantum number (m)
It describes the orientation of the orbital in space. It can have all the integral values between + l and – l through zero i.e. + l …….. 0 …….. – l. For each value of l, there will be
(2l + 1) values of m. actually the values of m gives us the information of degeneracy of orbitals in space.
4. Spin quantum number (s)
It describes the spin of electron in atom. Since an electron can spin clockwise or anti clockwise, thus two possible values are + and – depending upon the spin of electron.
Q.25 What is n + l rule?
Ans.
This rule says that sub–shells are arranged in the increasing order of (n + l) values and if any two sub–shells have the same (n + l) values, then the sub–shell is filled first whose n values is smaller.
Q.26 What is the origin of line spectrum?
Ans.
According to Bohr’s theory each bright line in a line spectrum results from the downward jump of electron from a higher energy E2 to lower energy E1. This difference in energy (E2 – E1) is emitted as radiation of definite frequency in the form of spectral line.
According to the quantum theory of radiation,
E1 – E2 = hu
Or u =
Q.27 When is Zeeman effect?
Ans.
When the excited atoms of hydrogen are placed in a magnetic field, its spectral line are further split up in to closely spaced lines. This type of splitting of spectral lines is called Zeeman effect.
Q.28 What is stark effect?
Ans.
When the excited hydrogen atom are placed in an electric field, its spectral lines are further split up into closely spaced lines. This type of splitting of spectral lines is called stark effect.
Q.29 What is Mosely’s Law?
Ans.
Mosely’s law states that the frequency of spectral line in
x–ray spectrum varies as the square of atomic number of an element emitting it. This law convinces us that it is the atomic number and not the atomic mass of the element which determines its characteristic properties, both physical and chemical.
Q.30 Describe Summerfield’s modification of Bohr’s model atom.
Ans.
Summerfield suggested that the moving electron revolves in elliptical orbits in addition to circular orbit, with the nucleus situated at one of the foci of the ellipse. The elliptical paths of the moving electron go on changing their position in space, and the nucleus is buried by the electronic cloud from all the sides.
Q.31 Which of these orbitals, 3d or 4s has higher energy level?
Ans.
For 3d, n + l = 3 + 2 = 5 and for 4s, n + l = 4 + 0 = 4. Therefore 3d orbital has higher energy, than 4s orbital.
Q.32 How many maximum number of electron can have an orbital and a shell?
Ans.
An orbital can have maximum two electrons with opposite spins. A shell can have maximum of 2n2 electrons, where “n” is the principal quantum number. First shell can have maximum 2 electrons, 2nd shell have 8 electrons 3rd shell have 18 electrons etc.
Q.33 Distribute electrons in orbitals of 19K, 29Cu, 24Cr, 53I.
Ans.
19K ® 1s2, 2s2, 2p6, 3s2, 3p6, 4s1
29Cu ® 1s2, 2s2, 2p6, 3s2 3p6, 3d10, 4s1
24Cr ® 1s2, 2s2, 2p6, 3s2 3p6, 3d5, 4s1
53I ® 1s2, 2s2, 2p6, 3s2, 3p6, 3d10, 4s2, 4p6, 4d10, 5s2, 5p5
Q.34 What does it mean, when we say energy is quantized?
Ans.
Quantization means that energy can only be absorbed or emitted in specific amounts or multiples of these amounts. This minimum amount of energy is equal to a constant times the frequency of the radiation absorbed or emitted E = hv.
Q.35 Why do not we notice the quantization of energy in every day activities?
Ans.
In everyday activities, macroscopic objects such as our bodies gain or lose total amounts of energy much larger than a single quantum, hv. The gain or loss of the relatively minuscule quantum of energy in unnoticed.
Q.36 Explain the existence of line spectra is consistent with Bohr’s theory of quantized energies for the electron in the hydrogen atom.
Ans.
When applied to atoms, the notion of quantized energies means that only certain values of D E are allowed. These are represented by the lines in the emission spectra of excited atoms.
Q.37 In what ways does de Broglie’s hypothesis require revision of our picture of the H–atom based on Bohr’s model?
Ans.
De Broglie’s hypothesis not electrons have a characteristic wavelength requires, revision of Bohr’s particle only model. For example the idea of a fixed orbit for the electron in hydrogen is hard, to reconcile with the wave properties of electron.
Q.38 (a) For n = 4 what are possible values of l?
(b) For l = 2 what are the possible values of m.
Ans.
(a) n = 4 l = 3, 2, 1, 0
(b) l = 2 m = – 2, – 1, 0, 1, 2
Q.39 Which of the following are permissible sets of quantum numbers for an electron in a hydrogen atom?
(a) n = 2 l = 1 m = 1
(b) n = 1 l = 0 m = – 1
(c) n = 4 l = 2 m = – 2
(d) n = 3 l = 3 m = 0
Ans.
(a) permissible 2p (b) not permissible
(c) Permissible 4d (d) not permissible
Q.40 (a) What are the possible values of the electron spin quantum numbers?
(b) What piece of experimental equipment can be used to distinguish electrons that have different values of the electron spin quantum number?
(c) Two electrons in an atom both occupy the Is orbital. What quantity must be different for the two electrons? What principle governs the answer to this question?
Ans.
(a) + , –
(b) A magnet with a strong inhomogeneous magnetic field.
(c) They must have different spin quantum number values. The Pauli exclusion principle.
Q.41 Give region of different spectral lines.
Ans.
1. Lyman series (U. V. region)
2. Balmer series (visible region)
3. Paschen series (I. R. region)
4. Bracket series (I. R. region)
5. Pfund series (I. R. region)
Q.1 Why it is necessary to decrease the pressure in the discharge tube to get the cathode rays?
Ans.
The current does not flow through the gas at ordinary pressure even at high voltage about 500 volts. However when the pressure inside the tube is decreased, the gas in the tube begins to conduct electricity at low pressure. Therefore it is necessary to decrease the pressure in the discharge tube to get the cathode rays.
Q.2 Which ever gas is used in the discharge tube the nature of the cathode rays remains the same why?
Ans.
A cathode ray consists of beam of electrons and electrons are constituents of all matter so, cathode rays do not depend upon the nature of the gas. Therefore, whichever gas is used in the discharge tube, the nature of cathode rays remains the same.
Q.3 Why e/m value of cathode rays is just equal to that of electrons?
Ans.
A cathode ray consists of beam of electrons, so cathode rays are actually electrons. Therefore e/m value of cathode ray is just equal to that of electron.
Q.4 The bending of the cathode rays in the electric and magnetic field show that they are negatively charged.
Ans.
The cathode ray beam travels in a straight line from the cathode to anode. The beam bends toward the south pole of the magnet when it passes through the magnetic field, which shows the cathode rays are negatively charged.
Q.5 Why positive rays are also called canal rays?
Ans.
Since positive rays produced in the discharge tube passed through the canals or holes of cathode, therefore positive rays are also called canal rays.
Q.6 The e/m values of positive rays for different gases are different but those for cathode rays, the e/m value is the same.
Ans.
The e/m value of positive rays depends upon the nature of gas used in the discharge tube. The characteristic of the gas varies from gas to gas, but for cathode rays e/m value is independent of the nature of the gas. Therefore, e/m values of positive rays for different gases are different but those for cathode rays the e/m value is the same.
Q.7 The e/m value for positive rays obtained from hydrogen gas 1836 times more than that of an electron?
Ans.
The mass of hydrogen gas is 1836 times more than that of an electron. Cathode rays consist of beam of electrons. The e/m value for positive rays depends upon the gas used in the tube, and e/m value for cathode rays is independent of the nature of the gas. Therefore e/m value for positive rays obtained from H2 gas is 1836 times less than that of cathode rays. Heavier the gas, the smaller the e/m value for positive rays.
Q.8 Justify, that cathode rays are material particles.
Ans.
Cathode rays drive a small paddle, wheel which shows that these rays posses momentum. From this observation, it is inferred that cathode rays are not rays but particles having a definite mass and velocity. Therefore cathode rays are material particles.
Q.9 How neutrons are produced?
Ans.
When a stream of a–particles from a polonium source is directed at beryllium target, penetrating radiations are produced, which are called neutrons.
He + Be ® C + n
Q.10 Why the neutrons are used as projectile?
Ans:
The particles, which hit the nucleus and can change its nature are called projectile. A projectile must be chargeless otherwise it will be captured or repelled by the nucleus. The slow moving neutrons cause nuclear reactions like fission and are used in artificial radioactivity. They are chargeless; therefore they can be used as projectile in nuclear research.
n + Cu → Cu + hv (γ- radiations)
Cu → Zn +-1e (β-particle)
Q.10 How are x–rays produced?
Ans.
X–rays are produced when fast moving electrons collide with heavy metal anode in the discharge tube.
Q.11 why the potential energy of bounded electron is negative in Bohr’s model?
Ans.
The potential energy of bounded electron is negative, because the energy of separated nucleus and electron is taken to be zero. As electron is brought from infinity towards the nucleus to form a stable state of the atom, energy is released because of attractive forces and the energy becomes less than zero, or negative. Therefore, the energy of the bounded electron is negative.
Q.12 Why the total energy of bounded electron in negative in Bohr’s model?
Ans.
The total energy of bounded electron is negative because the electron is under the force of attraction of the nucleus to have a stable state of the atom. More over when we calculate the total energy of the bounded electron, which is the sum of K.E. and P. E comes which is also negative.
Q.13 Explain that energy of an electron is inversely proportional to n2, but energy of higher orbits are always greater than those of the lower orbits in Bohr’s model.
Ans.
The energy of an electron in the nth orbit is
En = –
where e, m, 00 and h are all constants, thus En µ
The more negative the energy is the more stable will be the atom. The energy becomes successively less negative, therefore the energy values of higher orbits are always greater than those of the lower orbits.
Q.14 Explain the energy difference between adjacent levels goes on decreasing sharply in Bohr’s model.
Ans.
The energy difference between adjacent levels goes on decreasing, because the distance between the adjacent orbits increases.
Q.15 why does cathode rays produce shadow of an opaque object placed in their path.
Ans.
Any object which is material in nature, produces its shadow. Since cathode rays are material in nature, therefore, they produce shadow of an opaque object placed in their path.
Q.16 Give the main points of quantum theory of radiation.
Ans.
1. Energy is emitted or absorbed by atoms only in the form of packets called quantum.
2. The amount of energy associated with a quantum of radiation is proportional to the frequency (u) of the radiation.
E µ u
or E = hu
3. A body can emit or absorb energy only in terms of integral multiples of quantum.
E = nhu (where n = 1, 2, 3, 4, 5, ……..)
Q.17 Define frequency, wavelength and wave number.
Ans. Frequency (u):
The number of waves passing through a point per second is called frequency (u). Its units are cycles s–1.
Wavelength (l):
The distance between two successive crests or troughs is called wavelength “l” and is expressed in Ao or nm.
Wave number:
The number of waves per unit length is called wave number and is reciprocal of wave length.
=
The wave number is expressed (m–1) or per meter.
Q.18 What is spectrum? Differentiate between continuous spectrum and line spectrum.
Ans.
The dispersion of the components of white light, when it is passed through prism is called spectrum. The distribution among various wavelengths of the radiant energy emitted or absorbed by an object is also called spectrum.
Continuous spectrum:
A spectrum containing light of all wavelengths is called continuous spectrum.
In this type of spectrum, the boundary line between the colours cannot be marked. The colours diffuse into each other. One colour merges into another without any dark space. The best example of continuous spectrum is rainbow.
Line spectrum:
When an element or its compound is volatilized on a flame and the light emitted is seen through, a spectrometer. We see distinct lines separated by dark spaces. This type of spectrum is called line spectrum. This is the characteristic of an atom.
Q.19 Describe briefly Rutherford’s atomic model.
Ans.
According to Rutherford’s model most of the mass of the atom (99.95%) is concentrated in a positively charged centre, called nucleus around which the negatively charged electrons move.
Q.20 On which experiment Rutherford’s atomic model is based on, describe it briefly?
Ans.
Rutherford’s atomic model is based on the scattering of a–particles emitted from radioactive substances pass through the metal atoms of the foil undeflected by the light weight electrons. When an a–particle does happen to hit a metal–atom nucleus. However, it is scattered at a wide angle because it is repelled by the massive positively charged nucleus.
Q.21 Define orbit and orbital.
Ans. Orbit:
A definite circular path at a definite distance from the nucleus in which the electrons revolve around the nucleus is called an orbit.
K, L, M, N are orbits.
Orbital:
A three dimensional region or space around the nucleus, within which the probability of finding an electron is maximum called an orbital, s, p, d and f are atomic orbitals.
Q.22 What do you understand by wave particle duality and what is the de Broglei relation?
Ans.
According to de Broglei, all matter particles in motion have a dual character. It means that electrons, protons, neutrons, atoms, and molecules, possess the characteristics of both the material particle and a wave. This is called wave particle duality in matter.
De Broglei derived a mathematical equation which relates the wavelength (l) of the electron to the momentum of electron (mv)
l =
Where l = wavelength v = velocity of electron
M = mass of electron and h is Planck’s constant.
This equation l = is called de Broglie relation.
Q.23 What is Heisenberg’s uncertainty principles?
Ans.
Heisenberg showed that it is impossible to determine simultaneously both the position and momentum of an electron. Suppose that Dx is the uncertainty in the measurement of the position and Dp is the uncertainty in the measurement of momentum of an electron.
Dx . Dp ³
This relationship is called uncertainty principle.
Q.24 What are quantum numbers?
Ans.
The dimensionless numbers, rise naturally when the Schrodinger wave equation is solved for electron wave patterns and their energies are called quantum numbers.
These numbers describe the behaviour of electron in an atom completely.
There are four quantum numbers.
1. Principal quantum number “n”
It describes the energy of an electron in an atom. The value of n represents the shell or energy level in which the electron revolves around the nucleus. These shells are named as K, L, M, N, O, P, having the values of n, 1, 2, 3, 4, 5 and 6 respectively. The greater the value of n, the greater will be the distance from the nucleus and greater will be the energy of electron in the shell.
2. Azimuthal quantum number “l”
It determines the shape of orbital, it can have any integer value from 0 to n–l. this quantum number is used to represent the sub–shells, and these value are l = 0, 1, 2, 3. These values represent different sub–shells which are designated as s, p, d, and f, with values of l = 0, 1, 2, 3 respectively.
3. Magnetic quantum number (m)
It describes the orientation of the orbital in space. It can have all the integral values between + l and – l through zero i.e. + l …….. 0 …….. – l. For each value of l, there will be
(2l + 1) values of m. actually the values of m gives us the information of degeneracy of orbitals in space.
4. Spin quantum number (s)
It describes the spin of electron in atom. Since an electron can spin clockwise or anti clockwise, thus two possible values are + and – depending upon the spin of electron.
Q.25 What is n + l rule?
Ans.
This rule says that sub–shells are arranged in the increasing order of (n + l) values and if any two sub–shells have the same (n + l) values, then the sub–shell is filled first whose n values is smaller.
Q.26 What is the origin of line spectrum?
Ans.
According to Bohr’s theory each bright line in a line spectrum results from the downward jump of electron from a higher energy E2 to lower energy E1. This difference in energy (E2 – E1) is emitted as radiation of definite frequency in the form of spectral line.
According to the quantum theory of radiation,
E1 – E2 = hu
Or u =
Q.27 When is Zeeman effect?
Ans.
When the excited atoms of hydrogen are placed in a magnetic field, its spectral line are further split up in to closely spaced lines. This type of splitting of spectral lines is called Zeeman effect.
Q.28 What is stark effect?
Ans.
When the excited hydrogen atom are placed in an electric field, its spectral lines are further split up into closely spaced lines. This type of splitting of spectral lines is called stark effect.
Q.29 What is Mosely’s Law?
Ans.
Mosely’s law states that the frequency of spectral line in
x–ray spectrum varies as the square of atomic number of an element emitting it. This law convinces us that it is the atomic number and not the atomic mass of the element which determines its characteristic properties, both physical and chemical.
Q.30 Describe Summerfield’s modification of Bohr’s model atom.
Ans.
Summerfield suggested that the moving electron revolves in elliptical orbits in addition to circular orbit, with the nucleus situated at one of the foci of the ellipse. The elliptical paths of the moving electron go on changing their position in space, and the nucleus is buried by the electronic cloud from all the sides.
Q.31 Which of these orbitals, 3d or 4s has higher energy level?
Ans.
For 3d, n + l = 3 + 2 = 5 and for 4s, n + l = 4 + 0 = 4. Therefore 3d orbital has higher energy, than 4s orbital.
Q.32 How many maximum number of electron can have an orbital and a shell?
Ans.
An orbital can have maximum two electrons with opposite spins. A shell can have maximum of 2n2 electrons, where “n” is the principal quantum number. First shell can have maximum 2 electrons, 2nd shell have 8 electrons 3rd shell have 18 electrons etc.
Q.33 Distribute electrons in orbitals of 19K, 29Cu, 24Cr, 53I.
Ans.
19K ® 1s2, 2s2, 2p6, 3s2, 3p6, 4s1
29Cu ® 1s2, 2s2, 2p6, 3s2 3p6, 3d10, 4s1
24Cr ® 1s2, 2s2, 2p6, 3s2 3p6, 3d5, 4s1
53I ® 1s2, 2s2, 2p6, 3s2, 3p6, 3d10, 4s2, 4p6, 4d10, 5s2, 5p5
Q.34 What does it mean, when we say energy is quantized?
Ans.
Quantization means that energy can only be absorbed or emitted in specific amounts or multiples of these amounts. This minimum amount of energy is equal to a constant times the frequency of the radiation absorbed or emitted E = hv.
Q.35 Why do not we notice the quantization of energy in every day activities?
Ans.
In everyday activities, macroscopic objects such as our bodies gain or lose total amounts of energy much larger than a single quantum, hv. The gain or loss of the relatively minuscule quantum of energy in unnoticed.
Q.36 Explain the existence of line spectra is consistent with Bohr’s theory of quantized energies for the electron in the hydrogen atom.
Ans.
When applied to atoms, the notion of quantized energies means that only certain values of D E are allowed. These are represented by the lines in the emission spectra of excited atoms.
Q.37 In what ways does de Broglie’s hypothesis require revision of our picture of the H–atom based on Bohr’s model?
Ans.
De Broglie’s hypothesis not electrons have a characteristic wavelength requires, revision of Bohr’s particle only model. For example the idea of a fixed orbit for the electron in hydrogen is hard, to reconcile with the wave properties of electron.
Q.38 (a) For n = 4 what are possible values of l?
(b) For l = 2 what are the possible values of m.
Ans.
(a) n = 4 l = 3, 2, 1, 0
(b) l = 2 m = – 2, – 1, 0, 1, 2
Q.39 Which of the following are permissible sets of quantum numbers for an electron in a hydrogen atom?
(a) n = 2 l = 1 m = 1
(b) n = 1 l = 0 m = – 1
(c) n = 4 l = 2 m = – 2
(d) n = 3 l = 3 m = 0
Ans.
(a) permissible 2p (b) not permissible
(c) Permissible 4d (d) not permissible
Q.40 (a) What are the possible values of the electron spin quantum numbers?
(b) What piece of experimental equipment can be used to distinguish electrons that have different values of the electron spin quantum number?
(c) Two electrons in an atom both occupy the Is orbital. What quantity must be different for the two electrons? What principle governs the answer to this question?
Ans.
(a) + , –
(b) A magnet with a strong inhomogeneous magnetic field.
(c) They must have different spin quantum number values. The Pauli exclusion principle.
Q.41 Give region of different spectral lines.
Ans.
1. Lyman series (U. V. region)
2. Balmer series (visible region)
3. Paschen series (I. R. region)
4. Bracket series (I. R. region)
5. Pfund series (I. R. region)
CHAPTER#6 ******************CHEMICAL BONDING*********************
SHORT QUESTION WITH ANSWERS
Q.1 Dipole moments of chlorobenzene is 1.70 D and of chlorobenzene is 2.5 D while that of paradichlorbenzene is zero; why?
Ans.
Benzene has zero dipole moment as it is a symmetrical planar hexagonal molecule. The substitution of benzene ring with two Cl–atoms at the para positions does not add any dipole moment to the benzene as the dipoles created being equal and opposite cancel out each other’s effect. As, there is no such cancellation in monochlorobenzene. These molecules have resultant dipole moments of 1.70 D and 2.5 D respectively.
Q.2 What is covalent bond?
Ans.
A chemical bond formed by the sharing of a pair of electrons between atoms is called covalent bond.
H + H ® H – H
Q.3 What is meant by a coordinate covalent bond?
Ans.
A coordinate covalent bond is for a bond formed, when both electrons of the bond are denoted by one atom.
Q.4 What is the difference between a localized p bond and a delocalized one?
Ans.
In a localized p bond, the electron density is concentrated between the two atoms forming the bond. In a delocalized p bond, the electron density is spread over all the atoms that contribute p orbitals to the net work.
Q.5 How will you differentiate between a polar covalent bond and non–polar covalent bond?
Ans.
A covalent bond between two dissimilar atoms in which the shared electron pair is not attracted equally by the two atoms and the bonded atoms acquire a partial positive and negative charge is called polar covalent bond.
A covalent bond between two like atoms such as H–H,
Cl–Cl, in which shared electron pair is attracted equally by both the atoms is called a non–polar covalent bond.
Q.6 Indicate the hybridization and bond angles associated with each of the following (a) linear (b) tetrahedral (c) trigonal planar.
Ans.
(a) Sp. 180o (b) Sp3. 109o (c) Sp2. 120o
Q.7 What are the similarities and differences between atomic orbitals and molecular orbitals?
Ans.
Both atomic and molecular orbitals have a characteristic energy and shape each can hold a maximum of the two electrons. Atomic orbitals are localized and their energies are the result of interactions between the subatomic particles in a single atom. Molecular orbitals can be delocalized and their energies are influenced by interactions between electrons on several atoms.
Q.8 Why is the bonding molecular orbital of H2 at lower energy than the electron in a hydrogen atom?
Ans.
There is a net lowering in energy that accompanies bond formation because the electrons in H2are strongly attracted to both H nuclei, while in H–atom the electron is attracted by only one nucleus.
Q.9 How many electrons can be placed into each molecular orbital of a molecule?
Ans.
Two electrons can be placed into each molecular orbital of a molecule.
Q.10 What is meant by an ionic bond?
Ans.
An ionic bond is a chemical bond formed by the electzostatic attraction between positive and negative ions. The bond formed between two atoms when one or more electrons are transformed from valence shell of one atom to the valence shell of the other. The atom that loses electrons becomes cation and the atom that gains electrons becomes anion. The electronic configurations of both ions are those of noble gas atoms. All metals react with non–metals to form ionic compounds.
Na(g) ® Na+ + e–
Cl(g) + e– ® Cl–(g)
Na+(g) + Cl–(g) ® NaCl
Q.11 Can a molecule have polar bonds and not a dipole?
Ans.
Yes a molecule can have polar bonds and not a dipole, if the orientation of the polar bonds in the molecule cancel each other’s effect e.g. CO2.
Q.12 Can a molecule have non–polar bonds only and have a dipole?
Ans.
No. A molecule which have non–polar bonds only cannot have a dipole.
Q.13 State the difference between a polar bond and a polar molecule.
Ans.
A polar bond is a covalent bond that exists between two atoms having an electronegativity difference greater than 0.2 a polar molecule results, if one or more polar bonds in a molecule is not balanced by other polar bonds in the
Q.14 According to molecular orbital theory, would Be2 be expected to exist? Explain.
Ans.
Be2 is not expected to exist, it has bond order of zero and is not energetically favoured over isolated Be atoms.
Q.15 What is meant by bond order?
Ans.
The bond order is half the difference between the number of bonding electrons and the number of antibonding electrons.
Bond order =
(No. of bonding electrons – No. of antibonding electrons)
Q.16 What is meant by paramagnetic substance?
Ans.
A paramagnetic substance is a substance that is attracted by a magnetic field and this attraction is generally the result of unpaired electrons. The more unpaired electrons in a species (substance), the stronger the force of attraction. This type of magnetic behaviour is called paramagnetism.
Q.17 What is meant by a diamagnetic substance?
Ans.
A diamagnetic substance is a substance that is not attracted by a magnetic field. This property is called diamagnetism. This property generally means that substance has only paired electrons.
Q.18 (a) what is hybridization at carbon atom in CH4, C2H4 and C2H2.
(b) The carbon atom in CH4 cannot participate in multiple bonding, whereas that is C2H4 can. Explain.
Ans.
(a) The hybridization at the carbon atom in CH4 is Sp3, in C2H4 is sp2 and in C2H2 is sp.
(b) The C atom in CH4 is sp3 hybridized, there are no un–hybridized p orbitals, available for the poverlap required for multiple bonds. In C2H4, the C atom is sp2 hybridized with two p atomic orbitals (one on each C atom), available to form the p overlap in the C = C double bond.
Q.19 Describe briefly the VSEPR theory.
Ans.
The VSEPR theory predicts the shape of molecules and ions in which valence shell electron pairs are arranged around the central atom of a molecule or ions in such a way that there is maximum separation so that electron repulsions are minimized and electron nucleus attractions are maximized. Some of these electron pairs are bonding and some are lone pairs. The direction in space of the bonding pairs gives the molecular geometry. A lone pair of electrons occupy more space than a bonding pair. Repulsive forces decrease sharply with increasing interpair angle. They are strong at 90o much weaker at 120o and very weak at 180o. in VSEPR model, each multiple bond is treated as though it were a single electron pair.
Q.20 A lone pair of electrons occupies more space than a bond pair?
Ans:
A lone pair of electrons occupies more space than bond pair because lone pair is attracted by only one nucleus while bond pair is attracted by two nuclei. Due to less nuclear attraction to lone pair its electronic charge is spread ot more in space than that of bond pair.
Q.21 Predict the geometry of (a) BeCl2 (b) BF3 (c) SiH4 molecules.
Ans.
(a) The two bond pairs of electrons in BeCl2 molecule arrange themselves as far apart as possible to minimize the repulsion between them. The only arrangement which can satisfy this condition is linear i.e. at angle of 180o.
: Cl ––– –– Cl
(b) The BF3 molecule containing three bond pairs of electrons is trigonal planar because, this structure gives maximum separation among the three bonding electron pairs.
::: :
(c) In SiH4 molecule, the electrostatic repulsion between four bonding electron pairs will be minimum, when they are present at corners of a regular tetrahedron making angle 109.5o with each other.
Q.22 Describe briefly the valence bond theory.
Ans.
According to valence bond theory, a covalent bond is formed by pairing of electrons by the overlap of half (partially) filled atomic orbitals of two atoms. The two overlapping orbitals must be valence orbitals, must be half filled and must retain their identities. By overlap means that the electrons of overlapping orbitals share a common region of high electron density, along the line between two nuclei called bond axis.
This theory explains the bonding in terms of overlapping of atomic orbitals and mixing of atomic orbitals called hybridization. Multiple bond occurs via the overlap of atomic orbitals to give sbonds and p bonds.
Q.23 Why monoatomic cations smaller than their corresponding neutral atoms?
Ans.
Electrostatic repulsions are reduced by removing an electron from a neutral atom, effective neutral charge increases and therefore the cation is smaller than their corresponding neutral atoms.
Q.24 Why are monoatomic anions larger than their corresponding neutral atoms?
Ans.
The additional electrostatic repulsion produced by adding an electron to a neutral atom decreases the effective nuclear charge experienced by the valence electrons and increases the size of anions.
Q.25 Why does the size of ions increase as one produced down a column in the periodic table?
Ans.
Going down a column, valence electrons are further from the nucleus and they experience greater shielding by core electrons. The greater radial extent of the valence electrons outweigh the increase in atomic number. Therefore the size of ions increase as one proceeds down a column.
Q.26 What is an isoelectronic series?
Ans.
An isoelectronic series is a group of atoms or ions that have the same number of electrons, and thus the same electronic configuration.
Q.27 Why noble gases are most stable?
Ans.
Noble gases are most stable because their s and p orbitals are completely filled.
Q.28 Why O2 molecule is paramagnetic in nature?
Ans.
O2 molecule is paramagnetic in nature due to presence of unpaired electrons in its molecule.
Q.29 Why CO2 is non–polar molecule although C–O bond is polar?
Ans.
Each C–O bond in CO2 is polar. The two bond dipoles in CO2 are equal in magnitude and are exactly opposite in direction. The bond dipoles cancel each other. Therefore, the overall dipole moment of CO2 is zero. Thus CO2 is a non–polar molecule.
Od– = C = Od+
Q.30 Why H2O is a polar molecule?
Ans.
H2O is a bent molecule with two polar bonds. Both the bonds are identical, so the bond dipoles are equal in magnitude. Because the molecule is bent, however the bond dipoles do not directly oppose to each other and therefore, do not cancel each other.
Hence the H2O molecule has an overall dipole moment
(m = 1.85D), because H2O has dipole moment, it is polar molecule.
Q.31 The melting and boiling points of electrovalent compounds are very high as compared with those of covalent compounds. Explain.
Ans.
The melting and boiling points of electrovalent compounds are very high, because the ions are tightly packed in the crystal lattice by strong attractive forces and high thermal energy is required to separate them from one another. The atoms of molecules in covalent compounds are held together by weak intermolecular forces and less energy is required to separate the atoms or molecules in a solid or liquid.
Q.32 (a) Why solid NaCl does not conduct electricity?
(b) What will happen if electric currents is passed through molten NaCl or its aqueous solution? Explain.
Ans.
(a) Solid NaCl does not conduct electricity because the oppositely charged ions are held together by strong electrostatic forces and the ions do not free to move.
(b) However, in the molten NaCl or its aqueous solution the ions because quite free to migrate in an electric field and conduct electricity when passed through them.
Q.33 What is molecular orbital theory?
Ans.
According to molecular orbital theory, atomic orbitals overlap to form molecular orbitals n atomic orbitals combine to form n molecular orbitals. Half of them are bonding molecular orbitals and half antibonding molecular orbitals. In this combination, the individual atomic orbital character is lost in order to form an entirely new orbital that belongs to whole molecule. The theory successfully explains bond order and paramagnetic property of O2.
Q.34 In many cases, the distinction between a coordinate covalent and a covalent bond vanishes after bond formation. Explain with the help of an example.
Ans.
A coordinate covalent bond is not essentially different from other covalent bonds, it involves the sharing of pair of electrons between two atoms. An example is formation of NH4+ ion in which all bonds are identical, so the distinction between a coordinate covalent bond and covalent bond vanishes after bond formation.
H+ + :NH3 ® H – ® H+
or
Q.35 PF3 is a polar molecule with dipole moment 1.02 D and thus the P–F bond is polar. Si, being in proximity of P in the periodic table, it is expected that Si–F bond would also be polar, but SiF4 has no dipole moment. Explain why it is so?
Ans.
PF3 has one lone pair of electrons and one P atom lies out of plane of rest of the atom and thus PF3 is pyramidal molecule and has 1.02 D dipole moment and thus the P–F bond is polar.
SiF4 molecular has tetrahedral shape and the SiF bonds are directed by Symmetrically about the central Si atom. Although Si–F bonds are polar, but all the four bond moments cancel out one another and give SiF4 molecule of zero dipole moment.
Q.36 NaCl is a harder substance at room temperature than glucose explain.
Ans.
The hardness of substance depends on the strength of the forces between the particles forming a substance. NaCl is an ionic compound and consists of Na+ and Cl– ions which are held together by strong electrostatic forces of attraction while glucose consists of molecules which are held together by weak intermolecular forces. Therefore, NaCl is a harder substance at room temperature than glucose.
Q.37 The linear of BeCl2 suggests that central Be atom is sp–hybridized. What type of hybridization a central atom undergoes when the atoms bonded to it are located at the corners of (a) an equilateral triangle and (b) a regular tetrahedron.
Ans.
(a) The central atom undergoes Sp2 hybridization when the atoms bonded to it are located at the corners of an equilateral triangle.
(b) The central atom undergoes Sp3 hybridization when the atoms bounded to it are located at the corners of a regular tetrahedron.
Q.38 A double bond is shorter and stronger than a single bond.
Ans.
The greater the number of electron pairs in the bond, the shorter and stronger will be the bond because of greater nuclei–electron attractions. A double bond has two shared electron pairs, while a single bond has only one shared electron pair. Thus a double bond has a greater nuclei electron attraction, than a single bond. Therefore, a double bond is shorter and stronger than a single bond.
Q.39 NH3 and H2O can form coordinate covalent bond with H+ but CH4 cannot do so.
Ans.
NH3 and H2O both have lone pair of electrons on N and O atoms which can donate to a H+(electron deficient) to from a coordinate covalent bond. In CH4 there is no lone pair of electrons to donate H+ for the formation of coordinate covalent bond. Therefore, NH3 and H2O can form coordinate covalent bond with H+ but CH4 cannot do so.
H3N: + H+ ® [H3N ® H]+
H2O + H+ ® H3O ® [H2O+ ® H]+
Q.40 Covalent bond may be non–polar but coordinate covalent bond is always polar.
Ans.
A covalent bond between two like atoms is always non–polar, whereas between two unlike atoms, it is a polar. In coordinate covalent bond the shared electron pair is denoted by only one of the two bonded atoms. The atom which denotes the electron pair acquires partial positive charge and the atom which accepts the electron pair acquires partial negative charge. Therefore, coordinate covalent bond is always polar. Hence a covalent bond may be non–polar, but coordinate covalent bond is always polar.
Q.41 Molecule of O2 is paramagnetic in nature. Explain.
Ans.
A substance is paramagnetic, when it has unpaired electrons.
According to molecular orbital theory, O2 has two unpaired electrons in the degenerate orbitalsp2 py and p*2 px. Due to the presence of these two unpaired electrons O2 molecule is paramagnetic.
Q.42 Dipole moment of CO2 is zero but that of CO is o.12 Debye.Why?
Ans:
CO2 is a linear molecule and the tow diploes cancel the effect of each other. In CO there is a single dipole directed from carbon to oxygen and it not cancelled.
Q.43 Why dipole moment of benzene is zero?
Ans:
Benzene is symmetrical planar molecule. It has six C – H there are six dipole moments. All the dipole moments cancel the effect of each other and net value is zero.
Q.44 I.E is index of metallic character why?
Ans:
Metals have loosely held electrons which are delocalized and are responsible for the properties of metals. So, metals have low ionization energies.
Q.45 The abnormality of the bond length and bond strength in HI is less prominent than that of HCl, give reason?
Ans:
Chlorine has higher electronegative than iodine. So, the polarities of HCI and HI bonds are unequal. Therefore, abnormality of bond length and bond strength of HCI is more prominent than HI.
Q.46 How does electro negativity difference decide the nature of ionic bond?
Ans:
When the electro negativity difference between two bonded atoms is 1.7 or move than that, then the bond is said to be ionic, otherwise, covalent. The % age of ionic character is more the 51% when the electro negativity difference is 1.7.
Q.47 Why NH3 and PH3 give coordinate covalent bond?
Ans:
NH3 and PH3 have lone pairs of electrons, which can be donated to H+ to make a coordinate covalent bond. In this way, NH4+ and PH4+ are produced which have perfect tetrahedral structure and all the four bonds have perfectly equal status.
Q.48 Most of the elements of the periodic table attain the electronic configuration of inert gases during bond formation. Justify it
Ans:
Inert gases are not reactive due to complete octet except He, Most of the S- and P- block elements may attain eight electrons in the outermost orbitals they do so either by losing gaining or sharing the electrons.
Q.49 Define octet rule?
Ans:
The tendency of the atoms to attain a maximum of eight electrons in the valence shell is called octet rule.
Q.50 What is an ionic radius?
Ans:
The ionic radius of an ion is the radius of the ion while considering it to be spherical in shape.
Q.51 What is a covalent radius?
The covalent radius of an element is defined as half of the single bond length between tow similar atoms covalently bonded in a molecule.
Q.52 What is difference between sigma & pi bond?
Ans:
SIGMA BOND:
The bond which is formed by the head to head overlapping called sigma bond. The electron density is present between two nuclei.
Pi BOND:
The bond which is formed by the sideways overlapping of two half filled orbitals. The electron density is present above and below the line joining the two nuclei.
Q.53 Why the size of an atom can not be measured directly?
Ans:
The size of an atom can not be measured directly due following reasons:
(i)There is no sharp boundary of an atom. The probability of finding an electron never becomes zero even at larger distances from the nucleus.
(ii) The electronic probability distribution is affected by neighbouring atoms. For this reason the size of an atom may change from one compound to another.
Q.54 Why E.A of Flourine is less than the expected value?
Ans
Since the size of fluorine is very small when electron is added in the fluorine it is strongly repelled by the already existing electrons. An extra amount of energy is provided to add an electron therefore its electron affinity is less than the expected value.
Q.55 Why sigma bond is stronger than pi bond? Or Why pi bond is more diffused than sigma bond?
Ans:
Sigma bond is more diffused than pi bond due to the linear overlapping of orbitals. Moreover electron density is present between two nuclei which is strongly attracted by two nuclei. While in pi bond electron density is not strongly attracted by two nuclei therefore it is weak than pi bond
Q.56 Define bond length.
Ans.
The distance between the nuclei of two atoms forming a covalent bond is called bond length. In general it is the sum of the covalent radii of the combined atoms.
Q.57 What is dipole moment? What are its units?
Ans.
The dipole moment may be defined as the product of electric charge (q) and distance (r) between the two oppositely charged centres. It is vector quantity as it has magnitude and direction. It plays a major role in determining the % age ionic character of a covalent bond and the shapes of molecules.
The dipole moment is measured in Debye units (D). It is denoted by symbol m.
Q.58 Define bond energy Give its units.
Ans.
The bond energy is defined as the average amount of energy required to break all bonds of particular type in one mole of substance. It is determined by measuring the heat involved in a chemical reaction.
It is also defined as the energy required to break Avogadro’s number (6.02 x 1023) of bonds or the energy released when an Avagadro number of bonds are formed. It is a measure of strength of bonds. The bond energy is measured in KJ mol–1.
Q.59 Define the following terms.
(a) Ionization energy
(b) Electron affinity
(c) electronegativity
Ans.
(a) Ionization energy:
The minimum amount of energy required to remove an electron from an atom is called ionization energy.
It depends upon the atomic size, nuclear charge and shielding effect of electrons.
(b) Electron Affinity:
The minimum amount of energy released when an electron is added to an isolated neutral gaseous atom in the lowest energy state to produce an anion is called electron affinity. It is measured in kJ mol–1.
(c) Electronegativity:
The tendency of an atom to attract shared pair of electron towards itself is called electronegativity. It is measured in electron volts.
Q.1 Dipole moments of chlorobenzene is 1.70 D and of chlorobenzene is 2.5 D while that of paradichlorbenzene is zero; why?
Ans.
Benzene has zero dipole moment as it is a symmetrical planar hexagonal molecule. The substitution of benzene ring with two Cl–atoms at the para positions does not add any dipole moment to the benzene as the dipoles created being equal and opposite cancel out each other’s effect. As, there is no such cancellation in monochlorobenzene. These molecules have resultant dipole moments of 1.70 D and 2.5 D respectively.
Q.2 What is covalent bond?
Ans.
A chemical bond formed by the sharing of a pair of electrons between atoms is called covalent bond.
H + H ® H – H
Q.3 What is meant by a coordinate covalent bond?
Ans.
A coordinate covalent bond is for a bond formed, when both electrons of the bond are denoted by one atom.
Q.4 What is the difference between a localized p bond and a delocalized one?
Ans.
In a localized p bond, the electron density is concentrated between the two atoms forming the bond. In a delocalized p bond, the electron density is spread over all the atoms that contribute p orbitals to the net work.
Q.5 How will you differentiate between a polar covalent bond and non–polar covalent bond?
Ans.
A covalent bond between two dissimilar atoms in which the shared electron pair is not attracted equally by the two atoms and the bonded atoms acquire a partial positive and negative charge is called polar covalent bond.
A covalent bond between two like atoms such as H–H,
Cl–Cl, in which shared electron pair is attracted equally by both the atoms is called a non–polar covalent bond.
Q.6 Indicate the hybridization and bond angles associated with each of the following (a) linear (b) tetrahedral (c) trigonal planar.
Ans.
(a) Sp. 180o (b) Sp3. 109o (c) Sp2. 120o
Q.7 What are the similarities and differences between atomic orbitals and molecular orbitals?
Ans.
Both atomic and molecular orbitals have a characteristic energy and shape each can hold a maximum of the two electrons. Atomic orbitals are localized and their energies are the result of interactions between the subatomic particles in a single atom. Molecular orbitals can be delocalized and their energies are influenced by interactions between electrons on several atoms.
Q.8 Why is the bonding molecular orbital of H2 at lower energy than the electron in a hydrogen atom?
Ans.
There is a net lowering in energy that accompanies bond formation because the electrons in H2are strongly attracted to both H nuclei, while in H–atom the electron is attracted by only one nucleus.
Q.9 How many electrons can be placed into each molecular orbital of a molecule?
Ans.
Two electrons can be placed into each molecular orbital of a molecule.
Q.10 What is meant by an ionic bond?
Ans.
An ionic bond is a chemical bond formed by the electzostatic attraction between positive and negative ions. The bond formed between two atoms when one or more electrons are transformed from valence shell of one atom to the valence shell of the other. The atom that loses electrons becomes cation and the atom that gains electrons becomes anion. The electronic configurations of both ions are those of noble gas atoms. All metals react with non–metals to form ionic compounds.
Na(g) ® Na+ + e–
Cl(g) + e– ® Cl–(g)
Na+(g) + Cl–(g) ® NaCl
Q.11 Can a molecule have polar bonds and not a dipole?
Ans.
Yes a molecule can have polar bonds and not a dipole, if the orientation of the polar bonds in the molecule cancel each other’s effect e.g. CO2.
Q.12 Can a molecule have non–polar bonds only and have a dipole?
Ans.
No. A molecule which have non–polar bonds only cannot have a dipole.
Q.13 State the difference between a polar bond and a polar molecule.
Ans.
A polar bond is a covalent bond that exists between two atoms having an electronegativity difference greater than 0.2 a polar molecule results, if one or more polar bonds in a molecule is not balanced by other polar bonds in the
Q.14 According to molecular orbital theory, would Be2 be expected to exist? Explain.
Ans.
Be2 is not expected to exist, it has bond order of zero and is not energetically favoured over isolated Be atoms.
Q.15 What is meant by bond order?
Ans.
The bond order is half the difference between the number of bonding electrons and the number of antibonding electrons.
Bond order =
(No. of bonding electrons – No. of antibonding electrons)
Q.16 What is meant by paramagnetic substance?
Ans.
A paramagnetic substance is a substance that is attracted by a magnetic field and this attraction is generally the result of unpaired electrons. The more unpaired electrons in a species (substance), the stronger the force of attraction. This type of magnetic behaviour is called paramagnetism.
Q.17 What is meant by a diamagnetic substance?
Ans.
A diamagnetic substance is a substance that is not attracted by a magnetic field. This property is called diamagnetism. This property generally means that substance has only paired electrons.
Q.18 (a) what is hybridization at carbon atom in CH4, C2H4 and C2H2.
(b) The carbon atom in CH4 cannot participate in multiple bonding, whereas that is C2H4 can. Explain.
Ans.
(a) The hybridization at the carbon atom in CH4 is Sp3, in C2H4 is sp2 and in C2H2 is sp.
(b) The C atom in CH4 is sp3 hybridized, there are no un–hybridized p orbitals, available for the poverlap required for multiple bonds. In C2H4, the C atom is sp2 hybridized with two p atomic orbitals (one on each C atom), available to form the p overlap in the C = C double bond.
Q.19 Describe briefly the VSEPR theory.
Ans.
The VSEPR theory predicts the shape of molecules and ions in which valence shell electron pairs are arranged around the central atom of a molecule or ions in such a way that there is maximum separation so that electron repulsions are minimized and electron nucleus attractions are maximized. Some of these electron pairs are bonding and some are lone pairs. The direction in space of the bonding pairs gives the molecular geometry. A lone pair of electrons occupy more space than a bonding pair. Repulsive forces decrease sharply with increasing interpair angle. They are strong at 90o much weaker at 120o and very weak at 180o. in VSEPR model, each multiple bond is treated as though it were a single electron pair.
Q.20 A lone pair of electrons occupies more space than a bond pair?
Ans:
A lone pair of electrons occupies more space than bond pair because lone pair is attracted by only one nucleus while bond pair is attracted by two nuclei. Due to less nuclear attraction to lone pair its electronic charge is spread ot more in space than that of bond pair.
Q.21 Predict the geometry of (a) BeCl2 (b) BF3 (c) SiH4 molecules.
Ans.
(a) The two bond pairs of electrons in BeCl2 molecule arrange themselves as far apart as possible to minimize the repulsion between them. The only arrangement which can satisfy this condition is linear i.e. at angle of 180o.
: Cl ––– –– Cl
(b) The BF3 molecule containing three bond pairs of electrons is trigonal planar because, this structure gives maximum separation among the three bonding electron pairs.
::: :
(c) In SiH4 molecule, the electrostatic repulsion between four bonding electron pairs will be minimum, when they are present at corners of a regular tetrahedron making angle 109.5o with each other.
Q.22 Describe briefly the valence bond theory.
Ans.
According to valence bond theory, a covalent bond is formed by pairing of electrons by the overlap of half (partially) filled atomic orbitals of two atoms. The two overlapping orbitals must be valence orbitals, must be half filled and must retain their identities. By overlap means that the electrons of overlapping orbitals share a common region of high electron density, along the line between two nuclei called bond axis.
This theory explains the bonding in terms of overlapping of atomic orbitals and mixing of atomic orbitals called hybridization. Multiple bond occurs via the overlap of atomic orbitals to give sbonds and p bonds.
Q.23 Why monoatomic cations smaller than their corresponding neutral atoms?
Ans.
Electrostatic repulsions are reduced by removing an electron from a neutral atom, effective neutral charge increases and therefore the cation is smaller than their corresponding neutral atoms.
Q.24 Why are monoatomic anions larger than their corresponding neutral atoms?
Ans.
The additional electrostatic repulsion produced by adding an electron to a neutral atom decreases the effective nuclear charge experienced by the valence electrons and increases the size of anions.
Q.25 Why does the size of ions increase as one produced down a column in the periodic table?
Ans.
Going down a column, valence electrons are further from the nucleus and they experience greater shielding by core electrons. The greater radial extent of the valence electrons outweigh the increase in atomic number. Therefore the size of ions increase as one proceeds down a column.
Q.26 What is an isoelectronic series?
Ans.
An isoelectronic series is a group of atoms or ions that have the same number of electrons, and thus the same electronic configuration.
Q.27 Why noble gases are most stable?
Ans.
Noble gases are most stable because their s and p orbitals are completely filled.
Q.28 Why O2 molecule is paramagnetic in nature?
Ans.
O2 molecule is paramagnetic in nature due to presence of unpaired electrons in its molecule.
Q.29 Why CO2 is non–polar molecule although C–O bond is polar?
Ans.
Each C–O bond in CO2 is polar. The two bond dipoles in CO2 are equal in magnitude and are exactly opposite in direction. The bond dipoles cancel each other. Therefore, the overall dipole moment of CO2 is zero. Thus CO2 is a non–polar molecule.
Od– = C = Od+
Q.30 Why H2O is a polar molecule?
Ans.
H2O is a bent molecule with two polar bonds. Both the bonds are identical, so the bond dipoles are equal in magnitude. Because the molecule is bent, however the bond dipoles do not directly oppose to each other and therefore, do not cancel each other.
Hence the H2O molecule has an overall dipole moment
(m = 1.85D), because H2O has dipole moment, it is polar molecule.
Q.31 The melting and boiling points of electrovalent compounds are very high as compared with those of covalent compounds. Explain.
Ans.
The melting and boiling points of electrovalent compounds are very high, because the ions are tightly packed in the crystal lattice by strong attractive forces and high thermal energy is required to separate them from one another. The atoms of molecules in covalent compounds are held together by weak intermolecular forces and less energy is required to separate the atoms or molecules in a solid or liquid.
Q.32 (a) Why solid NaCl does not conduct electricity?
(b) What will happen if electric currents is passed through molten NaCl or its aqueous solution? Explain.
Ans.
(a) Solid NaCl does not conduct electricity because the oppositely charged ions are held together by strong electrostatic forces and the ions do not free to move.
(b) However, in the molten NaCl or its aqueous solution the ions because quite free to migrate in an electric field and conduct electricity when passed through them.
Q.33 What is molecular orbital theory?
Ans.
According to molecular orbital theory, atomic orbitals overlap to form molecular orbitals n atomic orbitals combine to form n molecular orbitals. Half of them are bonding molecular orbitals and half antibonding molecular orbitals. In this combination, the individual atomic orbital character is lost in order to form an entirely new orbital that belongs to whole molecule. The theory successfully explains bond order and paramagnetic property of O2.
Q.34 In many cases, the distinction between a coordinate covalent and a covalent bond vanishes after bond formation. Explain with the help of an example.
Ans.
A coordinate covalent bond is not essentially different from other covalent bonds, it involves the sharing of pair of electrons between two atoms. An example is formation of NH4+ ion in which all bonds are identical, so the distinction between a coordinate covalent bond and covalent bond vanishes after bond formation.
H+ + :NH3 ® H – ® H+
or
Q.35 PF3 is a polar molecule with dipole moment 1.02 D and thus the P–F bond is polar. Si, being in proximity of P in the periodic table, it is expected that Si–F bond would also be polar, but SiF4 has no dipole moment. Explain why it is so?
Ans.
PF3 has one lone pair of electrons and one P atom lies out of plane of rest of the atom and thus PF3 is pyramidal molecule and has 1.02 D dipole moment and thus the P–F bond is polar.
SiF4 molecular has tetrahedral shape and the SiF bonds are directed by Symmetrically about the central Si atom. Although Si–F bonds are polar, but all the four bond moments cancel out one another and give SiF4 molecule of zero dipole moment.
Q.36 NaCl is a harder substance at room temperature than glucose explain.
Ans.
The hardness of substance depends on the strength of the forces between the particles forming a substance. NaCl is an ionic compound and consists of Na+ and Cl– ions which are held together by strong electrostatic forces of attraction while glucose consists of molecules which are held together by weak intermolecular forces. Therefore, NaCl is a harder substance at room temperature than glucose.
Q.37 The linear of BeCl2 suggests that central Be atom is sp–hybridized. What type of hybridization a central atom undergoes when the atoms bonded to it are located at the corners of (a) an equilateral triangle and (b) a regular tetrahedron.
Ans.
(a) The central atom undergoes Sp2 hybridization when the atoms bonded to it are located at the corners of an equilateral triangle.
(b) The central atom undergoes Sp3 hybridization when the atoms bounded to it are located at the corners of a regular tetrahedron.
Q.38 A double bond is shorter and stronger than a single bond.
Ans.
The greater the number of electron pairs in the bond, the shorter and stronger will be the bond because of greater nuclei–electron attractions. A double bond has two shared electron pairs, while a single bond has only one shared electron pair. Thus a double bond has a greater nuclei electron attraction, than a single bond. Therefore, a double bond is shorter and stronger than a single bond.
Q.39 NH3 and H2O can form coordinate covalent bond with H+ but CH4 cannot do so.
Ans.
NH3 and H2O both have lone pair of electrons on N and O atoms which can donate to a H+(electron deficient) to from a coordinate covalent bond. In CH4 there is no lone pair of electrons to donate H+ for the formation of coordinate covalent bond. Therefore, NH3 and H2O can form coordinate covalent bond with H+ but CH4 cannot do so.
H3N: + H+ ® [H3N ® H]+
H2O + H+ ® H3O ® [H2O+ ® H]+
Q.40 Covalent bond may be non–polar but coordinate covalent bond is always polar.
Ans.
A covalent bond between two like atoms is always non–polar, whereas between two unlike atoms, it is a polar. In coordinate covalent bond the shared electron pair is denoted by only one of the two bonded atoms. The atom which denotes the electron pair acquires partial positive charge and the atom which accepts the electron pair acquires partial negative charge. Therefore, coordinate covalent bond is always polar. Hence a covalent bond may be non–polar, but coordinate covalent bond is always polar.
Q.41 Molecule of O2 is paramagnetic in nature. Explain.
Ans.
A substance is paramagnetic, when it has unpaired electrons.
According to molecular orbital theory, O2 has two unpaired electrons in the degenerate orbitalsp2 py and p*2 px. Due to the presence of these two unpaired electrons O2 molecule is paramagnetic.
Q.42 Dipole moment of CO2 is zero but that of CO is o.12 Debye.Why?
Ans:
CO2 is a linear molecule and the tow diploes cancel the effect of each other. In CO there is a single dipole directed from carbon to oxygen and it not cancelled.
Q.43 Why dipole moment of benzene is zero?
Ans:
Benzene is symmetrical planar molecule. It has six C – H there are six dipole moments. All the dipole moments cancel the effect of each other and net value is zero.
Q.44 I.E is index of metallic character why?
Ans:
Metals have loosely held electrons which are delocalized and are responsible for the properties of metals. So, metals have low ionization energies.
Q.45 The abnormality of the bond length and bond strength in HI is less prominent than that of HCl, give reason?
Ans:
Chlorine has higher electronegative than iodine. So, the polarities of HCI and HI bonds are unequal. Therefore, abnormality of bond length and bond strength of HCI is more prominent than HI.
Q.46 How does electro negativity difference decide the nature of ionic bond?
Ans:
When the electro negativity difference between two bonded atoms is 1.7 or move than that, then the bond is said to be ionic, otherwise, covalent. The % age of ionic character is more the 51% when the electro negativity difference is 1.7.
Q.47 Why NH3 and PH3 give coordinate covalent bond?
Ans:
NH3 and PH3 have lone pairs of electrons, which can be donated to H+ to make a coordinate covalent bond. In this way, NH4+ and PH4+ are produced which have perfect tetrahedral structure and all the four bonds have perfectly equal status.
Q.48 Most of the elements of the periodic table attain the electronic configuration of inert gases during bond formation. Justify it
Ans:
Inert gases are not reactive due to complete octet except He, Most of the S- and P- block elements may attain eight electrons in the outermost orbitals they do so either by losing gaining or sharing the electrons.
Q.49 Define octet rule?
Ans:
The tendency of the atoms to attain a maximum of eight electrons in the valence shell is called octet rule.
Q.50 What is an ionic radius?
Ans:
The ionic radius of an ion is the radius of the ion while considering it to be spherical in shape.
Q.51 What is a covalent radius?
The covalent radius of an element is defined as half of the single bond length between tow similar atoms covalently bonded in a molecule.
Q.52 What is difference between sigma & pi bond?
Ans:
SIGMA BOND:
The bond which is formed by the head to head overlapping called sigma bond. The electron density is present between two nuclei.
Pi BOND:
The bond which is formed by the sideways overlapping of two half filled orbitals. The electron density is present above and below the line joining the two nuclei.
Q.53 Why the size of an atom can not be measured directly?
Ans:
The size of an atom can not be measured directly due following reasons:
(i)There is no sharp boundary of an atom. The probability of finding an electron never becomes zero even at larger distances from the nucleus.
(ii) The electronic probability distribution is affected by neighbouring atoms. For this reason the size of an atom may change from one compound to another.
Q.54 Why E.A of Flourine is less than the expected value?
Ans
Since the size of fluorine is very small when electron is added in the fluorine it is strongly repelled by the already existing electrons. An extra amount of energy is provided to add an electron therefore its electron affinity is less than the expected value.
Q.55 Why sigma bond is stronger than pi bond? Or Why pi bond is more diffused than sigma bond?
Ans:
Sigma bond is more diffused than pi bond due to the linear overlapping of orbitals. Moreover electron density is present between two nuclei which is strongly attracted by two nuclei. While in pi bond electron density is not strongly attracted by two nuclei therefore it is weak than pi bond
Q.56 Define bond length.
Ans.
The distance between the nuclei of two atoms forming a covalent bond is called bond length. In general it is the sum of the covalent radii of the combined atoms.
Q.57 What is dipole moment? What are its units?
Ans.
The dipole moment may be defined as the product of electric charge (q) and distance (r) between the two oppositely charged centres. It is vector quantity as it has magnitude and direction. It plays a major role in determining the % age ionic character of a covalent bond and the shapes of molecules.
The dipole moment is measured in Debye units (D). It is denoted by symbol m.
Q.58 Define bond energy Give its units.
Ans.
The bond energy is defined as the average amount of energy required to break all bonds of particular type in one mole of substance. It is determined by measuring the heat involved in a chemical reaction.
It is also defined as the energy required to break Avogadro’s number (6.02 x 1023) of bonds or the energy released when an Avagadro number of bonds are formed. It is a measure of strength of bonds. The bond energy is measured in KJ mol–1.
Q.59 Define the following terms.
(a) Ionization energy
(b) Electron affinity
(c) electronegativity
Ans.
(a) Ionization energy:
The minimum amount of energy required to remove an electron from an atom is called ionization energy.
It depends upon the atomic size, nuclear charge and shielding effect of electrons.
(b) Electron Affinity:
The minimum amount of energy released when an electron is added to an isolated neutral gaseous atom in the lowest energy state to produce an anion is called electron affinity. It is measured in kJ mol–1.
(c) Electronegativity:
The tendency of an atom to attract shared pair of electron towards itself is called electronegativity. It is measured in electron volts.
Chapter#7 **Thermo-Chemistry*****
SHORT QUESTION AND ANSWERS
Q.1 Define the following terms and give three examples of each.
Ans.
(i) System:
The substance which is under experiment or under observation is called as system.
Examples:
(i) Pb(NO3)2 in decomposition of Pb(NO3)2.
(ii) Zn and CuSO4 solution, the reaction mixture in the vessel.
(iii) CaCO3 in thermal decomposition of CaCO3
(ii) Surroundings:
Everything around the system which is not a part of system is called surroundings.
For example
During the reaction between Zn and CuSO4 solution vessel and air etc are surroundings.
(iii) State function:
A macroscopic property of a system which has some definite value for initial and final state and independent of the path followed e.g.
(i) Pressure (ii) Temperature (iii) Internal energy.
Note: Heat is not a state function.
Q.2 Describe the units of energy.
Ans.
Mostly Joule and calorie are used for the measurement of energy.
Calorie:
The amount of energy required to raise the temperature of one gram of water from 14.5oC to 15.5oC is called one calorie.
Joule:
It is SI unit of energy and defined as energy expanded when a force of one Newton moves a body through one meter in the direction in which force is applied.
Joule = Force x distance
1J = 1 N x 1 m
Q.3 What are rhermochemical reactions?
Ans:
Exothermic reaction:
Those thermochemical reactions in which heat is evolved as a result of reaction are called as exothermic reactions.
C(s) + O2(g) ® CO2(g) DH = – 393.7 kJ/mole
H2(g) + O2(g) ® H2O(l) DH = – 285.5 kJ/mole
N2 + 3H2(g) ® 2NH3(g) DH = – 41.6 kJ/mole
Endothermic reactions:
Those thermochemical reactions in which heat is absorbed as a result of reaction are called as endothermic reaction.
N2(g) + O2(g) ® 2NO(g) DH = + 180.51 kJ/mole
H2O(l) ® H2(g) + O2(g) DH = + 285.58 kJ/mole
H2(g) + I2(g) ® 2H I DH = + 52.96 kJ/mole
Q.4 Differentiate between internal energy and enthalpy of a system?
Ans:
Internal energy:
The total of all kinds of K.E and P.E of all the particles of a system is called as internal energy. It is denoted by “E’ e.g., kinetic energy may be in the form of translation, vibrational and rotational motion and potential energy is intermolecular and intramolecular forces of attraction.It is a sate function of system.
E= K.E + P.E
Enthalpy of the system:
The total heat contents of a system and denoted by H. The increase in the internal energy of a system plus work done is called as enthalpy i.e.
H = E + Pv
Q.5 Define the followings:
(i) Enthalpy of reaction
(ii) Enthalpy of neutralization
(iii) Enthalpy of combustion
Ans:
Standard Enthalpy of reaction:
The enthalpy change when no. of moles of reactants as indicated by the balanced chemical equation react completely together to give the products under the standard conditions.
H2(g) + O2(g) ® H2O(l) DH = – 285.5 kJ/m
N2(g) + O2(g) ® 2NO(g) DH = + 180.5 kJ/m
Standard Enthalpy of Combustion:
The enthalpy change when one mole of a substance is completely burnt in excess of oxygen under standard conditions.
` C2H5OH(l) + 3O2(g) ® 2CO2(g) + 3H2O(l) DH
= – 1368 kJ/m
C(s) + O2(g) ® CO2(g) DH = – 393.7 kJ/m
2Al(s) + O2(g) ® Al2O3(s) DH = – 1675.7 kJ/m
Standard enthalpy of atomization:
The enthalpy change when one mole of gaseous atoms are formed from the elements under the standard conditions is called enthalpy of atomization.
® H(g) DH = 218 kJ/mole
Cl2 ® Cl(g) DH = + 121 kJ/mole
Q.6 Why it is essential to mention the physical states of reactants and products in a thermochemical equation?
Ans.
The heat of reaction depends upon the physical states of the reactants and products, heat of reaction is different in different physical states therefore, while writing a thermochemical equation it is essential to mention the physical states of the reactants and products.e.g
H2(g) + 1/2O2(g) → H2O(g) ∆H = -241.5 KJ mol-1
H2(g) + 1/2O2(g) → H2O(l) ∆ H= -285.8 KJ mol-1
Q.7 Differentiate between spontaneous and non–spontaneous reaction?
Ans.
The process which takes place on its own without any outside help and moves from a non–equilibrium state to equilibrium state is called spontaneous or natural process. It is real, unidirectional and irreversible e.g. water flows from higher level to low leve, reaction between acid and base etc.
There are certain reactions which need energy to start but once they start they proceed their own for example burning of candle.
The process which does not take place on its own and does not occur in nature is called as non–spontaneous. It is reverse of the spontaneous process i.e. pumping of water uphill, flow of heat from colder to hotter region etc.
Q.8 Prove that change in enthalpy is equal to heat of reaction? / prove that qp =DH?
Ans.
We know that enthalpy is equal to the internal energy plus product of pressure and volume.
H = E + Pv
According to first law of the thermodynamics
q = DE + w
At constant pressure w = PDv
qp = DE + PDv
qp = (E2 – E1) + (v2 – v1)p
qp = E2 – E1 + Pv2 – Pv1
qp = (E2 + Pv2) – (E1 + Pv1) H2 = E2 + Pv2
qp =H2 – H1 H1 = E1 + Pv1
qp =DH
This shows that enthalpy change is equal to amount of heat absorbed at constant pressure.
Q.9 Briefly explain laws of thermochemistry.
Ans. First law of thermochemistry:
The enthalpy of formation of a compound to the enthalpy of decomposition of that compound. e.g.
H2(g) + O2(g) ® H2O(l) DH = – 285.5 kJ/mole
H2O(l) ® H2(g) + O2(g) DH = + 285.5 kJ/mole
Second law of thermochemistry (Hess’s law):
The amount of heat evolved or absorbed in a chemical reaction is same whether the reaction takes place in one or several steps. e.g. single step process.
C(s) + O2(g) ® CO2(g) DH1 = – 393.7 kJ
Two steps process:
C(s) + O2(g) ® CO(g) DH2 = – 110.7kJ
DH1 = DH2 + DH3
– 393.7 = – 110.7 – 283
– 393.7 = – 393.7
Q.10 Draw a complete, fully labeled Born–Haber cycle for the formation of KBr.
Reactions:
K(s) + Br2(l) ® K+ Br– D H/kJ mole–1
– 392 kJ mole–1
K(s) ® K(g) + 90 kJ mole–1
K(g) ® K+ e– + 420 kJ mole–1
Br2(l) ® Br(g) +112 kJ mole–1
Br(q) + e– ® Br–(g) – 342 kJ mole–1
Solution:
The heat of formation of KBr is equal to sum of all the enthalpies.
DHf = – 392 kJ mole–1 (Heat of formation)
DHs = + 90 kJ mole–1 (Heat of sublimation)
DHi = + 420 kJ mole–1 (Ionization pot of k)
DHD/2 = + 112 kJ mole–1 (Dissociation energy of Br2)
DHe = – 342 kJ mole (electron affinity)
DHl = ? (Lattice energy of kBr)
DHf = DHat + DHi + DHd + DHe + DHl
– 392 = + 90 + 420 + 112 + (– 342) + DHl
DHl = – 672 kJ mole–1
Q.11 Heat is evolved in exothermic reactions and absorbed in endothermic relations.
Ans.
When bond formation energy is greater than the bond breaking energy then the excess of energy is evolved making the reaction exothermic. When the bond breaking energy is greater than the bond formation energy then the difference of energy is supplied from surrounding making the reaction is endothermic.
Q.12 How would you explain that change in enthalpy is a state function?
Ans.
As H = E + PV
As E, P and v are state functions as they are independent of path and depend only on the initial and final state of the system therefore enthalpy of a system is also a state function because enthalpy depends on E, P and V.
Q.13 How can you prove that (w = – PDV)
Ans.
We know that
Work = Force x distance
Initial volume= V1
Final volume = V2
Change in volume DV = V2 – V1
In the expansion of gases, the work is expressed in terms of pressure and change in volume. So,
Work = P x DV
= w = – PDV
The work is negative because it is done by the system.
Q14. What is difference between heat and temperature? Write a mathematical relationship between these two parameters?
Ans:
Heat: The measure of total energy of a substance is called heat. It is property of a body which flows from a body at higher temperature to a body at lower temperature. It is denoted by ‘q’. It depends upon the quantity of a substance. It is measured by calorimeter. It is not a state function.
Temperature: It is measure of average K.E of the molecules in the system. It is denoted by ‘T’. It is independent of the quantity of a substance. It is measured by thermometer. It is a state function.
Realtionship: q= m x S x ∆T
Q15. What is enthalpy of neutralization? Why enthalpy of neutralization of strong acid and base is always -57.4 KJ mol-1?
Ans:
It is the amount of heat eveolved or absorbed when one mole of H+ ions from an acid reacts with one mole of OH- from a base to form one mole of H2O. Under standard conditions it is called standard enthalpy of neutralization, and it is denoted by ∆Hn.
H+ + OH- → H2O ∆Hn = -57.4 KJ mol-1
The heat of neutralization of strong acid or base is always -57.4 KJ mol-1 because strong acid or base is completely ionized and when acid and base is mixed no bond has to be broken
Q.1 Define the following terms and give three examples of each.
Ans.
(i) System:
The substance which is under experiment or under observation is called as system.
Examples:
(i) Pb(NO3)2 in decomposition of Pb(NO3)2.
(ii) Zn and CuSO4 solution, the reaction mixture in the vessel.
(iii) CaCO3 in thermal decomposition of CaCO3
(ii) Surroundings:
Everything around the system which is not a part of system is called surroundings.
For example
During the reaction between Zn and CuSO4 solution vessel and air etc are surroundings.
(iii) State function:
A macroscopic property of a system which has some definite value for initial and final state and independent of the path followed e.g.
(i) Pressure (ii) Temperature (iii) Internal energy.
Note: Heat is not a state function.
Q.2 Describe the units of energy.
Ans.
Mostly Joule and calorie are used for the measurement of energy.
Calorie:
The amount of energy required to raise the temperature of one gram of water from 14.5oC to 15.5oC is called one calorie.
Joule:
It is SI unit of energy and defined as energy expanded when a force of one Newton moves a body through one meter in the direction in which force is applied.
Joule = Force x distance
1J = 1 N x 1 m
Q.3 What are rhermochemical reactions?
Ans:
Exothermic reaction:
Those thermochemical reactions in which heat is evolved as a result of reaction are called as exothermic reactions.
C(s) + O2(g) ® CO2(g) DH = – 393.7 kJ/mole
H2(g) + O2(g) ® H2O(l) DH = – 285.5 kJ/mole
N2 + 3H2(g) ® 2NH3(g) DH = – 41.6 kJ/mole
Endothermic reactions:
Those thermochemical reactions in which heat is absorbed as a result of reaction are called as endothermic reaction.
N2(g) + O2(g) ® 2NO(g) DH = + 180.51 kJ/mole
H2O(l) ® H2(g) + O2(g) DH = + 285.58 kJ/mole
H2(g) + I2(g) ® 2H I DH = + 52.96 kJ/mole
Q.4 Differentiate between internal energy and enthalpy of a system?
Ans:
Internal energy:
The total of all kinds of K.E and P.E of all the particles of a system is called as internal energy. It is denoted by “E’ e.g., kinetic energy may be in the form of translation, vibrational and rotational motion and potential energy is intermolecular and intramolecular forces of attraction.It is a sate function of system.
E= K.E + P.E
Enthalpy of the system:
The total heat contents of a system and denoted by H. The increase in the internal energy of a system plus work done is called as enthalpy i.e.
H = E + Pv
Q.5 Define the followings:
(i) Enthalpy of reaction
(ii) Enthalpy of neutralization
(iii) Enthalpy of combustion
Ans:
Standard Enthalpy of reaction:
The enthalpy change when no. of moles of reactants as indicated by the balanced chemical equation react completely together to give the products under the standard conditions.
H2(g) + O2(g) ® H2O(l) DH = – 285.5 kJ/m
N2(g) + O2(g) ® 2NO(g) DH = + 180.5 kJ/m
Standard Enthalpy of Combustion:
The enthalpy change when one mole of a substance is completely burnt in excess of oxygen under standard conditions.
` C2H5OH(l) + 3O2(g) ® 2CO2(g) + 3H2O(l) DH
= – 1368 kJ/m
C(s) + O2(g) ® CO2(g) DH = – 393.7 kJ/m
2Al(s) + O2(g) ® Al2O3(s) DH = – 1675.7 kJ/m
Standard enthalpy of atomization:
The enthalpy change when one mole of gaseous atoms are formed from the elements under the standard conditions is called enthalpy of atomization.
® H(g) DH = 218 kJ/mole
Cl2 ® Cl(g) DH = + 121 kJ/mole
Q.6 Why it is essential to mention the physical states of reactants and products in a thermochemical equation?
Ans.
The heat of reaction depends upon the physical states of the reactants and products, heat of reaction is different in different physical states therefore, while writing a thermochemical equation it is essential to mention the physical states of the reactants and products.e.g
H2(g) + 1/2O2(g) → H2O(g) ∆H = -241.5 KJ mol-1
H2(g) + 1/2O2(g) → H2O(l) ∆ H= -285.8 KJ mol-1
Q.7 Differentiate between spontaneous and non–spontaneous reaction?
Ans.
The process which takes place on its own without any outside help and moves from a non–equilibrium state to equilibrium state is called spontaneous or natural process. It is real, unidirectional and irreversible e.g. water flows from higher level to low leve, reaction between acid and base etc.
There are certain reactions which need energy to start but once they start they proceed their own for example burning of candle.
The process which does not take place on its own and does not occur in nature is called as non–spontaneous. It is reverse of the spontaneous process i.e. pumping of water uphill, flow of heat from colder to hotter region etc.
Q.8 Prove that change in enthalpy is equal to heat of reaction? / prove that qp =DH?
Ans.
We know that enthalpy is equal to the internal energy plus product of pressure and volume.
H = E + Pv
According to first law of the thermodynamics
q = DE + w
At constant pressure w = PDv
qp = DE + PDv
qp = (E2 – E1) + (v2 – v1)p
qp = E2 – E1 + Pv2 – Pv1
qp = (E2 + Pv2) – (E1 + Pv1) H2 = E2 + Pv2
qp =H2 – H1 H1 = E1 + Pv1
qp =DH
This shows that enthalpy change is equal to amount of heat absorbed at constant pressure.
Q.9 Briefly explain laws of thermochemistry.
Ans. First law of thermochemistry:
The enthalpy of formation of a compound to the enthalpy of decomposition of that compound. e.g.
H2(g) + O2(g) ® H2O(l) DH = – 285.5 kJ/mole
H2O(l) ® H2(g) + O2(g) DH = + 285.5 kJ/mole
Second law of thermochemistry (Hess’s law):
The amount of heat evolved or absorbed in a chemical reaction is same whether the reaction takes place in one or several steps. e.g. single step process.
C(s) + O2(g) ® CO2(g) DH1 = – 393.7 kJ
Two steps process:
C(s) + O2(g) ® CO(g) DH2 = – 110.7kJ
DH1 = DH2 + DH3
– 393.7 = – 110.7 – 283
– 393.7 = – 393.7
Q.10 Draw a complete, fully labeled Born–Haber cycle for the formation of KBr.
Reactions:
K(s) + Br2(l) ® K+ Br– D H/kJ mole–1
– 392 kJ mole–1
K(s) ® K(g) + 90 kJ mole–1
K(g) ® K+ e– + 420 kJ mole–1
Br2(l) ® Br(g) +112 kJ mole–1
Br(q) + e– ® Br–(g) – 342 kJ mole–1
Solution:
The heat of formation of KBr is equal to sum of all the enthalpies.
DHf = – 392 kJ mole–1 (Heat of formation)
DHs = + 90 kJ mole–1 (Heat of sublimation)
DHi = + 420 kJ mole–1 (Ionization pot of k)
DHD/2 = + 112 kJ mole–1 (Dissociation energy of Br2)
DHe = – 342 kJ mole (electron affinity)
DHl = ? (Lattice energy of kBr)
DHf = DHat + DHi + DHd + DHe + DHl
– 392 = + 90 + 420 + 112 + (– 342) + DHl
DHl = – 672 kJ mole–1
Q.11 Heat is evolved in exothermic reactions and absorbed in endothermic relations.
Ans.
When bond formation energy is greater than the bond breaking energy then the excess of energy is evolved making the reaction exothermic. When the bond breaking energy is greater than the bond formation energy then the difference of energy is supplied from surrounding making the reaction is endothermic.
Q.12 How would you explain that change in enthalpy is a state function?
Ans.
As H = E + PV
As E, P and v are state functions as they are independent of path and depend only on the initial and final state of the system therefore enthalpy of a system is also a state function because enthalpy depends on E, P and V.
Q.13 How can you prove that (w = – PDV)
Ans.
We know that
Work = Force x distance
Initial volume= V1
Final volume = V2
Change in volume DV = V2 – V1
In the expansion of gases, the work is expressed in terms of pressure and change in volume. So,
Work = P x DV
= w = – PDV
The work is negative because it is done by the system.
Q14. What is difference between heat and temperature? Write a mathematical relationship between these two parameters?
Ans:
Heat: The measure of total energy of a substance is called heat. It is property of a body which flows from a body at higher temperature to a body at lower temperature. It is denoted by ‘q’. It depends upon the quantity of a substance. It is measured by calorimeter. It is not a state function.
Temperature: It is measure of average K.E of the molecules in the system. It is denoted by ‘T’. It is independent of the quantity of a substance. It is measured by thermometer. It is a state function.
Realtionship: q= m x S x ∆T
Q15. What is enthalpy of neutralization? Why enthalpy of neutralization of strong acid and base is always -57.4 KJ mol-1?
Ans:
It is the amount of heat eveolved or absorbed when one mole of H+ ions from an acid reacts with one mole of OH- from a base to form one mole of H2O. Under standard conditions it is called standard enthalpy of neutralization, and it is denoted by ∆Hn.
H+ + OH- → H2O ∆Hn = -57.4 KJ mol-1
The heat of neutralization of strong acid or base is always -57.4 KJ mol-1 because strong acid or base is completely ionized and when acid and base is mixed no bond has to be broken
Chapter#8 *CHEMICAL EQUILIBRIUM*
SHORT QUESTION WITH ANSWERS
Q.1 What is weak electrolyte?
Ans.
A compound which is only partially ionized in aqueous solution is called as weak electrolyte.e.g CH3COOH(Acetic aci)
Q.2 What is meant by state of chemical equilibrium?
Ans.
The state of reversible reaction, in which forward and reverse rates are equal.(Rf=Rr)
Q.3 What are reversible reaction?
Ans.
Those reactions in which the reactants products are converted into each other under same set of conditions.
Q.4 Define Le–Chatlier’s principle.
Ans.
If a system in equilibrium is disturbed at equilibrium it will move in that direction where change is minimized.
Q.5 What are conjugate acid and bases?
Ans.
When an acid is dissolved in water it will form H3O+ ions and an anion the anion of an acid is called conjugate base and H3O+ is called conjugate acid e.g.
CH3OOH + H2O H3O+ + CH3COO–
Acid Base Conjugate acid Conjugate base
Q.6 Why we need buffer solution in daily life?
Ans.
Buffers can resist the charge of its PH value therefore they are required in chemical analysis, pharmaceuticals, electroplating, photography, beverage industry, microbiology molecular biology, soil science and in quantitative analysis.Our blood is best example of buffer solution having PH 7.35. If it decreases upto 7 or goes upto 8 death may occur.
Q.7 Discuss the factors on which buffer PH depends.
Ans.
PH of buffer depends upon two factors.
(i) Pka or ka value of an acid
(ii) Concentration of salt and acid taken for the buffer.
Q.8 What is buffer capacity?
Ans.
The amount of acid or base which a buffer can absorb without significant charge in PH is called buffer capacity. Buffer capacity is the ability of buffer to resist the change in its PH value.
Q.9 How can we calculate the PH of buffer?
Ans.
We can calculate the PH of a buffer the PH of Henderson equation which can be derived as
HA ® H+ + A–
Kc =
H+ =
Take log on both sides
Log [H+] = log
Log [H+] = log ka +
Multiply equation by (– 1) on both sides
– log [H+] = – log ka – log
PH = Pka – log
or
PH = Pka + log
PH = Pka + log
In similar way
POH = Pkb + log
Q.10 How Kc can be applied to calculate direction of chemical reaction?
Ans.
(i) If < Kc
Then the reaction will proceed in the forward direction until the equilibrium is established.
(ii) If > Kc
Then the reaction will proceed in the backward direction until the equilibrium is established.
If = kc
Then the reaction is already in equilibrium.
Q.11 How value of Kc can help to predict the extent to which a chemical reaction can take place?
Ans.
The extent of reaction depends upon the magnitude of Kc. i.e.
(i) Very small value of Kc:
when Kc is very small the forward reaction will not occur to an appreciable extent and the reverse reaction will go almost to complete.
(ii) Kc is very large:
When Kc is very large the reverse reaction will not occur to an appreciable extent and the forward reaction is almost complete.
(iii) Kc has an intermediate value:
When the value of Kc is neither very large nor very small, the equilibrium mixture contains appreciable amounts of both reactant and products.
Q.12 What are optimum conditions for the formation of NH3 by Haber’s process?
N2(g) + 3H2(g) 2NH3(g) DH = -92.46 kJ/m
Ans.
(i) As the reaction is exothermic max yield is obtained at low temp.
(ii) By applying pressure, more product is obtained. According four moles of reactants produce two moles of product so by applying pressure more yield of NH3 is gained.
(iii) Continuous withdrawl of NH3 will also increase the yield.
(vi) Rate of reaction can also be increase with the help of catalyst like iron or iron oxide.
Q.13 A catalyst does not change the position of equilibrium but this equilibrium position approach earlier. Why?
Ans.
A catalyst decrease the energy of activation required by reacting substance so the position of equilibrium reaches earlier.
Q.14 How the change in pressure at equilibrium position effect the following reaction?
PCl5 PCl3 + Cl2
Ans.
According to Le–chatlier’s principle, by increasing pressure the reaction will move toward less number of moles. In the above reaction, by increasing pressure the reaction will move in backward direction.
Q.15 What is the relationship b/w kc, kp kc & kn?
Ans.
These four parameters are related as.
Kp = Kc (RT)Dn = Kx (P)Dn = Kn
When mole of products are reactant are equal i.e. Dn = 0 their
Kp = kc = dx = kn
Q.16 Prove that Kw is ionic product of water and its value is 1 x 10–14 at 25o.
Ans.
Water undergoes self ionization as
2H2O ® H3O+ + OH–
or H2O ® H+ + OH–
Kc [H2O] = [H+] [OH–]
Kw = [H+] [OH–]
At 25oC the value of Kw ionic product have been measured. Its value is 1 x 10–14.
Q.17 How can you purify NaCl by common ion effect?
Ans.
The impurities present in the common salt are CaCl2 MgCl2 and Na2SO4 in it. These impurities can be removed by common Ion effect. For this purpose a saturated solution of NaCl is prepared. Then HCl gas is passed through it. HCl is stronger electrolyte than NaCl Ionization of NaCl is suppressed by passing HCl gas and it is precipitated out.
Q.18 A buffer consists of a weak acid and is salt.
Ans.
(i) CH3COOH + CH3COONa
(ii) C6H5COOH + C6H5COONa
(iii) H2CO3 + NaHCO3
(iv) H3PO3 + NaH2PO4
Q.19 How can you co–relate Ka with Kc?
Ans.
Ka is dissociation constant of an acid while Kc is the equilibrium constant suppose an acid HA Ionize as
HA + H2O ® H3O+ + A–
Kc =
concentration of [H2O] almost remains constant so we can write.
Kc =
Kc [H2O] = ka
Ka =
Q.20 How can you calculate % Ionization of an acid?
Ans.
Percentage Ionization of an acid can be determined by the following formula.
% Ionization = x 100
initially available.
Q.21 N2(g) + 3H2(g) 2NH3(g) DH = -92.46 KJ/mol
This reaction is favourable at low temperature then why its production is favourable at 4500C?
Ans:
This is industrial condition favourable at the industrial level because at low temperature the reaction becomes very slow because both are gases.
Q.22 What is common ion effect?
Ans:
The decrease in ionization of a weak electrolyte by adding another strong electrolyte having common ion.e.g ionization of NaCl is decreased by passing HCl through NaCl solution.
Q.23 What is solubility product?
Ans:
The product of molar solubilities of the ions of weak electrolyte at equilibrium stage is called solubility product. It is represented by Ksp
Q.1 What is weak electrolyte?
Ans.
A compound which is only partially ionized in aqueous solution is called as weak electrolyte.e.g CH3COOH(Acetic aci)
Q.2 What is meant by state of chemical equilibrium?
Ans.
The state of reversible reaction, in which forward and reverse rates are equal.(Rf=Rr)
Q.3 What are reversible reaction?
Ans.
Those reactions in which the reactants products are converted into each other under same set of conditions.
Q.4 Define Le–Chatlier’s principle.
Ans.
If a system in equilibrium is disturbed at equilibrium it will move in that direction where change is minimized.
Q.5 What are conjugate acid and bases?
Ans.
When an acid is dissolved in water it will form H3O+ ions and an anion the anion of an acid is called conjugate base and H3O+ is called conjugate acid e.g.
CH3OOH + H2O H3O+ + CH3COO–
Acid Base Conjugate acid Conjugate base
Q.6 Why we need buffer solution in daily life?
Ans.
Buffers can resist the charge of its PH value therefore they are required in chemical analysis, pharmaceuticals, electroplating, photography, beverage industry, microbiology molecular biology, soil science and in quantitative analysis.Our blood is best example of buffer solution having PH 7.35. If it decreases upto 7 or goes upto 8 death may occur.
Q.7 Discuss the factors on which buffer PH depends.
Ans.
PH of buffer depends upon two factors.
(i) Pka or ka value of an acid
(ii) Concentration of salt and acid taken for the buffer.
Q.8 What is buffer capacity?
Ans.
The amount of acid or base which a buffer can absorb without significant charge in PH is called buffer capacity. Buffer capacity is the ability of buffer to resist the change in its PH value.
Q.9 How can we calculate the PH of buffer?
Ans.
We can calculate the PH of a buffer the PH of Henderson equation which can be derived as
HA ® H+ + A–
Kc =
H+ =
Take log on both sides
Log [H+] = log
Log [H+] = log ka +
Multiply equation by (– 1) on both sides
– log [H+] = – log ka – log
PH = Pka – log
or
PH = Pka + log
PH = Pka + log
In similar way
POH = Pkb + log
Q.10 How Kc can be applied to calculate direction of chemical reaction?
Ans.
(i) If < Kc
Then the reaction will proceed in the forward direction until the equilibrium is established.
(ii) If > Kc
Then the reaction will proceed in the backward direction until the equilibrium is established.
If = kc
Then the reaction is already in equilibrium.
Q.11 How value of Kc can help to predict the extent to which a chemical reaction can take place?
Ans.
The extent of reaction depends upon the magnitude of Kc. i.e.
(i) Very small value of Kc:
when Kc is very small the forward reaction will not occur to an appreciable extent and the reverse reaction will go almost to complete.
(ii) Kc is very large:
When Kc is very large the reverse reaction will not occur to an appreciable extent and the forward reaction is almost complete.
(iii) Kc has an intermediate value:
When the value of Kc is neither very large nor very small, the equilibrium mixture contains appreciable amounts of both reactant and products.
Q.12 What are optimum conditions for the formation of NH3 by Haber’s process?
N2(g) + 3H2(g) 2NH3(g) DH = -92.46 kJ/m
Ans.
(i) As the reaction is exothermic max yield is obtained at low temp.
(ii) By applying pressure, more product is obtained. According four moles of reactants produce two moles of product so by applying pressure more yield of NH3 is gained.
(iii) Continuous withdrawl of NH3 will also increase the yield.
(vi) Rate of reaction can also be increase with the help of catalyst like iron or iron oxide.
Q.13 A catalyst does not change the position of equilibrium but this equilibrium position approach earlier. Why?
Ans.
A catalyst decrease the energy of activation required by reacting substance so the position of equilibrium reaches earlier.
Q.14 How the change in pressure at equilibrium position effect the following reaction?
PCl5 PCl3 + Cl2
Ans.
According to Le–chatlier’s principle, by increasing pressure the reaction will move toward less number of moles. In the above reaction, by increasing pressure the reaction will move in backward direction.
Q.15 What is the relationship b/w kc, kp kc & kn?
Ans.
These four parameters are related as.
Kp = Kc (RT)Dn = Kx (P)Dn = Kn
When mole of products are reactant are equal i.e. Dn = 0 their
Kp = kc = dx = kn
Q.16 Prove that Kw is ionic product of water and its value is 1 x 10–14 at 25o.
Ans.
Water undergoes self ionization as
2H2O ® H3O+ + OH–
or H2O ® H+ + OH–
Kc [H2O] = [H+] [OH–]
Kw = [H+] [OH–]
At 25oC the value of Kw ionic product have been measured. Its value is 1 x 10–14.
Q.17 How can you purify NaCl by common ion effect?
Ans.
The impurities present in the common salt are CaCl2 MgCl2 and Na2SO4 in it. These impurities can be removed by common Ion effect. For this purpose a saturated solution of NaCl is prepared. Then HCl gas is passed through it. HCl is stronger electrolyte than NaCl Ionization of NaCl is suppressed by passing HCl gas and it is precipitated out.
Q.18 A buffer consists of a weak acid and is salt.
Ans.
(i) CH3COOH + CH3COONa
(ii) C6H5COOH + C6H5COONa
(iii) H2CO3 + NaHCO3
(iv) H3PO3 + NaH2PO4
Q.19 How can you co–relate Ka with Kc?
Ans.
Ka is dissociation constant of an acid while Kc is the equilibrium constant suppose an acid HA Ionize as
HA + H2O ® H3O+ + A–
Kc =
concentration of [H2O] almost remains constant so we can write.
Kc =
Kc [H2O] = ka
Ka =
Q.20 How can you calculate % Ionization of an acid?
Ans.
Percentage Ionization of an acid can be determined by the following formula.
% Ionization = x 100
initially available.
Q.21 N2(g) + 3H2(g) 2NH3(g) DH = -92.46 KJ/mol
This reaction is favourable at low temperature then why its production is favourable at 4500C?
Ans:
This is industrial condition favourable at the industrial level because at low temperature the reaction becomes very slow because both are gases.
Q.22 What is common ion effect?
Ans:
The decrease in ionization of a weak electrolyte by adding another strong electrolyte having common ion.e.g ionization of NaCl is decreased by passing HCl through NaCl solution.
Q.23 What is solubility product?
Ans:
The product of molar solubilities of the ions of weak electrolyte at equilibrium stage is called solubility product. It is represented by Ksp
chap 9 ***SOLUTIONS**
SHORT QUESTIONS WITH ANSWER
Q.1 Binary solution can be homogenous or heterogeneous explain?
Ans.
The solutions which contain two components only are called as binary solution. If binary solution has single phase, it is called homogenous solution e.g. glucose in water. If binary solutions has more than one phase it is called heterogeneous solution e.g. oil in water.
Q.2 What is phase?
Ans.
Every sample of matter with uniform properties and fixed composition is called phase.
Q.3 Define molarity.
Ans.
Molarity is the number of moles of solute dissolved per dm3 of solution.
Q.4 What is molarity?
Ans.
Molarity is the no. of moles of solute in 1000 gm (1 kg) of the solvent.
Q.5 Explain mole fraction (x).
Ans.
The mole fraction of any component in a mixture is the ratio of the number of moles of it to the total number of moles of all the components present.
Q.6 Define PPm?
Ans.
It is defined as number of parts (by weight or volume) of a solute per million parts (by weight or volume) of the solution.
Q.7 Like dissolve like. Explain.
Ans.
The inter–Ionic forces of attraction are very strong in Ionic solids, so equally strong polar solvents are needed to dissolve them. such solids cannot be dissolved by moderately polar solvents e.g. acetone.
Q.8 What is critical solution or upper consulate temperature?
Ans.
The temperature at which two conjugate solutions merge into one another is called critical solution temperature.
Q.9 What is difference between ideal and non–ideal solution.
Ans.
IDEAL SOLUTION
NON–IDEAL SOLUTION
1.
Ideal solution obeys the Raoult’s law.
1.
Non–Ideal solution does not obey Raoult’s law.
2.
It has zero enthalpy change as their heat of solution D H = 0.
2.
It has exothermic or endothermic enthalpy change as their heat of solution D H ¹ 0
3.
Volume of solution is sum of volume of solvent and solute and change in volume is zero D v = 0.
3.
Volume of solution is not sum of the volume of solute and solvents and change in volume D V ¹ 0.
4.
Examples
(i) Benzene – ether
(ii) Chlorobenzene–bromobenzene
4.
Example
(i) Alcohol – water
(ii) Water HCl
Q.10 Define Zeotropic and Azeotropic mixture.
Ans.
The liquid mixture which boils at constant temperature and distills over without change in composition at any pressure like a chemical compound is called azeotropic mixture e.g. water HCl system, water–ethanol system. The liquid mixture which can be distilled over with a change in composition is called zeotropic mixture e.g. Benzene–ether system.
Q.11 How can you relate the dynamic equilibrium with recrystallization?
Ans.
When some solute is added to the solvent, the force of attraction between the solute particles breakdown. This process is called dissolution. Some particles of solute may combine again and converted to a solid substance. This process is called recrystalization when a saturated solution is prepared there is an equilibrium b/w dissolution and recrystallization.
Q.12 Prove that HCl form an Azeotropic mixture with water?
Ans.
HCl form an azeotropic mixture with water boiling at 110o and containing 20.24% of acid. This pair of liquids showing positive deviation has azeotropic mixture with boiling point in comparison with pure components.
Q.13 What are conjugate solutions?
Ans.
As the mutual solubilities are limited, the liquids are only partially miscible on shaking equal volumes of water and ether, two layers are formed. Each liquid layer is a saturated solution of the other liquid. Such solution are called conjugate solutions.
Q.14 State Raoult’s law in different way.
Ans. 1st statement:
The vapour pressure of a solvent above a solution is equal to the product of the vapour pressure of pure solvent and the mole fraction of solvent in solution.
2nd statement:
The lowering of vapour pressure is directly proportional to the mole fraction of solute.
3rd statement:
The relative lowering of vapour pressure is equal to the mole fraction of solute.
Q.15 The relative lowering of v.p is better than lowering of v.p why?
Ans.
The relative lowering of v.p:
(i) is independent of temp.
(ii) depends upon the cone. of solute.
(iii) is constant when equimolar proportions of different solutes are dissolved in the same mass of same solvent.
Q.16 What is positive deviation?
Ans.
If a graph is plotted b/w composition and vapour pressure of a solution the total vapour pressure curve rises to maximum, which is above the vapour pressure of either of the pure components, shows positive deviation from Raoult’s law.
Q.17 What is negative deviation?
Ans.
If the vapour pressure show a minimum. When vapour pressure of solution becomes less than either of component the solution will show negative deviation from Raoult’s law.
Q.18 Define solubility?
Ans.
It is defined as the concentration of solute in the solution when it is in equilibrium with the solid substance at a particular temperature.
Q.19 Ce2(SO4) shows exceptional behaviour toward solubility.
Ans.
Ce2 (SO4) shows exceptional behaviour whose solubility decreases with the increase in temperature becomes constant from 40oC onwards. Anyhow, it shows continuous solubility curve.
Q.20 Na2SO4 . 10H2O and CaCl2 . 6H2O show discontinuous solubility curves?
Ans.
Sometimes the solubility curves show sudden changes of solubilities and these curves are called discontinuous solubility curves. Actually these curves are combination of many curves. At the break a new solid phase appears and another solubility curve of new phase begins.
Q.21 What are colligative properties of solution?
Ans.
The colligative properties are properties of solution that depend on the no. of solute and solvent molecules or ions. Some of the more important such properties are as follows:
(i) Lowering of vapour pressure.
(ii) Elevation of boiling point.
(iii) Depression of freezing point.
Q.22 What are practical application of colligative properties?
Ans.
It has following applications.
(i) Methods of molecular mass determination.
(ii) Contributed to the development of solution theory.
(iii) Use of NaCl or KNO3 to lower the melting point of ice, this freezing mixture of ice and salt is used in ice cream machine.
(iv) It protects automobile’s radiator when antifreeze glycol is added in the radiators.
Q.23 What are ebullioscopic and cryoscopic constants?
Ans.
The elevation in boiling point when one mole of solute is dissolved to the one kg. of solvent is called molar boiling point constant or ebullioscopic constant.
The depression in freezing point when one mole of solute is dissolved in 1 kg. Of solvent is called molar freezing point constant or cryoscopic constant.
Q.24 Colligative properties are obeyed when solutions are dilute, explain.
Ans.
When solution is concentrated, Raoult’s law is not obeyed when concentration of solute is high then these molecules may associate or combine with each other and all colligative properties show non–ideality.
Q.25 Boiling point of solvent increases due to the pressure of solutes or impurities?
Ans.
Boiling point of a substance depends upon the external pressure and v.p of liquid. When some solute is added to the solvent, the vapour pressure of solvent lowers. Lowering of v.p depends upon the concentration of solution.
Q.26 what are hydration energy and solvation?
Ans.
The enthalpy change when one mole of solute is dissolved in specific amount of water at given temp. is called hydration energy. The enthalpy changes when one mole of solute is dissolved in specific amount of solvent at given temp. is called energy of solvation.
Q.27 Explain why lattice energy of ionic solids is always higher than molecular solids?
Ans.
Lattice energy of ionic solid is large due greater force attraction between the –ve and +ve ions. In case of molecular solid, less amount of energy is required to separate the molecules, because they have less intermolecular forces.
Q.28 What are hydrates?
Ans.
Those substances which have some water of crystallization in them, are called as hydrates e.g., MgSO4 . 7H2O.
Q.29 Colligative properties are obeyed when the solute is non–electrolyte
Ans.
Colligative properties depend upon the no. of particles. In the case of electrolytes, ions are produced and so the number of particles of the solutions increase and amount of colligative properties also increase. If the electrolyte is not 100% dissociated then the amount of colligative properties have to assessed according to their degree of dissociation.
Q.30 Relative lowering of v.p is independent of temperature.
Ans.
The formula for the relative lowering of vapour pressure and mole fraction of solute is
= x2
Vapour pressure depends upon temperature and lowering of v.p also depends upon temp. So when the temp of a solution is increased both the factors increase in such a way that the ratio remains the same.
Think! Why molal Solution is more dilute than Molar solution?
How sum of all mole fraction is unity?
Q.1 Binary solution can be homogenous or heterogeneous explain?
Ans.
The solutions which contain two components only are called as binary solution. If binary solution has single phase, it is called homogenous solution e.g. glucose in water. If binary solutions has more than one phase it is called heterogeneous solution e.g. oil in water.
Q.2 What is phase?
Ans.
Every sample of matter with uniform properties and fixed composition is called phase.
Q.3 Define molarity.
Ans.
Molarity is the number of moles of solute dissolved per dm3 of solution.
Q.4 What is molarity?
Ans.
Molarity is the no. of moles of solute in 1000 gm (1 kg) of the solvent.
Q.5 Explain mole fraction (x).
Ans.
The mole fraction of any component in a mixture is the ratio of the number of moles of it to the total number of moles of all the components present.
Q.6 Define PPm?
Ans.
It is defined as number of parts (by weight or volume) of a solute per million parts (by weight or volume) of the solution.
Q.7 Like dissolve like. Explain.
Ans.
The inter–Ionic forces of attraction are very strong in Ionic solids, so equally strong polar solvents are needed to dissolve them. such solids cannot be dissolved by moderately polar solvents e.g. acetone.
Q.8 What is critical solution or upper consulate temperature?
Ans.
The temperature at which two conjugate solutions merge into one another is called critical solution temperature.
Q.9 What is difference between ideal and non–ideal solution.
Ans.
IDEAL SOLUTION
NON–IDEAL SOLUTION
1.
Ideal solution obeys the Raoult’s law.
1.
Non–Ideal solution does not obey Raoult’s law.
2.
It has zero enthalpy change as their heat of solution D H = 0.
2.
It has exothermic or endothermic enthalpy change as their heat of solution D H ¹ 0
3.
Volume of solution is sum of volume of solvent and solute and change in volume is zero D v = 0.
3.
Volume of solution is not sum of the volume of solute and solvents and change in volume D V ¹ 0.
4.
Examples
(i) Benzene – ether
(ii) Chlorobenzene–bromobenzene
4.
Example
(i) Alcohol – water
(ii) Water HCl
Q.10 Define Zeotropic and Azeotropic mixture.
Ans.
The liquid mixture which boils at constant temperature and distills over without change in composition at any pressure like a chemical compound is called azeotropic mixture e.g. water HCl system, water–ethanol system. The liquid mixture which can be distilled over with a change in composition is called zeotropic mixture e.g. Benzene–ether system.
Q.11 How can you relate the dynamic equilibrium with recrystallization?
Ans.
When some solute is added to the solvent, the force of attraction between the solute particles breakdown. This process is called dissolution. Some particles of solute may combine again and converted to a solid substance. This process is called recrystalization when a saturated solution is prepared there is an equilibrium b/w dissolution and recrystallization.
Q.12 Prove that HCl form an Azeotropic mixture with water?
Ans.
HCl form an azeotropic mixture with water boiling at 110o and containing 20.24% of acid. This pair of liquids showing positive deviation has azeotropic mixture with boiling point in comparison with pure components.
Q.13 What are conjugate solutions?
Ans.
As the mutual solubilities are limited, the liquids are only partially miscible on shaking equal volumes of water and ether, two layers are formed. Each liquid layer is a saturated solution of the other liquid. Such solution are called conjugate solutions.
Q.14 State Raoult’s law in different way.
Ans. 1st statement:
The vapour pressure of a solvent above a solution is equal to the product of the vapour pressure of pure solvent and the mole fraction of solvent in solution.
2nd statement:
The lowering of vapour pressure is directly proportional to the mole fraction of solute.
3rd statement:
The relative lowering of vapour pressure is equal to the mole fraction of solute.
Q.15 The relative lowering of v.p is better than lowering of v.p why?
Ans.
The relative lowering of v.p:
(i) is independent of temp.
(ii) depends upon the cone. of solute.
(iii) is constant when equimolar proportions of different solutes are dissolved in the same mass of same solvent.
Q.16 What is positive deviation?
Ans.
If a graph is plotted b/w composition and vapour pressure of a solution the total vapour pressure curve rises to maximum, which is above the vapour pressure of either of the pure components, shows positive deviation from Raoult’s law.
Q.17 What is negative deviation?
Ans.
If the vapour pressure show a minimum. When vapour pressure of solution becomes less than either of component the solution will show negative deviation from Raoult’s law.
Q.18 Define solubility?
Ans.
It is defined as the concentration of solute in the solution when it is in equilibrium with the solid substance at a particular temperature.
Q.19 Ce2(SO4) shows exceptional behaviour toward solubility.
Ans.
Ce2 (SO4) shows exceptional behaviour whose solubility decreases with the increase in temperature becomes constant from 40oC onwards. Anyhow, it shows continuous solubility curve.
Q.20 Na2SO4 . 10H2O and CaCl2 . 6H2O show discontinuous solubility curves?
Ans.
Sometimes the solubility curves show sudden changes of solubilities and these curves are called discontinuous solubility curves. Actually these curves are combination of many curves. At the break a new solid phase appears and another solubility curve of new phase begins.
Q.21 What are colligative properties of solution?
Ans.
The colligative properties are properties of solution that depend on the no. of solute and solvent molecules or ions. Some of the more important such properties are as follows:
(i) Lowering of vapour pressure.
(ii) Elevation of boiling point.
(iii) Depression of freezing point.
Q.22 What are practical application of colligative properties?
Ans.
It has following applications.
(i) Methods of molecular mass determination.
(ii) Contributed to the development of solution theory.
(iii) Use of NaCl or KNO3 to lower the melting point of ice, this freezing mixture of ice and salt is used in ice cream machine.
(iv) It protects automobile’s radiator when antifreeze glycol is added in the radiators.
Q.23 What are ebullioscopic and cryoscopic constants?
Ans.
The elevation in boiling point when one mole of solute is dissolved to the one kg. of solvent is called molar boiling point constant or ebullioscopic constant.
The depression in freezing point when one mole of solute is dissolved in 1 kg. Of solvent is called molar freezing point constant or cryoscopic constant.
Q.24 Colligative properties are obeyed when solutions are dilute, explain.
Ans.
When solution is concentrated, Raoult’s law is not obeyed when concentration of solute is high then these molecules may associate or combine with each other and all colligative properties show non–ideality.
Q.25 Boiling point of solvent increases due to the pressure of solutes or impurities?
Ans.
Boiling point of a substance depends upon the external pressure and v.p of liquid. When some solute is added to the solvent, the vapour pressure of solvent lowers. Lowering of v.p depends upon the concentration of solution.
Q.26 what are hydration energy and solvation?
Ans.
The enthalpy change when one mole of solute is dissolved in specific amount of water at given temp. is called hydration energy. The enthalpy changes when one mole of solute is dissolved in specific amount of solvent at given temp. is called energy of solvation.
Q.27 Explain why lattice energy of ionic solids is always higher than molecular solids?
Ans.
Lattice energy of ionic solid is large due greater force attraction between the –ve and +ve ions. In case of molecular solid, less amount of energy is required to separate the molecules, because they have less intermolecular forces.
Q.28 What are hydrates?
Ans.
Those substances which have some water of crystallization in them, are called as hydrates e.g., MgSO4 . 7H2O.
Q.29 Colligative properties are obeyed when the solute is non–electrolyte
Ans.
Colligative properties depend upon the no. of particles. In the case of electrolytes, ions are produced and so the number of particles of the solutions increase and amount of colligative properties also increase. If the electrolyte is not 100% dissociated then the amount of colligative properties have to assessed according to their degree of dissociation.
Q.30 Relative lowering of v.p is independent of temperature.
Ans.
The formula for the relative lowering of vapour pressure and mole fraction of solute is
= x2
Vapour pressure depends upon temperature and lowering of v.p also depends upon temp. So when the temp of a solution is increased both the factors increase in such a way that the ratio remains the same.
Think! Why molal Solution is more dilute than Molar solution?
How sum of all mole fraction is unity?
CHAP 10 ****ELECTROCHEMISTRY***
SHORT QUESTION WITH ANSWERS
Q.1 Differentiate electrolytic and voltaic cell.
Ans.
An electrochemical cell in which electric current is used to drive a non–spontaneous redox reaction is called electrolytic cell. e.g. electrolysis of molten NaCl.
An electrochemical cell in which electricity is produced as a result of spontaneous redox reaction is called Galvanic or voltaic cell e.g. Zn – Cu cell.
Q.2 Differentiate b/w ionization and electrolysis.
Ans.
Ionization Electrolysis 1. The process of splitting up of an ionic compound into charged particles when fused or dissolved in water is called ionization. 1. The process of decomposing a substance usually in solution or in molten state is called electrolysis. 2. In this process, there is no need of electrodes. 2. In this process, electrodes are required. 3. There is no need of electricity for ionization. 3. There is a need of electricity for electrolysis. 4. During ionization ions are produced. 4. In electrolysis, ions are oxidized or reduced to neutral atoms or molecules. Q.3 What are conductors? Define electronic conductors and electrolytic conductors.
Ans. Conductor:
The substance through which electricity can pass is called conductor. There are two types of conductors.
Electronic Conductor:
The conductors in which electricity is passed due to the movement of free electrons e.g. metals.
Electrolytic Conductors:
The conductors in which electricity is passed, due to movement of ions to their respective electrodes.
Q.4 Differentiate b/w oxidation and reduction.
Ans. Oxidation:
(i) Addition of oxygen
(ii) Removal of electrons or
(iii) Removal of hydrogen is called as oxidation.
Reduction:
(i) Addition of hydrogen
(ii) Addition of electrons or
(iii) Removal of oxygen is called as reduction.
Q.5 Define oxidation number. Write down the rules for assigning oxidation number.
Ans. Oxidation Number:
The apparent charge present on an atom in a molecule or ionic compound is called oxidation number.
Rules:
(1) Oxidation state of all elements in free state is zero e.g.
O eq \a\co1(0,2) , N eq \a\co1(0,2) , F eq \a\co1(0,2) , Cl eq \a\co1(0,2) etc.
(2) Oxidation number of hydrogen is + 1 except in case of metal hydride where its is – 1 H+1Cl and NaH–1.
(3) Oxidation number of oxygen is – 2 except in case of per oxides and super oxide. Where it is – 1 and – eq \f(1,2) respectively.
(4) In binary compounds of halogens, the charge of halogen is
– 1 e.g. NaCl–1.
(5) Oxidation number of IA group elements in combined form is + 1, Group II is + 2 and group III is + 3.
(6) The sum of oxidation state of all elements in neutral compounds is zero e.g. K+1M+7nO eq \a\co1(–2(4),4 ) .
(7) The sum of oxidation states of ion is equal to charge present on it P+5O eq \a\co1(–3,4 ) eq \x( P = + 5 )
Q.6 K2Cr2O7 + HCl ( KCl + CrCl3 + Cl2 + H2O
Calculate the oxidation number of each element in above equation?
Ans.
K eq \a\co1(+1(2),2 ) Cr eq \a\co1(+6(2),2 ) O eq \a\co1(–2(7),7 ) + H+1Cl–1 (
K+1Cl–1 + Cr+3 Cl eq \a\co1(–1(3),3 ) + Cl eq \a\co1(0,2) + H eq \a\co1(+1(2),2 ) O
Q.7 Calculate oxidation number of Cr in the following compounds.
(i) CrCl3 (ii) Cr2(SO4)3 (iii) K2CrO4
(iv) K2Cr2O7 (v) CrO3 (vi) Cr2O3
(vii) Cr2O eq \a\co1(2–,7 )
Ans.
(i) CrCl3 ( x + (–1) 3 = 0 x = 0 + 3
(ii) Cr2(SO4)3 ( 2x + 6 (3) + (–2) 4 (3) = 0
= 2x + 18 – 24 = 0
2x – 6 = 0
2x = 6 eq \x( x = 3 )
(iii) K2CrO4 ( + 1 (2) + x + – 2 (4) = 0
= 2 + x – 8 = 0 eq \x( x = + 6 )
(iv) K2Cr2O7 ( + 1 (2) + 2x + – 2 (7) = 0
+ 2 + 2x – 14 = 0
2x – 12 = 0
2x = 12
eq \x( x = + 6 )
(v) CrO3 ( x + (–2) 3 = 0
eq \x( x = + 6 )
(vi) Cr2O3 ( 2x + (–2) 3 = 0
2x – 6 = 0
2x = + 6
eq \x(x = + 3 )
(vii) Cr2O eq \a\co1(–2,7 ) ( 2x + (–2) 7 = – 2
2x – 14 = – 2
2x = + 12
eq \x( x = + 6 )
Q.8 In a Galvanic cell chemical energy is converted to electrical energy. How?
Ans.
In galvanic cell chemical energy changes to electrical energy because in this cell redox reaction takes place spontaneously. As a result of which electrons are transferred from anode to cathode. These electron flow is called current or electrical energy.
Q.9 Zn acts as anode when connected to Cu but as cathode when connected to Al?
Ans.
Zn acts as anode when connected to Cu because reduction potential of Cu is greater than Zn. So Zn will be oxidised and acts as anode Cu acts as cathode. But in case of Al connected to Zn. Zinc acts strong oxidising agent (with greater reduction potential) than Al so Al will act as anode and Zn acts as cathode.
Q.10 Explain through equations how lead battery is discharged.
Ans.
In this process, the anode and cathode of an external source are joined to the anode and cathode of the cell respectively. Then the redox cathode of the cell respectively. Then the redox reactions at respective electrodes are reversed. Hence cell begins to recharge reaction at cathode:
Pb eq \a\co1(+2 ,(aq)) + Ze– ( Pb eq \a\co1(0 ,(s)) (Reduction)
PbSO4(s)
SO eq \a\co1(2–,4 ) (aq)
Reaction at Anode:
+2H2O
Pb+2(aq) ( PbO2 + OH+ + 2e–
PbSO4(s)
SO eq \a\co1(2– ,4(aq))
Overall Reaction:
2PbSO4(s) + 2H2O(l) ( Pb(s) + PbO2(s) + 4H+ + 2SO eq \a\co1(2– ,4(aq))
Q.11 What is density of acid used in lead accumulator and voltage of cell?
Ans.
When lead accumulator is in use it acid conc. falls and density decrease to 1.15 g/cm3. Its e.m.f. also decreases. After recharging conc. of acid is increased to make up its density to 1.25 gcm–3 and voltage becomes 2 volts.
Q.12 What are advantages of Alkaline battery?
Ans.
(1) It can be used over a wider range of temp due to more stability of electrolysis.
(2) It requires very little electrolyte so it is very compact.
(3) It maintains a constant voltage for a longer period of time and therefore lasts longer.
Q.13 What are silver oxide batteries?
Ans.
These are tiny and expensive batteries, which are commonly used in electronic watches, auto exposure camera and calculators. These are small button shape cells. In these types of batteries silver oxide (Ag2O) mixed with NaOH or KOH acts as cathode and zinc functions as anode.
Q.14 Discuss reactions taking place in the Ni–Cd cell.
Ans.
It is strong cell. The anode is made up of cadmium at which oxidation reaction takes place while cathode consists of NiO2. In this cell, an alkaline electrolyte is used.
Reactions:
At anode:
Cd(s) + 2OH eq \a\co1(– ,(aq)) ( Cd(OH)2(s) + 2e–
NiO2(s) + 2H2O(l) + 2e– ( Ni(OH)2(s) + 20H–
Overall Reactions:
Cd(s) + NiO2(s) + 2H2O ( Cd(OH)2(s) + Ni(OH)2(s)
Q.15 How fuel is converted into electrical energy in fuel cells?
Ans.
A type of cell, which is similar to galvanic cell in which, fuel used may be gaseous or liquid substance e.g. hydrogen, hydrazine, methanol, ammonia etc. and oxidants are generally oxygen or air. Fuel and oxidant are supplied to respective electrodes, which are porous and activated with catalyst.
Reactions:
At anode:
H2(g) + 2OH eq \a\co1(– ,(aq)) ( 2H2O(l) + 2e– (oxidation)
At cathode:
O2(g) + 2H2O(l) + 4e– ( 4OH eq \a\co1(– ,(aq)) (Reduction)
Overall reaction
H2(g) + O2(g) ( 2H2O(l)
Q.16 What is importance of fuel cells?
Ans.
The fuel cells are of great importance and used in space vehicles. In these cells, electrodes are made up of porous compressed carbon, which is impregnated with platinum. Which acts as catalyst. This fuel cell works at high temp in order to evaporate water, which is formed in it. After condensation, this water can be used for drinking water for astronauts.
Q.17 How prediction about feasibility of chemical reaction is made by electrochemical series?
Ans.
For a particular reaction, it is easy to predict whether it will take place or not. This is done by using electrochemical series e.g. Cu+2 can oxidise zinc metal but Zn+2 cannot oxidise Cu metal because Cu is below Zn in the electrochemical series. i.e.
Cu eq \a\co1(+2 ,(aq)) + Zn(s) ( Cu(s) + Zn eq \a\co1(2+,(aq)) . E eq \a\co1(0 ,cell) = 1.10 v
Cu(s) + Zn+2 ( Cu+2 + Zn(s) E eq \a\co1(0 ,cell) = – 1 . 1
This is not feasible.
Q.18 How e.m.f. of galvanic cell is calculated with the help of electrochemical series?
Ans.
The standard electrode potential values can be used to determine e.m.f. of a given galvanic cell i.e. e.m.f. of Daniel cell is calculated as
Zn(s) | Zn eq \a\co1(+2,(aq)) || Cu eq \a\co1(2+,(aq)) | Cu(s)
(Anode) Cathode
(Oxidation half cell) Reduction half cell)
Reactions:
At anode
Zn ( Zn2+ + 2e– E0 = + 0.76 v
+2e–
Cu+2 ( Cu(s) E0 = + 0.34 v
E eq \a\co1(0 ,cell) = E eq \a\co1(0 ,Anode) + E eq \a\co1(0 ,Cathode)
= 0.76 + 0.34 = 1.10 v
Q.19 What are relative reactivity of metals? Explain by electrochemical series?
Ans.
Reactivity of a metal depends upon its ability to lose electrons to change into M+ cation. It is clear from electrochemical series that smaller the value of standard reduction potential, greater its tendency to lose electrons. Alkali metals (Li, NaK) having lower value of reduction potentials are highly reactive then coinage metals
(Cu, Ag, Au).
Q.20 How it is possible to select a element as anode or cathode?
Ans.
The electrode, which has lower value of reduction potential acts as anode and that which has higher value of reduction potential acts as cathode E0 eq \b( \f(Cu+2,Cu) ) = 0.34 has a lower value of reduction potential than eq \f(E\a\co1(0 ,Ag+),Ag) = 0.80 v anode and Ag as cathode.
Q.21 How we can predict whether it takes place or not?
Ans.
With the help of electrochemical series, it can be determined whether a reaction will occur spontaneously or not. Consider the following reaction
Pb+2 + 2Ag0 ( 2Ag+ + Pb0
The reaction will occur only if E eq \a\co1(0 ,Cell) is positive
E eq \a\co1(0 ,Cell) = E eq \a\co1(0 ,Anode) + E eq \a\co1(0 ,Cathode)
E eq \a\co1(0 ,Cell) = E eq \a\co1(0 ,2Ag/Ag+) + E eq \a\co1(0 ,Pb2+/Pb)
E eq \a\co1(0 ,Cell) = – 0.80 + (– 0.13)
= – 0.93
E eq \a\co1(0 ,Cell) is negative so reaction will not take place spontaneously.
Q.22 How can a non–metal displace another non–metal?
Ans.
A non–metal of higher standard reduction potential can displace other non–metal of lower standard reduction potential
Cl2 + MgBr2 ( MgCl2 + Br2
or Cl2(g) + 2Br eq \a\co1(– ,(aq)) ( 2Cl eq \a\co1(– ,(aq)) + Br2(g)
In above reaction E eq \a\co1(0 ,Cl2/2Cl–) = + 1.36 v
is greater than E eq \a\co1(0 ,Br2/2Br–) = + 1.06 v.
Q.23 What is standard hydrogen electrode (SHE)?
Ans.
Standard hydrogen electrode is used to determine the electrode potential of other electrodes. It is used as reference electrode and its value is taken as 0.0 v.
SHE consist of platinum foil, coated with finally divided black platinum and suspended in 1 M HCl solution. Pure H2 gas at one atm pressure is bubbled into 1M HCl solution. The value of this half cell will be zero either it is oxidised or reduced
2H+ + 2e– ( H2 0.00 V (Reduction)
H2 ( 2H+ + 2e– 0.00 V (Oxidation)
Q.24 What is difference b/w cell and battery?
Ans.
The arrangement in which electrical energy is converted to chemical energy or chemical energy is converted to electrical energy is called cell.
The combination of two or more cells is called battery.
Q.25 Outline the important applications of electrolysis?
Ans.
(1) Electroplating
(2) Electro–refining
(3) Electro–manufacturing
(4) Electrotyping
Q.26 What is salt bridge? Also give its function?
Ans:
It is U-shaped glass tube having saturated solution of strong electrolyte like KCl. The glass tube is filled with jelly type material
It has two major functions:
It connects the two solutions in two half cells
It maintains the electrical neutrality by the diffusion of ions through it.
Q.27 What is the difference between primary and secondary cell?
Ans:
The cell which is not rechargeable called primary cell e.g Dry cell
The cell which is rechargeable called secondary cell. e.g lead accumulator
Q.28 What is electrochemical series?
Ans:
A list of elements in which they are arranged in the order of their standard electrode potential on hydrogen scale is called electrochemical series.
Q.1 Differentiate electrolytic and voltaic cell.
Ans.
An electrochemical cell in which electric current is used to drive a non–spontaneous redox reaction is called electrolytic cell. e.g. electrolysis of molten NaCl.
An electrochemical cell in which electricity is produced as a result of spontaneous redox reaction is called Galvanic or voltaic cell e.g. Zn – Cu cell.
Q.2 Differentiate b/w ionization and electrolysis.
Ans.
Ionization Electrolysis 1. The process of splitting up of an ionic compound into charged particles when fused or dissolved in water is called ionization. 1. The process of decomposing a substance usually in solution or in molten state is called electrolysis. 2. In this process, there is no need of electrodes. 2. In this process, electrodes are required. 3. There is no need of electricity for ionization. 3. There is a need of electricity for electrolysis. 4. During ionization ions are produced. 4. In electrolysis, ions are oxidized or reduced to neutral atoms or molecules. Q.3 What are conductors? Define electronic conductors and electrolytic conductors.
Ans. Conductor:
The substance through which electricity can pass is called conductor. There are two types of conductors.
Electronic Conductor:
The conductors in which electricity is passed due to the movement of free electrons e.g. metals.
Electrolytic Conductors:
The conductors in which electricity is passed, due to movement of ions to their respective electrodes.
Q.4 Differentiate b/w oxidation and reduction.
Ans. Oxidation:
(i) Addition of oxygen
(ii) Removal of electrons or
(iii) Removal of hydrogen is called as oxidation.
Reduction:
(i) Addition of hydrogen
(ii) Addition of electrons or
(iii) Removal of oxygen is called as reduction.
Q.5 Define oxidation number. Write down the rules for assigning oxidation number.
Ans. Oxidation Number:
The apparent charge present on an atom in a molecule or ionic compound is called oxidation number.
Rules:
(1) Oxidation state of all elements in free state is zero e.g.
O eq \a\co1(0,2) , N eq \a\co1(0,2) , F eq \a\co1(0,2) , Cl eq \a\co1(0,2) etc.
(2) Oxidation number of hydrogen is + 1 except in case of metal hydride where its is – 1 H+1Cl and NaH–1.
(3) Oxidation number of oxygen is – 2 except in case of per oxides and super oxide. Where it is – 1 and – eq \f(1,2) respectively.
(4) In binary compounds of halogens, the charge of halogen is
– 1 e.g. NaCl–1.
(5) Oxidation number of IA group elements in combined form is + 1, Group II is + 2 and group III is + 3.
(6) The sum of oxidation state of all elements in neutral compounds is zero e.g. K+1M+7nO eq \a\co1(–2(4),4 ) .
(7) The sum of oxidation states of ion is equal to charge present on it P+5O eq \a\co1(–3,4 ) eq \x( P = + 5 )
Q.6 K2Cr2O7 + HCl ( KCl + CrCl3 + Cl2 + H2O
Calculate the oxidation number of each element in above equation?
Ans.
K eq \a\co1(+1(2),2 ) Cr eq \a\co1(+6(2),2 ) O eq \a\co1(–2(7),7 ) + H+1Cl–1 (
K+1Cl–1 + Cr+3 Cl eq \a\co1(–1(3),3 ) + Cl eq \a\co1(0,2) + H eq \a\co1(+1(2),2 ) O
Q.7 Calculate oxidation number of Cr in the following compounds.
(i) CrCl3 (ii) Cr2(SO4)3 (iii) K2CrO4
(iv) K2Cr2O7 (v) CrO3 (vi) Cr2O3
(vii) Cr2O eq \a\co1(2–,7 )
Ans.
(i) CrCl3 ( x + (–1) 3 = 0 x = 0 + 3
(ii) Cr2(SO4)3 ( 2x + 6 (3) + (–2) 4 (3) = 0
= 2x + 18 – 24 = 0
2x – 6 = 0
2x = 6 eq \x( x = 3 )
(iii) K2CrO4 ( + 1 (2) + x + – 2 (4) = 0
= 2 + x – 8 = 0 eq \x( x = + 6 )
(iv) K2Cr2O7 ( + 1 (2) + 2x + – 2 (7) = 0
+ 2 + 2x – 14 = 0
2x – 12 = 0
2x = 12
eq \x( x = + 6 )
(v) CrO3 ( x + (–2) 3 = 0
eq \x( x = + 6 )
(vi) Cr2O3 ( 2x + (–2) 3 = 0
2x – 6 = 0
2x = + 6
eq \x(x = + 3 )
(vii) Cr2O eq \a\co1(–2,7 ) ( 2x + (–2) 7 = – 2
2x – 14 = – 2
2x = + 12
eq \x( x = + 6 )
Q.8 In a Galvanic cell chemical energy is converted to electrical energy. How?
Ans.
In galvanic cell chemical energy changes to electrical energy because in this cell redox reaction takes place spontaneously. As a result of which electrons are transferred from anode to cathode. These electron flow is called current or electrical energy.
Q.9 Zn acts as anode when connected to Cu but as cathode when connected to Al?
Ans.
Zn acts as anode when connected to Cu because reduction potential of Cu is greater than Zn. So Zn will be oxidised and acts as anode Cu acts as cathode. But in case of Al connected to Zn. Zinc acts strong oxidising agent (with greater reduction potential) than Al so Al will act as anode and Zn acts as cathode.
Q.10 Explain through equations how lead battery is discharged.
Ans.
In this process, the anode and cathode of an external source are joined to the anode and cathode of the cell respectively. Then the redox cathode of the cell respectively. Then the redox reactions at respective electrodes are reversed. Hence cell begins to recharge reaction at cathode:
Pb eq \a\co1(+2 ,(aq)) + Ze– ( Pb eq \a\co1(0 ,(s)) (Reduction)
PbSO4(s)
SO eq \a\co1(2–,4 ) (aq)
Reaction at Anode:
+2H2O
Pb+2(aq) ( PbO2 + OH+ + 2e–
PbSO4(s)
SO eq \a\co1(2– ,4(aq))
Overall Reaction:
2PbSO4(s) + 2H2O(l) ( Pb(s) + PbO2(s) + 4H+ + 2SO eq \a\co1(2– ,4(aq))
Q.11 What is density of acid used in lead accumulator and voltage of cell?
Ans.
When lead accumulator is in use it acid conc. falls and density decrease to 1.15 g/cm3. Its e.m.f. also decreases. After recharging conc. of acid is increased to make up its density to 1.25 gcm–3 and voltage becomes 2 volts.
Q.12 What are advantages of Alkaline battery?
Ans.
(1) It can be used over a wider range of temp due to more stability of electrolysis.
(2) It requires very little electrolyte so it is very compact.
(3) It maintains a constant voltage for a longer period of time and therefore lasts longer.
Q.13 What are silver oxide batteries?
Ans.
These are tiny and expensive batteries, which are commonly used in electronic watches, auto exposure camera and calculators. These are small button shape cells. In these types of batteries silver oxide (Ag2O) mixed with NaOH or KOH acts as cathode and zinc functions as anode.
Q.14 Discuss reactions taking place in the Ni–Cd cell.
Ans.
It is strong cell. The anode is made up of cadmium at which oxidation reaction takes place while cathode consists of NiO2. In this cell, an alkaline electrolyte is used.
Reactions:
At anode:
Cd(s) + 2OH eq \a\co1(– ,(aq)) ( Cd(OH)2(s) + 2e–
NiO2(s) + 2H2O(l) + 2e– ( Ni(OH)2(s) + 20H–
Overall Reactions:
Cd(s) + NiO2(s) + 2H2O ( Cd(OH)2(s) + Ni(OH)2(s)
Q.15 How fuel is converted into electrical energy in fuel cells?
Ans.
A type of cell, which is similar to galvanic cell in which, fuel used may be gaseous or liquid substance e.g. hydrogen, hydrazine, methanol, ammonia etc. and oxidants are generally oxygen or air. Fuel and oxidant are supplied to respective electrodes, which are porous and activated with catalyst.
Reactions:
At anode:
H2(g) + 2OH eq \a\co1(– ,(aq)) ( 2H2O(l) + 2e– (oxidation)
At cathode:
O2(g) + 2H2O(l) + 4e– ( 4OH eq \a\co1(– ,(aq)) (Reduction)
Overall reaction
H2(g) + O2(g) ( 2H2O(l)
Q.16 What is importance of fuel cells?
Ans.
The fuel cells are of great importance and used in space vehicles. In these cells, electrodes are made up of porous compressed carbon, which is impregnated with platinum. Which acts as catalyst. This fuel cell works at high temp in order to evaporate water, which is formed in it. After condensation, this water can be used for drinking water for astronauts.
Q.17 How prediction about feasibility of chemical reaction is made by electrochemical series?
Ans.
For a particular reaction, it is easy to predict whether it will take place or not. This is done by using electrochemical series e.g. Cu+2 can oxidise zinc metal but Zn+2 cannot oxidise Cu metal because Cu is below Zn in the electrochemical series. i.e.
Cu eq \a\co1(+2 ,(aq)) + Zn(s) ( Cu(s) + Zn eq \a\co1(2+,(aq)) . E eq \a\co1(0 ,cell) = 1.10 v
Cu(s) + Zn+2 ( Cu+2 + Zn(s) E eq \a\co1(0 ,cell) = – 1 . 1
This is not feasible.
Q.18 How e.m.f. of galvanic cell is calculated with the help of electrochemical series?
Ans.
The standard electrode potential values can be used to determine e.m.f. of a given galvanic cell i.e. e.m.f. of Daniel cell is calculated as
Zn(s) | Zn eq \a\co1(+2,(aq)) || Cu eq \a\co1(2+,(aq)) | Cu(s)
(Anode) Cathode
(Oxidation half cell) Reduction half cell)
Reactions:
At anode
Zn ( Zn2+ + 2e– E0 = + 0.76 v
+2e–
Cu+2 ( Cu(s) E0 = + 0.34 v
E eq \a\co1(0 ,cell) = E eq \a\co1(0 ,Anode) + E eq \a\co1(0 ,Cathode)
= 0.76 + 0.34 = 1.10 v
Q.19 What are relative reactivity of metals? Explain by electrochemical series?
Ans.
Reactivity of a metal depends upon its ability to lose electrons to change into M+ cation. It is clear from electrochemical series that smaller the value of standard reduction potential, greater its tendency to lose electrons. Alkali metals (Li, NaK) having lower value of reduction potentials are highly reactive then coinage metals
(Cu, Ag, Au).
Q.20 How it is possible to select a element as anode or cathode?
Ans.
The electrode, which has lower value of reduction potential acts as anode and that which has higher value of reduction potential acts as cathode E0 eq \b( \f(Cu+2,Cu) ) = 0.34 has a lower value of reduction potential than eq \f(E\a\co1(0 ,Ag+),Ag) = 0.80 v anode and Ag as cathode.
Q.21 How we can predict whether it takes place or not?
Ans.
With the help of electrochemical series, it can be determined whether a reaction will occur spontaneously or not. Consider the following reaction
Pb+2 + 2Ag0 ( 2Ag+ + Pb0
The reaction will occur only if E eq \a\co1(0 ,Cell) is positive
E eq \a\co1(0 ,Cell) = E eq \a\co1(0 ,Anode) + E eq \a\co1(0 ,Cathode)
E eq \a\co1(0 ,Cell) = E eq \a\co1(0 ,2Ag/Ag+) + E eq \a\co1(0 ,Pb2+/Pb)
E eq \a\co1(0 ,Cell) = – 0.80 + (– 0.13)
= – 0.93
E eq \a\co1(0 ,Cell) is negative so reaction will not take place spontaneously.
Q.22 How can a non–metal displace another non–metal?
Ans.
A non–metal of higher standard reduction potential can displace other non–metal of lower standard reduction potential
Cl2 + MgBr2 ( MgCl2 + Br2
or Cl2(g) + 2Br eq \a\co1(– ,(aq)) ( 2Cl eq \a\co1(– ,(aq)) + Br2(g)
In above reaction E eq \a\co1(0 ,Cl2/2Cl–) = + 1.36 v
is greater than E eq \a\co1(0 ,Br2/2Br–) = + 1.06 v.
Q.23 What is standard hydrogen electrode (SHE)?
Ans.
Standard hydrogen electrode is used to determine the electrode potential of other electrodes. It is used as reference electrode and its value is taken as 0.0 v.
SHE consist of platinum foil, coated with finally divided black platinum and suspended in 1 M HCl solution. Pure H2 gas at one atm pressure is bubbled into 1M HCl solution. The value of this half cell will be zero either it is oxidised or reduced
2H+ + 2e– ( H2 0.00 V (Reduction)
H2 ( 2H+ + 2e– 0.00 V (Oxidation)
Q.24 What is difference b/w cell and battery?
Ans.
The arrangement in which electrical energy is converted to chemical energy or chemical energy is converted to electrical energy is called cell.
The combination of two or more cells is called battery.
Q.25 Outline the important applications of electrolysis?
Ans.
(1) Electroplating
(2) Electro–refining
(3) Electro–manufacturing
(4) Electrotyping
Q.26 What is salt bridge? Also give its function?
Ans:
It is U-shaped glass tube having saturated solution of strong electrolyte like KCl. The glass tube is filled with jelly type material
It has two major functions:
It connects the two solutions in two half cells
It maintains the electrical neutrality by the diffusion of ions through it.
Q.27 What is the difference between primary and secondary cell?
Ans:
The cell which is not rechargeable called primary cell e.g Dry cell
The cell which is rechargeable called secondary cell. e.g lead accumulator
Q.28 What is electrochemical series?
Ans:
A list of elements in which they are arranged in the order of their standard electrode potential on hydrogen scale is called electrochemical series.
Chap 11 ***REACTION KINETICS**
SHORT QUESTION WITH ANSWERS
Q.1 What is meant by chemical kinetics?
Ans.
It is that branch of chemistry which helps to study the followings.
(i) Rates of chemical reactions.
(ii) Study the mechanism of reactions.
Q.2 Differentiate between instantaneous rate of reaction and average rate of reaction.
Ans.
The rate of a chemical reaction at a specific time or at any moment is called as instantaneous rate of a reaction. It is very fast at the start and very slow at the end.
Q.3 Rate of a chemical reaction is an ever changing parameter under the given conditions. Comment.
Ans.
According to law of mass action, the rate of chemical reaction is directly product of active masses of reactants. When the reaction progress, the reaction is very fast at the beginning slow, somewhere in the middle and very slow at the end. The reason is that rate depends upon concentration as conc. decreases every moment, rate decreases every moment.
Q.4 A student started to work on 200 gms of a compound A whose half life 10 hours. How much of the compound will be left after 50 hours.
Ans.
10 hours 10 hours 10 hours
200 gm ® 100 gm ® 50 gm ®
10 hours 10 hours
6.25 gm ¬ 12.5 gm ¬ 25 gm
Q.5 What is meant by rate law?
Ans.
The rate of a chemical reaction is directly proportional to the product of concentrations of reactants each raised to some power.
Consider the following reaction
aA + bB ® cC + dD
Rate of reaction µ [A]a [B]b
This equation is called rate law.
Q.6 What is meant by specific rate constant, or rate constant?
Ans.
The rate constant of a reaction is the rate when the concentration of reactant are unity.
For the reaction
A + B ® C + D
Rate of reaction = K [A] [B]
This equation is called rate law where K is called specific rate constant.
When [A] [B] = 1 then
Rate of reaction = K
Q.7 What is meant by order of reaction?
Ans.
The order of reaction is the sum of exponents of the concentration terms in the rate equation. OR
Number of reacting substances whose concentration changes during the chemical reaction.
Q.8 Explain zero order reaction?
Ans.
That chemical reaction whose rate is independent of concentration of reactant e.g. all photochemical reaction are zero order e.g.
chlorophyll
6CO2 + 6H2O ® C6H12O6 + 6CO2
light
Q.9 What are pseudo order reaction?
Ans.
When one of reactant is taken in large excess its concentration is consider as constant and the order becomes pseudo.
CH3 CH3
| |
CH2– C–Br + H2O ® CH3 – C–OH + HBr
| Excess |
CH3 CH3
Rate = K [(CH3)3 CBr]
The conc. of H2O is taken as constant.
Q.10 What is half life period?
Ans.
Half life period of a reaction is the time required to convert 50% of reactants into products e.g. half life of decomposition of N2O5 at 45oC is 24 minutes. For first order reaction, half life period is independent of the initial concentration of reactants.
Q.11 The radioactive decay is always first order reaction.
Ans.
The radioactive decay of a substance is directly proportions to the number of nuclei in the sample. Its half life period is independent of its initial conc. of reactant.
[t1/2] µ
[t1/2] µ For first order n = 1
Q.12 What is meant by rate determining step?
Ans.
If a reaction occurs in several steps, one of the steps is the slowest. The rate of this step determines the overall rate of reaction. This slowest step is called as rate determining step. The number of molecules taking part in the rate determining step appears in the rate equation of the reaction e.g.
NO2(g) + CO(g) ® NO(g) + CO2(g)
The rate equation of the reaction is found to be
Rate = K [NO2]2
This equation shows that rate of reaction is independent of conc. of CO. In other words the equation tells us that reaction involve more than one step.
Q.13 Describe electrical conductivity method for determination of rate of reaction.
Ans.
The rate of a reaction involving ions can be studied by electrical conductivity method the conductivity of such a solution depends upon the rate of change of conc. of the reacting ions or the ions formed during reaction.
Q.14 Explain briefly spectrometry for the rate determination.
Ans.
This method is applicable if a reactant or a product absorbs ultraviolet, visible or infrared radiations. The rate of reaction can be measured by measuring the amount of radiations absorbed.
Q.15 What is difference between dilatometric and refractrometric method?
Ans.
Dilatometric method is useful for those reactions, which involve small volume changes in solutions. The volume change is directly proportional to the extent of reaction, where refractrometric method id applicable to reactions in solutions where there are changes in refractive indices of the substances taking part in the chemical reaction.
Q.16 Describe briefly optical rotation method.
Ans.
In this method, the angle through which plane polarised light is rotated by the reacting mixture is measured by a polarimeter. The extent of rotation determines the concentration of optically active substance. If any of the species in the reaction mixture is optically active, then this method can be followed to find out the rate reaction.
Q.17 Describe Half–life method for finding the order of reaction?
Ans.
Order of reaction can be determined if half–life and initial concentration is inversely proportional to the initial concentration of reactants raised to the power one less than the order of reaction.
Therefore (t1/2)n µ
Let us perform a reaction twice by taking two different initial concentrations a1 and a2 and their half–life periods are found to be t1 and t2 respectively
(t1)n µ and t2 µ
= Dividing the two relations.
Taking log on both sides:
log = (n – 1) log
Rearranging n = 1 +
Q.18 Describe how surface area effect rate of reaction?
Ans.
The increased surface area of reaction increases the possibilities of atoms and molecules of reactants to come in contact with each other and rate increase e.g. Al foil reacts with NaOH moderately when warmed but powered Al reacts rapidly with cold NaOH and H2 is evolved.
Q.19 Differentiate between positive and negative catalyst.
Ans.
The substance which increase the rate of reaction is called positive catalyst
Pt
e.g. 2H2 + O2 ® 2H2O
without pt reaction is very slow when the rate of reaction is retarded by adding a substance, ten it is said to be a negative catalyst e.g. tetra ethyl lead is added to petrol, because it saves the petrol from pre–ignition.
Q.20 Differentiate between auto catalysis and promoter.
Ans.
In some reactions, a product formed act as catalyst and this phenomenon is called auto catalysis. e.g., In the hydrolysis of ethyl acetate, acetic acid acts as catalyst for further reaction.
CH3COOC2H5 + H2O CH3COOH + C2H5OH
where as promoters are those substances which increases the activity of a catalyst. It is therefore called as catalyst of catalyst. e.g. Hydrogenation of vegetable oils is accelerated by Nickle. The activity of Nickle can be increased by using copper and tellurium.
Q.21 Differentiate between homogeneous catalysis and heterogeneous catalysis.
Ans.
That process, in which catalysts and reactants have same phase is called as homogeneous catalysis e.g.,
NO(g)
2SO2(g) + O2(g) 2SO3(g)
but if the catalyst and reactants have different phases it is called as heterogeneous catalysis
Pt
4NH3(g) + 5O2 4NO(g) + 6H2O
Q.22 Can a catalyst can change the equilibrium constant?
Ans.
A catalyst cannot affect the equilibrium constant of a reaction but it helps the equilibrium to establish earlier. The rates of forward and backward steps are increased equally.
Q.23 What is meant by enzyme catalysis?
Ans.
Those reactions in which enzymes are used as catalyst are called as enzyme catalysis e.g.
O
|| Urease
NH2– C–NH2 + H2O ® 2NH3 + CO2
(Urea)
Urea undergoes hydrolysis into NH3 and CO2 in the presence of enzyme urease.
Q.24 What is poisoning of catalyst?
Ans.
Catalytic poisoning happens due to presence of trace amount of substances, which renders them in–effective, such substances are called poisons. The poisoning of a catalyst may be temporary or permanent. In permanent poisoning, the poison reacts chemically with the catalyst. e.g., the manufacturing of H2SO4 in the contact process needs pt as catalyst. The traces of arsenic present as impurities in the reacting gases make platinum ineffective.
Q.25 What is effect of temperature on catalyst?
Ans.
Some catalysts are physically changed by a change in temp and hence their power will be decreased e.g. colloidal catalysts like pt may be coagulated with rise in temp.
Q.1 What is meant by chemical kinetics?
Ans.
It is that branch of chemistry which helps to study the followings.
(i) Rates of chemical reactions.
(ii) Study the mechanism of reactions.
Q.2 Differentiate between instantaneous rate of reaction and average rate of reaction.
Ans.
The rate of a chemical reaction at a specific time or at any moment is called as instantaneous rate of a reaction. It is very fast at the start and very slow at the end.
Q.3 Rate of a chemical reaction is an ever changing parameter under the given conditions. Comment.
Ans.
According to law of mass action, the rate of chemical reaction is directly product of active masses of reactants. When the reaction progress, the reaction is very fast at the beginning slow, somewhere in the middle and very slow at the end. The reason is that rate depends upon concentration as conc. decreases every moment, rate decreases every moment.
Q.4 A student started to work on 200 gms of a compound A whose half life 10 hours. How much of the compound will be left after 50 hours.
Ans.
10 hours 10 hours 10 hours
200 gm ® 100 gm ® 50 gm ®
10 hours 10 hours
6.25 gm ¬ 12.5 gm ¬ 25 gm
Q.5 What is meant by rate law?
Ans.
The rate of a chemical reaction is directly proportional to the product of concentrations of reactants each raised to some power.
Consider the following reaction
aA + bB ® cC + dD
Rate of reaction µ [A]a [B]b
This equation is called rate law.
Q.6 What is meant by specific rate constant, or rate constant?
Ans.
The rate constant of a reaction is the rate when the concentration of reactant are unity.
For the reaction
A + B ® C + D
Rate of reaction = K [A] [B]
This equation is called rate law where K is called specific rate constant.
When [A] [B] = 1 then
Rate of reaction = K
Q.7 What is meant by order of reaction?
Ans.
The order of reaction is the sum of exponents of the concentration terms in the rate equation. OR
Number of reacting substances whose concentration changes during the chemical reaction.
Q.8 Explain zero order reaction?
Ans.
That chemical reaction whose rate is independent of concentration of reactant e.g. all photochemical reaction are zero order e.g.
chlorophyll
6CO2 + 6H2O ® C6H12O6 + 6CO2
light
Q.9 What are pseudo order reaction?
Ans.
When one of reactant is taken in large excess its concentration is consider as constant and the order becomes pseudo.
CH3 CH3
| |
CH2– C–Br + H2O ® CH3 – C–OH + HBr
| Excess |
CH3 CH3
Rate = K [(CH3)3 CBr]
The conc. of H2O is taken as constant.
Q.10 What is half life period?
Ans.
Half life period of a reaction is the time required to convert 50% of reactants into products e.g. half life of decomposition of N2O5 at 45oC is 24 minutes. For first order reaction, half life period is independent of the initial concentration of reactants.
Q.11 The radioactive decay is always first order reaction.
Ans.
The radioactive decay of a substance is directly proportions to the number of nuclei in the sample. Its half life period is independent of its initial conc. of reactant.
[t1/2] µ
[t1/2] µ For first order n = 1
Q.12 What is meant by rate determining step?
Ans.
If a reaction occurs in several steps, one of the steps is the slowest. The rate of this step determines the overall rate of reaction. This slowest step is called as rate determining step. The number of molecules taking part in the rate determining step appears in the rate equation of the reaction e.g.
NO2(g) + CO(g) ® NO(g) + CO2(g)
The rate equation of the reaction is found to be
Rate = K [NO2]2
This equation shows that rate of reaction is independent of conc. of CO. In other words the equation tells us that reaction involve more than one step.
Q.13 Describe electrical conductivity method for determination of rate of reaction.
Ans.
The rate of a reaction involving ions can be studied by electrical conductivity method the conductivity of such a solution depends upon the rate of change of conc. of the reacting ions or the ions formed during reaction.
Q.14 Explain briefly spectrometry for the rate determination.
Ans.
This method is applicable if a reactant or a product absorbs ultraviolet, visible or infrared radiations. The rate of reaction can be measured by measuring the amount of radiations absorbed.
Q.15 What is difference between dilatometric and refractrometric method?
Ans.
Dilatometric method is useful for those reactions, which involve small volume changes in solutions. The volume change is directly proportional to the extent of reaction, where refractrometric method id applicable to reactions in solutions where there are changes in refractive indices of the substances taking part in the chemical reaction.
Q.16 Describe briefly optical rotation method.
Ans.
In this method, the angle through which plane polarised light is rotated by the reacting mixture is measured by a polarimeter. The extent of rotation determines the concentration of optically active substance. If any of the species in the reaction mixture is optically active, then this method can be followed to find out the rate reaction.
Q.17 Describe Half–life method for finding the order of reaction?
Ans.
Order of reaction can be determined if half–life and initial concentration is inversely proportional to the initial concentration of reactants raised to the power one less than the order of reaction.
Therefore (t1/2)n µ
Let us perform a reaction twice by taking two different initial concentrations a1 and a2 and their half–life periods are found to be t1 and t2 respectively
(t1)n µ and t2 µ
= Dividing the two relations.
Taking log on both sides:
log = (n – 1) log
Rearranging n = 1 +
Q.18 Describe how surface area effect rate of reaction?
Ans.
The increased surface area of reaction increases the possibilities of atoms and molecules of reactants to come in contact with each other and rate increase e.g. Al foil reacts with NaOH moderately when warmed but powered Al reacts rapidly with cold NaOH and H2 is evolved.
Q.19 Differentiate between positive and negative catalyst.
Ans.
The substance which increase the rate of reaction is called positive catalyst
Pt
e.g. 2H2 + O2 ® 2H2O
without pt reaction is very slow when the rate of reaction is retarded by adding a substance, ten it is said to be a negative catalyst e.g. tetra ethyl lead is added to petrol, because it saves the petrol from pre–ignition.
Q.20 Differentiate between auto catalysis and promoter.
Ans.
In some reactions, a product formed act as catalyst and this phenomenon is called auto catalysis. e.g., In the hydrolysis of ethyl acetate, acetic acid acts as catalyst for further reaction.
CH3COOC2H5 + H2O CH3COOH + C2H5OH
where as promoters are those substances which increases the activity of a catalyst. It is therefore called as catalyst of catalyst. e.g. Hydrogenation of vegetable oils is accelerated by Nickle. The activity of Nickle can be increased by using copper and tellurium.
Q.21 Differentiate between homogeneous catalysis and heterogeneous catalysis.
Ans.
That process, in which catalysts and reactants have same phase is called as homogeneous catalysis e.g.,
NO(g)
2SO2(g) + O2(g) 2SO3(g)
but if the catalyst and reactants have different phases it is called as heterogeneous catalysis
Pt
4NH3(g) + 5O2 4NO(g) + 6H2O
Q.22 Can a catalyst can change the equilibrium constant?
Ans.
A catalyst cannot affect the equilibrium constant of a reaction but it helps the equilibrium to establish earlier. The rates of forward and backward steps are increased equally.
Q.23 What is meant by enzyme catalysis?
Ans.
Those reactions in which enzymes are used as catalyst are called as enzyme catalysis e.g.
O
|| Urease
NH2– C–NH2 + H2O ® 2NH3 + CO2
(Urea)
Urea undergoes hydrolysis into NH3 and CO2 in the presence of enzyme urease.
Q.24 What is poisoning of catalyst?
Ans.
Catalytic poisoning happens due to presence of trace amount of substances, which renders them in–effective, such substances are called poisons. The poisoning of a catalyst may be temporary or permanent. In permanent poisoning, the poison reacts chemically with the catalyst. e.g., the manufacturing of H2SO4 in the contact process needs pt as catalyst. The traces of arsenic present as impurities in the reacting gases make platinum ineffective.
Q.25 What is effect of temperature on catalyst?
Ans.
Some catalysts are physically changed by a change in temp and hence their power will be decreased e.g. colloidal catalysts like pt may be coagulated with rise in temp.
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