STIOCHIOMETRY

Definitions:

1. Limiting Reactant – The reactant that is completely consumed first in a chemical

reaction, determining the maximum amount of product that can be formed.
2. Excess Reactant – The reactant that remains after the reaction has stopped due to the
complete consumption of the limiting reactant.
3. Actual Yield – The amount of product actually obtained from a chemical reaction, which
is often less than the theoretical yield due to losses or incomplete reactions.
4. Theoretical Yield – The maximum possible amount of product that can be formed from
a given amount of reactants, calculated based on stoichiometry and assuming perfect
reaction conditions.

5.  Empirical Formula – The simplest whole-number ratio of atoms of each element in a

compound. It represents the basic proportion but not necessarily the actual number of
atoms in a molecule.
6.  Molecular Formula – The actual number of atoms of each element in a molecule of a
compound. It is a multiple of the empirical formula and provides the exact composition
of the substance.
Conceptual and calculation questions

Basics of Stoichiometry

1. What does stoichiometry primarily deal with?

a) The study of chemical bonds
b) The quantitative relationships between reactants and products in a chemical reaction
c) The classification of elements in the periodic table
d) The behavior of gases at different temperatures

Answer: b) The quantitative relationships between reactants and products in a chemical

reaction

2. In a balanced chemical equation, the coefficients represent:

a) The mass of each substance
b) The number of moles of each substance
c) The volume of each substance
d) The total number of atoms in the reaction

Answer: b) The number of moles of each substance

2. Mole Concept and Avogadro’s Number

3. One mole of any substance contains:

a) 6.02×10236.02×1023 molecules or atoms
b) 3.14×10223.14×1022 molecules
c) 12.01 grams of the substance
d) The same number of grams as its atomic number

Answer: a) 6.02×10236.02×1023 molecules or atoms

4. Which statement about the mole is correct?

a) One mole of different substances always has the same mass
b) One mole of different substances always has the same number of particles
c) One mole of oxygen gas contains the same number of atoms as one mole of helium gas
d) One mole of water molecules contains one mole of hydrogen atoms

Answer: b) One mole of different substances always has the same number of particles
3. Limiting and Excess Reactants

5. The limiting reactant in a chemical reaction is:

a) The reactant with the largest mass
b) The reactant that determines the amount of product formed
c) The reactant left over after the reaction stops
d) The catalyst used in the reaction

Answer: b) The reactant that determines the amount of product formed

6. If you have more of one reactant than is required to completely react with the other
reactant, the extra reactant is called the:
a) Limiting reactant
b) Excess reactant
c) Stoichiometric reactant
d) Catalyst

Answer: b) Excess reactant

4. Theoretical, Actual, and Percent Yield

7. The theoretical yield of a reaction is:

a) The amount of product actually obtained
b) The maximum possible amount of product that can be formed
c) The amount of reactants consumed in the reaction
d) The minimum possible amount of product that can be formed

Answer: b) The maximum possible amount of product that can be formed

8. Why is the actual yield of a reaction often less than the theoretical yield?
a) Some reactants may not react completely
b) Some product may be lost during purification
c) Side reactions may occur
d) All of the above

Answer: d) All of the above

5. Balancing Chemical Equations

10. Which law is the basis for balancing chemical equations?

a) Law of Definite Proportions
b) Law of Conservation of Mass
c) Law of Multiple Proportions
d) Law of Partial Pressures

Answer: b) Law of Conservation of Mass

11. In a balanced chemical equation, what must be equal on both sides?

a) The total number of molecules
b) The number of moles of reactants and products
c) The total mass and number of atoms of each element
d) The physical states of the reactants and products

Answer: c) The total mass and number of atoms of each element

6. Molar Mass and Conversions

12. The molar mass of a substance is:

a) The mass of one mole of its atoms or molecules
b) The number of atoms in one mole of the substance
c) The ratio of mass to volume
d) The weight of one molecule in kilograms

Answer: a) The mass of one mole of its atoms or molecules

13. How many grams are in one mole of water (H₂O)?

a) 18.02 g
b) 16.00 g
c) 2.02 g
d) 1.008 g

Answer: a) 18.02 g

7. Empirical and Molecular Formulas

14. The empirical formula of a compound represents:
a) The actual number of atoms in a molecule
b) The simplest whole-number ratio of atoms in the compound
c) The arrangement of atoms in the molecule
d) The mass of one mole of the compound

Answer: b) The simplest whole-number ratio of atoms in the compound

15. If the molecular formula of glucose is C₆H₁₂O₆, what is its empirical formula?
a) C₆H₁₂O₆
b) CHO
c) C₂H₄O₂
d) CH₂O

Answer: d) CH₂O

Calculations

How many moles of oxygen gas (O₂) are needed to completely react with 5.0 moles of
hydrogen gas (H₂) to form water (H₂O)?
Given the balanced equation:
2H2+O2→2H2O

2. How many grams of CO₂ are produced when 10.0 g of CH₄ (methane) is burned in
excess oxygen?
Given the reaction:
CH4+2O2→CO2+2H2O
(Molar masses: C = 12.01 g/mol, H = 1.008 g/mol, O = 16.00 g/mol)

3. If 4.00 moles of aluminum react with excess oxygen to form aluminum oxide (Al₂O₃),
how many moles of Al₂O₃ are produced?
Given:
4Al+3O2→2Al2O3

2. Limiting Reactant Problems

4. If 3.0 moles of N₂ react with 8.0 moles of H₂ to form NH₃, what is the limiting reactant?
Given:
N2+3H2→2NH3

5. 10.0 g of Fe reacts with 10.0 g of S to form FeS. What is the limiting reactant?
Given:
Fe+S→FeS

(Molar masses: Fe = 55.85 g/mol, S = 32.07 g/mol)

6. If 5.0 g of C₂H₆ reacts with 8.0 g of O₂, what is the limiting reactant?
Given:
2C2H6+7O2→4CO2+6H2O2C2H6+7O2→4CO2+6H2O
(Molar masses: C₂H₆ = 30.07 g/mol, O₂ = 32.00 g/mol)

3. Theoretical and Actual Yield Calculations

7. How many grams of NaCl are produced when 5.0 g of Na reacts with excess Cl₂?
Given:
2Na+Cl2→2NaCl2

(Molar masses: Na = 22.99 g/mol, Cl₂ = 70.90 g/mol, NaCl = 58.44 g/mol)

8. If the theoretical yield of a reaction is 25.0 g but only 18.5 g of product is obtained, what
is the percentage yield?
Percent Yield=Actual YieldTheoretical Yield×100Percent Yield=Theoretical YieldActual Yield
×100

9. When 12.0 g of C reacts with excess O₂ to form CO₂, what is the theoretical yield of
CO₂?
Given:
C+O2→CO2 (Molar masses: C = 12.01 g/mol, CO₂ = 44.01 g/mol)

4. Percentage Yield Problems

10. A reaction has a theoretical yield of 15.0 g but produces only 10.0 g of product. What is
the percentage yield?

11. When 20.0 g of CaCO₃ is decomposed, 8.5 g of CaO is obtained. If the theoretical yield
of CaO is 11.2 g, what is the percentage yield?
Given:
CaCO3→CaO+CO2

(Molar masses: CaCO₃ = 100.09 g/mol, CaO = 56.08 g/mol)

5. Empirical and Molecular Formula Problems

12. A compound contains 40.0% carbon, 6.7% hydrogen, and 53.3% oxygen by mass.
What is its empirical formula?
(Atomic masses: C = 12.01 g/mol, H = 1.008 g/mol, O = 16.00 g/mol)

13. A compound has an empirical formula of CH₂O and a molar mass of 180 g/mol. What
is its molecular formula?

14. A hydrocarbon contains 85.6% carbon and 14.4% hydrogen by mass. If the molar mass
of the compound is 56 g/mol, determine its molecular formula.
(Atomic masses: C = 12.01 g/mol, H = 1.008 g/mol)

6. Miscellaneous Stoichiometry Problems

15. How many liters of CO₂ gas (at STP) are produced from the complete combustion of
5.0 g of propane (C₃H₈)?
Given:
C3H8+5O2→3CO2+4H2O

(Molar mass of C₃H₈ = 44.10 g/mol, 1 mole of gas at STP = 22.4 L)

Chemical kinetics

 Chemical Kinetics – The branch of chemistry that studies the speed or rate of chemical
reactions, the factors affecting them, and the reaction mechanisms.

 Rate of Reaction – The change in concentration of reactants or products per unit time,
usually expressed in mol/L·s. It indicates how fast a reaction proceeds.

 Appearing Rate – The rate at which a product is formed in a chemical reaction, expressed
as the increase in concentration of the product per unit time.

 Disappearing Rate – The rate at which a reactant is consumed in a chemical reaction,

expressed as the decrease in concentration of the reactant per unit time.

 First-Order Reaction – A reaction in which the rate is directly proportional to the

concentration of one reactant. The rate law is expressed as:
Rate=k[A] where k is the rate constant and [A] is the reactant concentration.

 Second-Order Reaction – A reaction in which the rate depends on the square of the
concentration of one reactant or the product of two reactant concentrations. The rate law can be:
Rate=k[A]^2 or
Rate=k[A][B]

 Zero-Order Reaction – A reaction in which the rate is independent of the concentration of

the reactant(s). The rate law is:
Rate=k
meaning the reaction proceeds at a constant rate regardless of reactant concentration.
Collision theory

Collision theory explains how and why chemical reactions occur based on the behavior of
reacting particles. The key points are:

1. Particles Must Collide – Reactant molecules must physically collide for a reaction to
occur.
2. Proper Orientation – The colliding molecules must be oriented correctly for bonds to
break and form new products.
3. Sufficient Energy (Activation Energy, Ea) – The colliding molecules must have
enough kinetic energy to overcome the activation energy barrier and initiate the reaction.
4. Effective vs. Ineffective Collisions – Only collisions that meet both the orientation and
energy requirements lead to a successful reaction (effective collisions), while others do
not result in product formation.
5. Temperature Effect – Higher temperatures increase molecular speed, leading to more
frequent and energetic collisions, thus increasing the reaction rate.
6. Concentration Effect – Higher reactant concentration increases the frequency of
collisions, making reactions faster.
7. Catalysts Lower Activation Energy – A catalyst provides an alternative reaction
pathway with a lower activation energy, increasing the number of successful collisions
and speeding up the reaction.

Electrochemistry

1. Electrochemistry – The branch of chemistry that studies the relationship between

electricity and chemical reactions, including the generation of electrical energy from
redox reactions and the use of electricity to drive chemical changes.
2. Electrochemical Cell – A system that converts chemical energy into electrical energy (or
vice versa) through redox reactions, consisting of two electrodes and an electrolyte.
3. Electrolytic Cell – A type of electrochemical cell that uses an external electrical source
to drive a non-spontaneous redox reaction.
4. Anode – The electrode where oxidation occurs in an electrochemical cell. In galvanic
cells, it is the negative electrode, whereas in electrolytic cells, it is the positive electrode.
5. Cathode – The electrode where reduction occurs in an electrochemical cell. In galvanic
cells, it is the positive electrode, whereas in electrolytic cells, it is the negative electrode.
6. Redox Reaction – A chemical reaction that involves the transfer of electrons, consisting
of two half-reactions: oxidation (loss of electrons) and reduction (gain of electrons).

Five Definitions of Reduction and Oxidation:

Reduction:
 1. Gain of electrons
2. Decrease in oxidation state
3. Gain of hydrogen
4. Loss of oxygen
5. Occurs at the cathode

Oxidation:

1. Loss of electrons
2. Increase in oxidation state
3. Loss of hydrogen
4. Gain of oxygen
5. Occurs at the anode
6. Oxidizing Agent – A substance that accepts electrons and causes oxidation of another
substance while itself getting reduced.
7. Reducing Agent – A substance that donates electrons and causes reduction of another
substance while itself getting oxidized.
8. Counter Ions – Ions that balance the charge in a solution but do not participate in the
redox reaction (e.g., Na⁺ in NaCl solution).
9. Spectator Ions – Ions that remain unchanged in a reaction and do not take part in the
redox process (e.g., Cl⁻ in the reaction of NaCl and AgNO₃).
10. Net Ionic Equation – A simplified chemical equation that shows only the species
directly involved in the reaction, omitting spectator ions.
Conditions for spontaneous and non-spontaneous cell

Spontaneous Reaction (Galvanic/Voltaic Cell)

A reaction in an electrochemical cell is spontaneous when:

 The cell potential is positive:Ecell∘>0

 The Gibbs free energy change is negative:ΔG<0
 Equilibrium constant k>>>>>1
 The reaction occurs without an external power source, generating electrical energy.

2. Non-Spontaneous Reaction (Electrolytic Cell)

A reaction in an electrochemical cell is non-spontaneous when:

 The cell potential is negative:Ecell∘<0

 The Gibbs free energy change is positive:ΔG>0
 Equilibrium constant k<<<<<1
 An external power source is required to drive the reaction.

Nuclear chemistry

Nuclear chemistry is the branch of chemistry that deals with reactions involving atomic nuclei,
including nuclear decay, fission, fusion, and radiation. It focuses on changes in nuclear structure
and energy transformations.

Characteristics of Nuclear Reactions

1. Involves the Nucleus – Unlike chemical reactions, nuclear reactions affect the nucleus,
changing elements.
2. Mass-Energy Conversion – Follows Einstein’s equation E=mc2E=mc2, where some
mass converts into energy.
3. High Energy Release – Nuclear reactions release far more energy than chemical
reactions.
4. Transmutation – Atoms of one element can transform into another due to nuclear
changes.
5. Not Affected by External Conditions – Unlike chemical reactions, nuclear reactions are
unaffected by temperature, pressure, or catalysts.

Types of Nuclear Reactions

1. Radioactive Decay (Natural Transmutation)

o A nucleus emits radiation (alpha, beta, gamma) to become more stable.
2. Nuclear Fission
 o A heavy nucleus splits into smaller nuclei, releasing energy.
3. Nuclear Fusion
o Two light nuclei combine to form a heavier nucleus, releasing massive energy.
o Example (Sun’s Fusion Reaction):12H+13H→24He+01n12H+13H→24He+01n
4. Artificial Transmutation
o Bombarding a nucleus with high-energy particles (e.g., protons, neutrons) to
create a new element.
5. Electron Capture
o A nucleus absorbs an inner electron, converting a proton into a neutron.

Nuclear Stability and How to Determine It

1. Neutron-to-Proton Ratio (N/ZN/Z)

o Stable nuclei have an optimal N/ZN/Z ratio:
 Light elements (Z<20): N/Z≈1
 Heavy elements (Z>20): N/Z≈1.5
2. Binding Energy
o Higher binding energy per nucleon indicates greater stability.
3. Magic Numbers
o Nuclei with magic numbers (2, 8, 20, 28, 50, 82, 114, 126) of protons or
neutrons tend to be more stable.

Radioactive elements

Alpha particle

Alpha Particle (α)

 A helium nucleus consisting of 2 protons and 2 neutrons.

 Symbol: 4He2 or α
 Charge: +2
 Mass: 4 amu.
 Penetration Power: Low (stopped by paper or skin).

Beta particle

Beta Particle (β−)

 A high-speed electron emitted during beta decay.

 Symbol: 0e-1oe 0 β−1
 Charge: −1.
 Mass: ~0.00055 amu (negligible).
 Penetration Power: Moderate (stopped by aluminum foil).
Positron particle

Positron Particle (β+$

 A positively charged electron (antimatter of an electron).

 Symbol: β+.
 Charge: +1
 Mass: ~0.00055 amu.
 Penetration Power: Moderate.

Proton particle

Proton Particle (pp)

 A positively charged subatomic particle found in the nucleus.

 Symbol: 1P1 or 1H1
 Charge: +1
 Mass: 1.007 amu.
 Penetration Power: High energy needed to accelerate free protons.

Neutron particle

 Neutron Particle (nn)

 A neutral subatomic particle in the nucleus.

 Symbol: 0n1
 Charge: 00 (neutral).
 Mass: 1.008 amu.
 Penetration Power: Very high (stopped by thick concrete or water).

 Gamma Ray (γ)

 A high-energy electromagnetic wave emitted from an excited nucleus.

 Symbol: γ
 Charge: 00.
 Mass: 00 (pure energy).
 Penetration Power: Very high (stopped by lead or thick concrete).

Questions

What is nuclear chemistry?

A) The study of chemical reactions involving electrons

B) The study of changes in atomic nuclei, including radiation and energy transformations
C) The study of molecules in chemical bonding
D) The study of organic compounds

Answer: B) The study of changes in atomic nuclei, including radiation and energy
transformations

2. Which of the following is a characteristic of nuclear reactions?

A) They involve changes in electron configurations

B) They are affected by temperature and pressure
C) They involve changes in the nucleus and release large amounts of energy
D) They only occur in radioactive elements

Answer: C) They involve changes in the nucleus and release large amounts of energy

3. Which type of nuclear radiation has the highest penetrating power?

A) Alpha particles
B) Beta particles
C) Gamma rays
D) Positron particles

Answer: C) Gamma rays

4. What is the charge of an alpha particle (αα)?

A) +1
B) -1
C) +2
D) 0

Answer: C) +2

5. What happens in a nuclear fission reaction?

A) A heavy nucleus splits into smaller nuclei

B) Two light nuclei combine to form a heavier nucleus
C) An electron is captured by a nucleus
D) A nucleus emits only gamma radiation

Answer: A) A heavy nucleus splits into smaller nuclei

7. What is the function of neutrons in nuclear stability?

A) They stabilize the nucleus by reducing repulsion between protons

B) They increase the energy of the nucleus
C) They cause the nucleus to become unstable
D) They convert protons into electrons

Answer: A) They stabilize the nucleus by reducing repulsion between protons

8. What is a magic number in nuclear chemistry?

A) The atomic number of stable elements

B) A number of protons or neutrons that leads to extra stability
C) The number of electrons needed to complete an orbital
D) The number of isotopes of an element

Answer: B) A number of protons or neutrons that leads to extra stability

9. In an electrochemical cell, a reaction is spontaneous when the cell potential

(Ecell∘Ecell∘) is:

A) Positive
B) Negative
C) Zero
D) Dependent on temperature

Answer: A) Positive

11. What type of nuclear reaction occurs in the sun?

A) Fission
B) Fusion
C) Alpha decay
D) Beta decay

Answer: B) Fusion

12. Which of the following particles has no mass and no charge?

A) Alpha particle
B) Beta particle
C) Gamma ray
D) Neutron

Answer: C) Gamma ray

13. Which nuclear reaction is responsible for producing energy in nuclear power
plants?

A) Alpha decay
B) Fusion
C) Fission
D) Beta decay

Answer: C) Fission

14. In beta decay, what happens to a neutron?

A) It breaks into a proton and an electron

B) It combines with a proton to form an alpha particle
C) It converts into two electrons
D) It turns into an electron and a gamma ray

Answer: A) It breaks into a proton and an electron

15. Which of the following particles has a positive charge but a very small mass?
A) Neutron
B) Positron
C) Proton
D) Alpha particle

Answer: B) Positron

16. What does a net ionic equation represent?

A) All species, including spectator ions

B) Only the ions that undergo change in a reaction
C) The complete formula of all reactants and products
D) Only the counter ions in a reaction

Answer: B) Only the ions that undergo change in a reaction

17. Which of the following is the best indicator of nuclear stability?

A) The atomic radius

B) The neutron-to-proton ratio
C) The number of electrons
D) The electronegativity of the atom

Answer: B) The neutron-to-proton ratio

18. Which type of decay does carbon-14 undergo?

A) Alpha decay
B) Beta decay
C) Positron emission
D) Gamma emission

Answer: B) Beta decay

19. The energy released in nuclear reactions follows which equation?

A) E=mc2
B) PV=nRTPV=nRT
C) ΔG=ΔH−TΔS

D) F=m

Answer: A) E=mc2

20. Which of the following statements is true about an electrolytic cell?

A) It generates electrical energy spontaneously

B) It requires an external power source to drive a non-spontaneous reaction
C) It only works with elements in liquid form
D) It cannot undergo redox reactions

Answer: B) It requires an external power source to drive a non-spontaneous reaction

21. What is the symbol for a neutron?

A) 11p11p
B) −10e−10e
C) 01n01n
D) 24He24He

Answer: C) 01n01n

22. Which type of nuclear reaction involves the combination of two nuclei?

A) Fission
B) Fusion
C) Alpha decay
D) Beta decay

Answer: B) Fusion

23. The process of an atom capturing an electron and converting a proton into a
neutron is called:
A) Alpha decay
B) Electron capture
C) Beta emission
D) Gamma radiation

Answer: B) Electron capture

24. Which element is most commonly used in nuclear fission reactions?

A) Hydrogen
B) Carbon
C) Uranium
D) Oxygen

Answer: C) Uranium

25. If a nucleus emits an alpha particle, how does its atomic number change?

A) Increases by 2
B) Decreases by 2
C) Increases by 4
D) Remains unchanged

Answer: B) Decreases by 2

Chemical bonding and its theories

VBT, Lewis, Sidgwick Powell

According to Valence Bond Theory (VBT), covalent bonds are formed by:

A) Overlapping of atomic orbitals

B) Transfer of electrons
C) Complete loss of valence electrons
D) Weak Van der Waals forces

Answer: A) Overlapping of atomic orbitals

2. In Lewis theory, a covalent bond is formed when atoms:

A) Share electron pairs

B) Transfer electrons completely
C) Donate electrons to the nucleus
D) Share protons

Answer: A) Share electron pairs

3. According to the Sidgwick-Powell theory, the shape of a molecule depends on:

A) The atomic number of the central atom

B) The number of bonding and lone pairs around the central atom
C) The molecular mass
D) The number of neutrons in the nucleus

Answer: B) The number of bonding and lone pairs around the central atom

4. Which of the following is a limitation of Valence Bond Theory (VBT)?

A) It fails to explain molecular shape completely

B) It explains resonance well
C) It predicts bond lengths accurately
D) It considers electron delocalization

Answer: A) It fails to explain molecular shape completely

5. In the Lewis structure of CO2CO2, how many lone pairs are present on the
oxygen atoms?

A) 1
B) 2
C) 3
D) 4

Answer: B) 2
6. Which type of orbital overlap forms a sigma (σσ) bond?

A) Sidewise overlap of p orbitals

B) Head-on overlap of s or p orbitals
C) Overlap of d orbitals only
D) Overlap of f orbitals

Answer: B) Head-on overlap of s or p orbitals

7. In Lewis structures, the octet rule states that atoms tend to:

A) Gain, lose, or share electrons to achieve eight valence electrons

B) Transfer electrons to form ionic bonds
C) Maintain the same number of valence electrons
D) Form metallic bonds

Answer: A) Gain, lose, or share electrons to achieve eight valence electrons

8. According to Sidgwick-Powell theory, a molecule with 4 bond pairs and 1 lone

pair has what shape?

A) Tetrahedral
B) Trigonal pyramidal
C) Trigonal bipyramidal
D) Bent

Answer: C) Trigonal bipyramidal

9. The overlapping of two p orbitals sideways forms:

A) A sigma (σσ) bond

B) A pi (ππ) bond
C) A delta (δδ) bond
D) A hydrogen bond

Answer: B) A pi (ππ) bond

10. In Valence Bond Theory, which type of hybridization occurs in methane
(CH4CH4)?

A) sp
B) sp2
C) sp3
D) d2sp3d

Answer: C) sp3

11. What is the major limitation of Lewis theory?

A) It does not explain bond polarity

B) It does not predict the shape of molecules
C) It cannot explain the existence of ionic bonds
D) It cannot determine molecular mass

Answer: B) It does not predict the shape of molecules

12. Which molecule violates the octet rule?

A) CO2
B) NH3
C) SF6
D) CH4

Answer: C) SF6

13. Which of the following is NOT a postulate of Valence Bond Theory?

A) Covalent bonds form due to orbital overlap

B) Bond strength depends on the extent of overlap
C) Electrons remain localized between atoms
D) Molecular orbitals form from atomic orbitals

Answer: D) Molecular orbitals form from atomic orbitals

15. According to Lewis theory, why do noble gases rarely form bonds?

A) They have high atomic mass

B) They have a full valence shell
C) They contain unpaired electrons
D) They are unstable

Answer: B) They have a full valence shell

16. Which molecule has an angular (bent) shape according to the Sidgwick-
Powell theory?

A) CO2
B) H2O
C) BF3
D) CH4

Answer: B) H2O

17. Which type of hybridization is found in ethene (C2H4C2H4)?

A) sp3sp3
B) sp2
C) spsp
D) d2sp3d

Answer: B) sp2sp2

18. In Lewis structures, double bonds contain:

A) One sigma bond and two pi bonds

B) Two sigma bonds
C) One sigma bond and one pi bond
D) Two pi bonds

Answer: C) One sigma bond and one pi bond

19. What determines bond strength in Valence Bond Theory?

A) Number of unpaired electrons

B) Extent of orbital overlap
C) Type of molecular orbitals
D) Type of atom

Answer: B) Extent of orbital overlap

20. Which shape does a molecule with sp2sp2 hybridization adopt?

A) Tetrahedral
B) Trigonal planar
C) Octahedral
D) Linear

Answer: B) Trigonal planar

21. Which of the following statements about a pi bond (π) is true?

A) It is weaker than a sigma bond

B) It forms by head-on overlap of orbitals
C) It allows free rotation of bonded atoms
D) It forms in single bonds

Answer: A) It is weaker than a sigma bond

22. Which molecule follows the octet rule?

A) NO2
B) ClO3−
C) NH3NH3
D) PCl5

Answer: C) NH3

23. Which concept is introduced in Sidgwick-Powell theory?

A) Valence shell electron pair repulsion (VSEPR)
B) Molecular orbital theory
C) Crystal field theory
D) Electron cloud theory

Answer: A) Valence shell electron pair repulsion (VSEPR)

24. What is the hybridization of the central atom in PCl5PCl5?

A) sp2
B) sp3
C) sp3d
D) sp3d2

Answer: C) sp3d

25. Which of the following molecules has a trigonal planar shape?

A) NH3
B) BF3
C) H2O
D) XeF4

Answer: B) BF3

What does VSEPR theory primarily explain?

A) The formation of covalent bonds

B) The shape of molecules based on electron pair repulsion
C) The movement of electrons in an electric field
D) The color of transition metal compounds

Answer: B) The shape of molecules based on electron pair repulsion

2. According to VSEPR theory, electron pairs around the central atom:

A) Attract each other
B) Repel each other and arrange themselves as far apart as possible
C) Stay fixed in one position
D) Form only single bonds

Answer: B) Repel each other and arrange themselves as far apart as possible

3. What is the molecular geometry of a molecule with two bonding pairs

and zero lone pairs?

A) Trigonal planar
B) Bent
C) Linear
D) Tetrahedral

Answer: C) Linear

4. Which of the following molecules has a bent shape according to VSEPR

theory?

A) CO2
B) H2O
C) BF3
D) CH4

Answer: B) H2O

5. A molecule with four bonding pairs and zero lone pairs around the central
atom will have which shape?

A) Trigonal planar
B) Linear
C) Tetrahedral
D) Octahedral

Answer: C) Tetrahedral
6. What is the molecular shape of ammonia (NH3NH3) according to VSEPR
theory?

A) Trigonal pyramidal
B) Bent
C) Trigonal planar
D) Tetrahedral

Answer: A) Trigonal pyramidal

7. The bond angle in a perfect tetrahedral molecule is approximately:

A) 90°
B) 109.5°
C) 120°
D) 180°

Answer: B) 109.5°

8. Which factor determines the shape of a molecule in VSEPR theory?

A) The number of valence electrons of the central atom

B) The number of bonding and lone pairs around the central atom
C) The atomic number of the central atom
D) The mass of the molecule

Answer: B) The number of bonding and lone pairs around the central atom

9. What is the shape of a molecule with three bonding pairs and one lone pair
around the central atom?

A) Tetrahedral
B) Trigonal pyramidal
C) Trigonal planar
D) Linear

Answer: B) Trigonal pyramidal

10. Which of the following molecules has a trigonal bipyramidal geometry?

A) CH4
B) NH3
C) PCl5
D) H2O

Answer: C) PCl5

11. What is the molecular shape of sulfur hexafluoride (SF6SF6)?

A) Trigonal bipyramidal
B) Octahedral
C) Tetrahedral
D) Square planar

Answer: B) Octahedral

13. What happens to bond angles when lone pairs are present in a molecule?

A) They increase due to stronger repulsions

B) They decrease due to stronger repulsions from lone pairs
C) They remain unchanged
D) They become 180° in all cases

Answer: B) They decrease due to stronger repulsions from lone pairs

14. Which of the following molecules has a square planar shape?

A) XeF4
B) SF6
C) CH4
D) BF3

Answer: A) XeF4
Acid, Bases, pH, pOH

 Acid – A substance that donates protons (H+) or accepts electron pairs in a reaction.

 Acidity – The measure of an acid's strength, typically determined by its ability to donate
protons (H+H+) or lower the pH of a solution.

 Arrhenius Acid – A substance that increases the concentration of H+ (or H3O+H3O+) ions
in an aqueous solution. Example: HCl dissociates to form H+ and Cl−

 Brønsted Acid – A substance that donates a proton (H+H+) to another substance in a

chemical reaction. Example: H2SO4 donates H+ to water.

 Arrhenius Base – A substance that increases the concentration of OH−OH− (hydroxide ions)
in an aqueous solution. Example: NaOH dissociates to form Na+ and OH−.

 Brønsted Base – A substance that accepts a proton (H+H+) from another substance in a
chemical reaction. Example: NH3 accepts H+ to form NH4+

 Lewis Acid – A substance that accepts an electron pair during a reaction.

Example: AlCl3 accepts electron pairs.

 Lewis Base – A substance that donates an electron pair during a reaction.

Example: NH3NH3 donates a lone pair to form bonds.

 Conjugate Acid – The species formed when a Brønsted base gains a proton (H+H+).
Example: NH4+ is the conjugate acid of NH3

 Conjugate Base – The species formed when a Brønsted acid loses a proton (H+H+).
Example: Cl− is the conjugate base of HCl
Questions

According to Arrhenius theory, an acid is a substance that:

A) Increases OH− concentration in solution
B) Increases H+ (or H3O+) concentration in solution
C) Accepts electron pairs
D) Donates an electron pair

Answer: B) Increases H+ (or H3O+) concentration in solution

2. Which of the following is an example of a Brønsted-Lowry acid?

A) NH3NH3
B) OH−OH−
C) H2SO4H2SO4
D) Cl−Cl−

Answer: C) H2SO4H2SO4

3. What is the main difference between a Lewis acid and a Brønsted acid?

A) A Lewis acid donates H+ while a Brønsted acid accepts an electron pair

B) A Lewis acid accepts an electron pair, while a Brønsted acid donates H+H+
C) A Lewis acid donates an electron pair, while a Brønsted acid donates H+H+
D) A Lewis acid accepts H+ while a Brønsted acid donates an electron pair

Answer: B) A Lewis acid accepts an electron pair, while a Brønsted acid donates H+H+

4. Which of the following is a characteristic of an Arrhenius base?

A) It donates protons (H+)

B) It increases OH−concentration in water
C) It accepts electron pairs
D) It increases hydrogen gas in a reaction

Answer: B) It increases OH− concentration in water

5. Which of the following is a Lewis base?

A) BF3
B) NH3
C) HCl
D) SO3

Answer: B) NH3

6. Which of the following pairs correctly represents a conjugate acid-base pair?

A) H2O and H3O+

B) HCl and OH−
C) NH3NH3 and Cl−
D) NaOH and NaCl

Answer: A) H2O and H3O+

7. What is the conjugate base of H2SO4?

A) HSO4−
B) SO42−
C) H3O+
D) H2O

Answer: A) HSO4-

8. Which of the following is an example of an Arrhenius base?

A) HCl
B) NaOH
C) CO2
D) NH4+

Answer: B) NaOH

9. Which of the following statements is true about acidity?

A) Acidity is determined by the mass of an acid
B) Acidity is a measure of how much H+H+ an acid can donate
C) Acidity depends only on the number of valence electrons
D) Acidity remains the same regardless of concentration

Answer: B) Acidity is a measure of how much H+H+ an acid can donate

10. A conjugate base is formed when:

A) An acid gains a proton

B) A base gains an electron pair
C) An acid loses a proton
D) A base loses a hydroxide ion

Answer: C) An acid loses a proton

11. Which of the following substances acts as both a Brønsted acid and a
Brønsted base (amphoteric substance)?

A) HClHCl
B) NaOHNaOH
C) H2OH2O
D) CO2CO2

Answer: C) H2OH2O

12. Which of the following is NOT a property of Lewis acids?

A) They accept electron pairs

B) They can have incomplete octets
C) They donate protons
D) Examples include AlCl3 and BF3

Answer: C) They donate protons

13. What is the conjugate acid of NH3NH3?

A) NH2−
B) NH4+
C) N2
D) H2O

Answer: B) NH4+

14. Which of the following represents a Lewis acid-base reaction?

A) NaOH+HCl→NaCl+H2ON

B) NH3+BF3→NH3BF3

C) HCl→H+Cl−

D) H2O+H2SO4→H3O++HSO4−

Answer: B) NH3+BF3→NH3BF3

15. Which acid-base theory is most general and includes both Brønsted and
Arrhenius definitions?

A) Arrhenius Theory
B) Brønsted-Lowry Theory
C) Lewis Theory
D) Avogadro's Law

Answer: C) Lewis Theory

Calculations

What is the pH of a solution with a hydrogen ion concentration of 1.0×10−3M?

(A) 1
(B) 3
(C) 7
(D) 10
2. What is the pOH of a solution with a hydroxide ion concentration of 2.0×10−5?

(A) 4.7
(B) 5.3
(C) 9.7
(D) 10.3

3. A solution has a pOH of 4. What is its pH?

(A) 4
(B) 7
(C) 10
(D) 11

4. If a solution has a pH of 2, what is its hydrogen ion concentration [H+][H+]?

(A) 1.0×10−2M
(B) 1.0×10−4M
(C) 5.0×10−3M
(D) 2.0×10−2M

5. A solution has a hydroxide ion concentration of 1.0×10−8M1.0×10−8M. What

is its pH?

(A) 6
(B) 7
(C) 8
(D) 10

6. What is the pH of a 0.01 M solution of NaOH (a strong base)?

(A) 2
(B) 7
(C) 12
(D) 14
Answer: (C) 12

7. A 0.002 M solution of HCl (a strong acid) is prepared. What is its pH?

(A) 2.7
(B) 3.0
(C) 3.3
(D) 4.0

8. The pH of a solution is 9. What is its hydrogen ion concentration?

(A) 1.0×10−5M
(B) 1.0×10−9M
(C) 1.0×10−11
(D) 1.0×10−14

9. A 0.01 M acetic acid (CH3COOHCH3COOH) solution has a Ka of 1.8×10−5,

What is its approximate pH?

(A) 2.4
(B) 3.0
(C) 3.5
(D) 4.1

10. If the pKa of an acid is 4.75, what is its Ka?

(A) 1.8×10−5
(B) 5.6×10−4
(C) 1.0×10−3
(D) 2.5×10−6

11. What is the pH of a 0.025 M solution of KOH?

(A) 11.4
(B) 12.0
(C) 12.4
(D) 13.5

12. The pH of a buffer solution containing 0.1 M acetic acid and 0.1 M sodium
acetate is closest to:

(A) 3.8
(B) 4.8
(C) 5.8
(D) 6.8

13. If the pKa of a weak acid is 5.2, what is the pH of a buffer solution in which
the ratio of base to acid is 10:1?

(A) 4.2
(B) 5.2
(C) 6.2
(D) 7.2

14. A solution has a pH of 7. What is its hydroxide ion concentration?

(A) 1.0×10−7M

(B) 1.0×10−14M

(C) 1.0×10−5M

(D) 1.0×10−12M
15. The hydroxide ion concentration of a solution is 3.0×10−4M3.0×10−4M.
What is its pH?

(A) 10.5
(B) 11.2
(C) 11.5
(D) 12.5

Thermochemistry and thermodynamics

1. Thermochemistry – The study of heat energy associated with chemical reactions and
phase changes.
2. Thermodynamics – The branch of physics and chemistry that deals with energy
transformations, particularly involving heat, work, and internal energy.
3. Enthalpy (H) – A thermodynamic property representing the total heat content of a
system at constant pressure.
4. Enthalpy of Formation (ΔHf∘) – The heat change when one mole of a substance is
formed from its elements in their standard states.
5. Enthalpy of Combustion (ΔHc°) – The heat released when one mole of a substance
completely burns in oxygen under standard conditions.
6. First Law of Thermodynamics – States that energy cannot be created or destroyed, only
transferred or transformed. Mathematically, it is expressed as:

ΔU=Q+W

where ΔU is the change in internal energy, Q is heat added to the system, and W is work
done on the system.

7. Closed System – A system that allows energy exchange with its surroundings but not
matter transfer.
8. Open System – A system that allows both energy and matter to be exchanged with its
surroundings.
9. Isolated System – A system that does not exchange either energy or matter with its
surroundings.
10. Adiabatic Environment – A system where no heat exchange occurs with the
surroundings (Q=0).
11. Isochoric Environment – A system where volume remains constant (ΔV=0ΔV=0),
meaning no work is done (W=).
12. Isothermal Environment – A system where temperature remains constant (ΔT=0ΔT=0),
so internal energy (ΔU) remains unchanged.
13. Isobaric Environment – A system where pressure remains constant (ΔP=0ΔP=0),
meaning the heat change (Q) equals the enthalpy change (ΔH).
 Second Law of Thermodynamics – States that the total entropy of an isolated system
always increases over time. It implies that natural processes tend to move towards greater
disorder and that heat cannot spontaneously flow from a colder body to a hotter one.

1. Entropy (S) – A measure of the disorder or randomness in a system. Higher entropy

means greater disorder.
2. Gibbs Free Energy (G) – A thermodynamic quantity that indicates the maximum useful
work obtainable from a system at constant temperature and pressure. It determines the
spontaneity of a reaction.
3. Formula Linking Entropy, Enthalpy, and Gibbs Free Energy –

ΔG=ΔH−TΔS

where:

o ΔG = Gibbs free energy change

o ΔH = Enthalpy change
o T = Temperature (in Kelvin)
o ΔS = Entropy change

A reaction is:

 Spontaneous if ΔG<0 (negative)

 Non-spontaneous if ΔG>0(positive)
 At equilibrium if ΔG=0

Conceptual questions

What does the First Law of Thermodynamics state?

(A) Energy can be created and destroyed.

(B) Energy cannot be created or destroyed, only transferred or transformed.
(C) Entropy always decreases in an isolated system.
(D) The total energy of the universe is increasing.

Answer: (B)

2. Which of the following correctly represents the First Law of Thermodynamics?

(A) ΔU=Q+W
(B) ΔU=Q−W
(C) ΔS=Q/T
(D) ΔG=ΔH−TΔS

Answer: (A) ΔU=Q+W

3. In an adiabatic process, which quantity remains constant?

(A) Pressure
(B) Temperature
(C) Heat exchange
(D) Volume

Answer: (C) Heat exchange (Q=0)

4. Which of the following is an example of an open system?

(A) A sealed soda bottle

(B) A thermos flask
(C) A boiling pot of water
(D) A bomb calorimeter

Answer: (C) A boiling pot of water

5. What is enthalpy (HH)?

(A) A measure of disorder in a system

(B) The heat content of a system at constant pressure
(C) The energy required to break chemical bonds
(D) The amount of work done by a system at constant temperature

Answer: (B) The heat content of a system at constant pressure

6. Which of the following correctly defines Gibbs Free Energy (GG)?

(A) The energy required to heat a substance by 1°C

(B) The total energy of a system at constant volume
(C) The maximum useful work obtained from a system at constant pressure and temperature
(D) The heat change in a system at constant volume

Answer: (C) The maximum useful work obtained from a system at constant pressure and
temperature

7. What does the Second Law of Thermodynamics state?

(A) Energy cannot be created or destroyed.

(B) Every spontaneous process increases the total entropy of the universe.
(C) The total energy of an isolated system remains constant.
(D) The change in Gibbs Free Energy determines spontaneity.

Answer: (B) Every spontaneous process increases the total entropy of the universe.

8. Which of the following is a correct equation for Gibbs Free Energy?

(A) ΔG=ΔH+TΔS
(B) ΔG=ΔH−TΔS
(C) ΔG=TΔS−ΔH
(D) ΔG=ΔS−TΔH

Answer: (B) ΔG=ΔH−TΔS

9. If ΔG<0, what can be said about the reaction?

(A) It is non-spontaneous.
(B) It is at equilibrium.
(C) It is spontaneous.
(D) The reaction cannot proceed.

Answer: (C) It is spontaneous.

10. What is the entropy (SS) of a perfectly ordered crystalline substance at absolute zero (0
K)?

(A) 0 J/K
(B) 1 J/K
(C) A large positive value
(D) A large negative value

Answer: (A) 0 J/K (from the Third Law of Thermodynamics)

11. Which of the following processes will most likely increase entropy (ΔSΔS)?

(A) Freezing of water

(B) Condensation of steam
(C) Dissolving salt in water
(D) Formation of a solid from a gas lol

Answer: (C) Dissolving salt in water

12. What remains constant in an isobaric process?

(A) Pressure
(B) Temperature
(C) Volume
(D) Heat

Answer: (A) Pressure

13. In which type of system can energy but not matter be exchanged with the
surroundings?

(A) Open system

(B) Closed system
(C) Isolated system
(D) None of the above

Answer: (B) Closed system

14. The enthalpy of combustion is defined as:

(A) The heat required to break bonds in a molecule
(B) The heat released when one mole of a substance combusts completely in oxygen
(C) The heat absorbed when a substance melts
(D) The energy needed to form one mole of a compound

Answer: (B) The heat released when one mole of a substance combusts completely in
oxygen

15. If a process occurs in an isochoric environment, what remains constant?

(A) Temperature
(B) Volume
(C) Pressure
(D) Heat

Answer: (B) Volume

16. A reaction is spontaneous at all temperatures when:

(A) ΔH>0, ΔS<0

(B) ΔH<0, ΔS>0
(C) ΔH>0, ΔS>0
(D) ΔH<0, ΔS<0

Answer: (B) ΔH<0, ΔS>0

17. What happens in an isothermal process?

(A) Temperature changes

(B) Pressure remains constant
(C) Heat transfer occurs but temperature remains constant
(D) The system is isolated

Answer: (C) Heat transfer occurs but temperature remains constant

18. What is the main criterion for determining whether a reaction is spontaneous?
(A) Enthalpy (ΔH)
(B) Entropy (ΔS)
(C) Internal Energy (ΔU)
(D) Gibbs Free Energy (ΔG)

Answer: (D) Gibbs Free Energy (ΔGΔG)

19. A system where neither mass nor energy is exchanged with the surroundings is called:

(A) Open system

(B) Closed system
(C) Isolated system
(D) Adiabatic system

Answer: (C) Isolated system

20. If a reaction has ΔG=0, what does it mean?

(A) The reaction is spontaneous.

(B) The reaction is non-spontaneous.
(C) The system is at equilibrium.
(D) The reaction cannot proceed.

Answer: (C) The system is at equilibrium.

Calculations

Given the reaction:

2H2+O2→2H2O

If the enthalpy change (ΔH) for the formation of water is -286 kJ/mol, what is the total enthalpy
change for the reaction?

(A) -572 kJ
(B) -286 kJ
(C) -143 kJ
(D) 572 kJ
**2. The enthalpy of combustion of methane (CH4) is -890 kJ/mol. How much heat is
released when 3 moles of methane are burned?

(A) -2670 kJ
(B) -890 kJ
(C) -1780 kJ
(D) 890 kJ

4. A chemical reaction absorbs 450 J of heat and does 150 J of work on the surroundings.
What is the change in internal energy (ΔUΔU)?

(A) 600 J
(B) 300 J
(C) 450 J
(D) -300 J

Answer: (B) 300 J

5. How much heat is required to raise the temperature of 500 g of water from 25°C to
75°C?

(Specific heat capacity of water = 4.18 J/g°C)

(A) 10450 J
(B) 52350 J
(C) 104500 J
(D) 41800 J

Answer: (B) 52350 J

6. The standard enthalpy of formation (ΔHf∘) of NH₃(g) is -46 kJ/mol. What is the
enthalpy change for the formation of 3 moles of NH₃?

(A) -138 kJ
(B) -46 kJ
(C) 138 kJ
(D) 46 kJ
7. If 200 kJ of heat is absorbed by 50 g of a metal with a specific heat capacity of 0.385
J/g°C, what is the temperature change (ΔT)?

(A) 10.4°C
(B) 103.9°C
(C) 38.5°C
(D) 104°C

8. Calculate the standard enthalpy change (ΔH∘) for the reaction:

CH4+2O2→CO2+2H2O

Given:
ΔHf∘ΔHf∘ values (in kJ/mol):

 CH4=−75
 CO2=−394
 H2O=−286

(A) -890 kJ
(B) -1050 kJ
(C) -802 kJ
(D) -680 kJ

Answer: (A) -890 kJ

9. If a reaction releases 250 kJ of heat and does 50 kJ of work on the surroundings, what is
the change in internal energy (ΔUΔU)?

(A) 300 kJ
(B) 200 kJ
(C) -200 kJ
(D) -300 kJ

Answer: (B) 200 kJ

10. Calculate the heat released when 20 g of propane (C₃H₈) combusts, given its enthalpy
of combustion is -2220 kJ/mol.

(Molar mass of C₃H₈ = 44 g/mol)

(A) -1110 kJ
(B) -1008 kJ
(C) -1200 kJ
(D) -2220 kJ

Answer: (B) -1008 kJ

11. How much heat is released when 10 g of ethanol (C₂H₅OH) combusts, given its
enthalpy of combustion is -1367 kJ/mol?

(Molar mass of C₂H₅OH = 46 g/mol)

(A) -297 kJ
(B) -250 kJ
(C) -420 kJ
(D) -1367 kJ

12. A 200 g piece of metal is heated to 100°C and then placed in 500 g of water at 25°C. The
final temperature is 30°C. If the specific heat of the metal is 0.4 J/g°C, what was the heat
lost by the metal?

(A) 560 J
(B) 800 J
(C) 400 J
(D) 1200 J

13. For the reaction:

A+B→C+D

If 100 g of A releases 450 kJ, how much heat is released if 250 g of A reacts?

(A) 1125 kJ
(B) 450 kJ
(C) 750 kJ
(D) 250 kJ
15. The enthalpy of fusion of water is 6.01 kJ/mol. How much heat is needed to melt 50 g of
ice?

(Molar mass of water = 18 g/mol)

(A) 10.5 kJ
(B) 16.7 kJ
(C) 8.35 kJ
(D) 25 kJ

Answer: (B) 16.7 kJ

Periodicity and periodic trends

Periodicity

Periodicity refers to the repeating patterns and trends in the properties of elements as they are
arranged in the periodic table according to their atomic numbers. These trends arise due to the
periodic recurrence of similar electronic configurations in elements of the same group or period.

2. Periodic Trends and Their Directions

Periodic trends describe how certain chemical and physical properties change systematically
across a period (left to right) and down a group (top to bottom) in the periodic table.

Major Periodic Trends:

Trend Across a Period (Left to Right) Down a Group (Top to Bottom)

Atomic Radius Decreases Increases
Ionization Energy Increases Decreases
Electronegativity Increases Decreases
Electron Affinity Becomes more negative (increases) Becomes less negative (decreases)
Metallic Character Decreases Increases
Non-metallic Character Increases Decreases

3. The 18 Groups in the Periodic Table

The periodic table consists of 18 vertical columns called groups. Each group has elements with
similar chemical properties due to their valence electron configurations.
Group Number Group Name
Group 1 Alkali Metals
Group 2 Alkaline Earth Metals
Group 3 Scandium Group (Part of Transition Metals)
Group 4 Titanium Group
Group 5 Vanadium Group
Group 6 Chromium Group
Group 7 Manganese Group
Group 8 Iron Group
Group 9 Cobalt Group
Group 10 Nickel Group
Group 11 Coinage Metals (Copper Group)
Group 12 Zinc Group
Group 13 Boron Group
Group 14 Carbon Group
Group 15 Nitrogen Group (Pnictogens)
Group 16 Oxygen Group (Chalcogens)
Group 17 Halogens
Group 18 Noble Gases

4. The 7 Periods in the Periodic Table

The horizontal rows in the periodic table are called periods. Each period corresponds to the
number of electron shells an element has. The elements within the same period do not
necessarily have similar chemical properties but follow periodic trends.

Period Number Range of Atomic Numbers Blocks Found in This Period

Period 1 1-2 s-block
Period 2 3 - 10 s-block, p-block
Period 3 11 - 18 s-block, p-block
Period 4 19 - 36 s-block, d-block, p-block
Period 5 37 - 54 s-block, d-block, p-block
Period 6 55 - 86 s-block, d-block, f-block (Lanthanides), p-block
Period 7 87 - 118 s-block, d-block, f-block (Actinides), p-block

Would you like more details on any specific topic, such as why these trends
occur or examples of elements in each group?

4o
How many periods are present in the modern periodic table?

(A) 6
(B) 7
(C) 8
(D) 9

Answer: (B) 7

2. Which of the following elements is in Group 2 of the periodic table?

(A) Sodium
(B) Magnesium
(C) Fluorine
(D) Argon

Answer: (B) Magnesium

3. Which group in the periodic table contains noble gases?

(A) Group 1
(B) Group 2
(C) Group 17
(D) Group 18

Answer: (D) Group 18

4. As you move from left to right across a period, atomic radius generally:

(A) Increases
(B) Decreases
(C) Remains constant
(D) Increases, then decreases

Answer: (B) Decreases

5. The elements in the same group of the periodic table have similar:
(A) Atomic numbers
(B) Number of valence electrons
(C) Atomic masses
(D) Electronegativity

Answer: (B) Number of valence electrons

6. Which of the following elements has the highest ionization energy?

(A) Cesium
(B) Fluorine
(C) Oxygen
(D) Neon

Answer: (D) Neon

7. The most electronegative element in the periodic table is:

(A) Hydrogen
(B) Carbon
(C) Oxygen
(D) Fluorine

Answer: (D) Fluorine

8. Moving down a group, atomic radius:

(A) Increases
(B) Decreases
(C) Remains constant
(D) Increases, then decreases

Answer: (A) Increases

9. Which group of elements is also known as alkali metals?

(A) Group 1
(B) Group 2
(C) Group 16
(D) Group 17

Answer: (A) Group 1

10. The number of valence electrons in elements of Group 16 is:

(A) 2
(B) 4
(C) 6
(D) 8

Answer: (C) 6

11. The horizontal rows in the periodic table are called:

(A) Groups
(B) Blocks
(C) Periods
(D) Families

Answer: (C) Periods

12. Which of the following trends increases from left to right across a period?

(A) Atomic radius

(B) Metallic character
(C) Ionization energy
(D) Reactivity of metals

Answer: (C) Ionization energy

13. What is the general trend of electron affinity across a period from left to right?

(A) Increases
(B) Decreases
(C) Remains the same
(D) Increases, then decreases
Answer: (A) Increases

14. Which of the following statements about Group 1 elements is true?

(A) They have low ionization energies

(B) They are highly electronegative
(C) They form covalent compounds
(D) They are noble gases

Answer: (A) They have low ionization energies

15. Which group in the periodic table contains halogens?

(A) Group 1
(B) Group 2
(C) Group 17
(D) Group 18

Answer: (C) Group 17

16. Which of the following properties decreases as you move down Group 17?

(A) Atomic radius

(B) Reactivity
(C) Metallic character
(D) Electronegativity

Answer: (D) Electronegativity

17. The number of valence electrons in noble gases (except helium) is:

(A) 2
(B) 4
(C) 6
(D) 8

Answer: (D) 8
18. Which block of the periodic table contains transition metals?

(A) s-block
(B) p-block
(C) d-block
(D) f-block

Answer: (C) d-block

19. The element with atomic number 20 belongs to which group and period?

(A) Group 1, Period 3

(B) Group 2, Period 4
(C) Group 16, Period 3
(D) Group 18, Period 4

Answer: (B) Group 2, Period 4

20. Which of the following statements is true about elements in the same period?

(A) They have the same number of valence electrons

(B) They have the same number of electron shells
(C) They have similar chemical properties
(D) They belong to the same group

Answer: (B) They have the same number of electron shells

21. Which of the following elements has the smallest atomic radius?

(A) Lithium
(B) Carbon
(C) Fluorine
(D) Neon

Answer: (D) Neon

22. Which of the following increases as you move down a group in the periodic table?

(A) Ionization energy

(B) Electronegativity
(C) Atomic radius
(D) Electron affinity

Answer: (C) Atomic radius

23. What is the general trend of metallic character in the periodic table?

(A) Increases across a period

(B) Decreases down a group
(C) Increases down a group
(D) Remains constant

Answer: (C) Increases down a group

24. Which of the following elements belongs to the p-block?

(A) Magnesium
(B) Titanium
(C) Sulfur
(D) Cesium

Answer: (C) Sulfur

25. The periodic law states that:

(A) The chemical properties of elements are a function of their atomic mass
(B) The properties of elements repeat periodically when arranged by atomic number
(C) The number of neutrons determines periodicity
(D) Elements in the same group have different chemical properties

Answer: (B) The properties of elements repeat periodically when arranged by atomic
number

Structure of solids

Definition and Characteristics of Solids

Solids are one of the three primary states of matter, along with liquids and gases. They have:

 A definite shape and fixed volume due to tightly packed particles.

 High density, rigidity, and hardness because of strong intermolecular forces.
 Molecular motion limited to small vibrations around fixed positions.

2. Classification of Solids

Solids are broadly classified into:

1. Crystalline Solids – Particles are arranged in a well-ordered geometric pattern.

2. Amorphous Solids – Particles lack an ordered structure and have no defined geometric
pattern.

3. Crystalline Solids

Crystalline solids have a highly ordered structure and exhibit anisotropy (direction-dependent
properties). They possess:

 Sharp melting points.

 Definite heat of fusion.
 Cleavage planes (specific planes along which they split).

Types of Crystalline Solids:

1. Ionic Solids – Composed of oppositely charged ions (e.g., NaCl, NH₄NO₃). They
have high melting points and conduct electricity in solution.
2. Molecular Solids – Held together by weak Van der Waals forces (e.g., ice, sugar, HCl).
They are soft, non-conductive, and have low melting points.
3. Network Covalent Solids – Atoms form continuous covalent bonds (e.g., diamond,
quartz). They are hard, non-conductive, and have very high melting points.
4. Metallic Solids – Composed of metal atoms in a "sea of electrons", making them good
conductors of electricity and heat (e.g., iron, copper, alloys like steel and bronze).

4. Amorphous Solids

Unlike crystalline solids, amorphous solids lack a regular structure and

exhibit isotropy (uniform properties in all directions).

 They gradually melt over a temperature range instead of having a sharp melting point.
 Examples: Glass, rubber, plastics, and gels.

Question

Which of the following is NOT a characteristic of solids?

A) Definite shape
B) Fixed volume
C) Free-moving particles
D) High density
Answer: C) Free-moving particles

2. The major classification of solids includes:

A) Amorphous and Crystalline

B) Ionic and Metallic
C) Polar and Non-polar
D) Rigid and Non-rigid
Answer: A) Amorphous and Crystalline

3. In amorphous solids, the particles lack:

A) Fixed volume
B) An orderly arrangement
C) A definite melting point
D) Both B and C
Answer: D) Both B and C

4. Which of the following is an example of a crystalline solid?

A) Rubber
B) Glass
C) Sodium chloride
D) Plastic
Answer: C) Sodium chloride

5. Crystalline solids exhibit:

A) Isotropy
B) Anisotropy
C) No melting point
D) Irregular cleavage
Answer: B) Anisotropy
6. Amorphous solids are also called:

A) True solids
B) Pseudo solids
C) Metallic solids
D) Network solids
Answer: B) Pseudo solids

7. Which of these is an example of an amorphous solid?

A) Diamond
B) Ice
C) Glass
D) Sodium chloride
Answer: C) Glass

8. Crystalline solids have which of the following properties?

A) Sharp melting point

B) Random atomic arrangement
C) No definite cleavage
D) Isotropic behavior
Answer: A) Sharp melting point

9. What type of solid is diamond?

A) Ionic
B) Molecular
C) Network covalent
D) Metallic
Answer: C) Network covalent

10. Ionic solids are composed of:

A) Covalently bonded molecules
B) Oppositely charged ions
C) Free electrons
D) Weak intermolecular forces
Answer: B) Oppositely charged ions

11. Which property is unique to metallic solids?

A) Low melting point

B) High conductivity
C) Brittle nature
D) Dissolving in water
Answer: B) High conductivity

12. The main bonding in molecular solids is:

A) Ionic bonding
B) Metallic bonding
C) Van der Waals forces
D) Covalent bonding
Answer: C) Van der Waals forces

13. The property of cleavage in crystalline solids means they:

A) Can be shaped easily

B) Shatter randomly
C) Break along definite planes
D) Dissolve in water
Answer: C) Break along definite planes

14. Network covalent solids generally:

A) Have high melting points

B) Are soft and malleable
C) Conduct electricity
D) Dissolve in water
Answer: A) Have high melting points
15. Which type of solid is highly malleable and ductile?

A) Ionic solid
B) Network solid
C) Metallic solid
D) Molecular solid
Answer: C) Metallic solid

16. Why do ionic solids have high melting points?

A) Weak intermolecular forces

B) Strong electrostatic forces between ions
C) Free-moving electrons
D) Low density
Answer: B) Strong electrostatic forces between ions

17. Amorphous solids melt:

A) At a sharp, distinct temperature

B) Over a range of temperatures
C) At extremely low temperatures
D) Only under pressure
Answer: B) Over a range of temperatures

18. Which of the following is NOT a type of crystalline solid?

A) Molecular solid
B) Ionic solid
C) Liquid crystal
D) Network covalent solid
Answer: C) Liquid crystal

19. Which solid has a "sea of electrons" model?

A) Network covalent
B) Ionic
C) Metallic
D) Molecular
Answer: C) Metallic

20. What happens when an ionic solid dissolves in water?

A) It forms neutral molecules

B) It becomes a nonconductor
C) It releases free ions
D) It remains unchanged
Answer: C) It releases free ions

21. Which of the following is a property of molecular solids?

A) High electrical conductivity

B) Hard and brittle
C) Low melting points
D) Strong electrostatic forces
Answer: C) Low melting points

22. The orderly arrangement in crystalline solids extends:

A) Over a short distance

B) Over long distances
C) Randomly
D) In single directions only
Answer: B) Over long distances

23. Which of the following is NOT a feature of metallic solids?

A) Malleability
B) Ductility
C) Conductivity
D) Brittle nature
Answer: D) Brittle nature
24. What determines the shape of a crystalline solid?

A) The external environment

B) The arrangement of atoms
C) Temperature changes
D) The amount of pressure applied
Answer: B) The arrangement of atoms

25. The seven crystal systems include all EXCEPT:

A) Cubic
B) Tetragonal
C) Hexagonal
D) Amorphous
Answer: D) Amorphous

Electronic configuration, quantum numbers

Electronic configuration refers to the arrangement of electrons in the atomic orbitals of an

element. Electrons occupy orbitals in a specific order, following certain rules to ensure stability
and proper distribution.

Rules Guiding Electronic Configuration:

1. Aufbau Principle – Electrons fill orbitals in order of increasing energy levels (starting
from the lowest energy orbital). The order follows the sequence:
1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p → 5s → 4d → 5p → 6s → 4f → 5d → 6p
→ 7s → 5f → 6d → 7p
2. Pauli Exclusion Principle – No two electrons in an atom can have the same set of all
four quantum numbers. This means an orbital can hold a maximum of two electrons with
opposite spins.
3. Hund’s Rule – Electrons occupy orbitals of the same energy (degenerate orbitals) singly
first before pairing occurs. This minimizes electron repulsion and maximizes stability.

2. Quantum Numbers and Their Types

Quantum numbers describe the unique position and energy of an electron in an atom. Each
electron has a set of four quantum numbers:

1. Principal Quantum Number (n):

 Indicates the main energy level or shell of an electron.
o
It takes positive integer values (1, 2, 3, etc.).
o
A higher n value means the electron is farther from the nucleus and has more
o
energy.
2. Azimuthal (Angular Momentum) Quantum Number (l):
o Determines the shape of an orbital.
o Depends on n and can take values from 0 to (n-1).
o Values correspond to different orbital types:
 l = 0 → s-orbital (spherical)
 l = 1 → p-orbital (dumbbell)
 l = 2 → d-orbital (cloverleaf)
 l = 3 → f-orbital (complex shape)
3. Magnetic Quantum Number (mₗ):
o Determines the orientation of an orbital in space.
o Can have integer values from -l to +l (including 0).
4. Spin Quantum Number (mₛ):
o Describes the spin of an electron within an orbital.
o Can have two possible values: +½ (spin-up) or -½ (spin-down).

Questions

1. Which principle states that electrons fill orbitals in order of increasing energy
levels?
A) Hund’s Rule
B) Pauli Exclusion Principle
C) Aufbau Principle
D) Heisenberg Uncertainty Principle
Answer: C) Aufbau Principle
2. Which quantum number determines the shape of an orbital?
A) Principal quantum number (n)
B) Azimuthal quantum number (l)
C) Magnetic quantum number (mₗ)
D) Spin quantum number (mₛ)
Answer: B) Azimuthal quantum number (l)
3. How many electrons can the p-orbital hold in total?
A) 2
B) 6
C) 10
D) 14
Answer: B) 6
4. Which of the following electronic configurations violates Hund’s Rule?
A) ↑ ↑ ↑
B) ↑↓ ↑ ↑
C) ↑ ↑ ↓
D) ↑ ↑ ↑↓
Answer: B) ↑↓ ↑ ↑
 5. The electron configuration of an atom follows the Pauli Exclusion Principle, which
means:
A) Electrons fill orbitals in order of increasing energy.
B) No two electrons in an atom can have the same set of all four quantum numbers.
C) Electrons must occupy degenerate orbitals singly before pairing.
D) The number of protons equals the number of electrons in a neutral atom.
Answer: B) No two electrons in an atom can have the same set of all four quantum
numbers.
6. Which of the following atoms has a completely filled s-orbital in its valence shell?
A) Oxygen
B) Fluorine
C) Neon
D) Sodium
Answer: C) Neon
7. What is the electronic configuration of a neutral carbon atom (Z = 6)?
A) 1s² 2s² 2p²
B) 1s² 2s² 2p⁶
C) 1s² 2s² 2p⁴
D) 1s² 2s²
Answer: A) 1s² 2s² 2p²
8. How many valence electrons does sulfur (S, Z = 16) have?
A) 4
B) 6
C) 8
D) 16
Answer: B) 6
9. What is the noble gas notation for the electronic configuration of calcium (Z = 20)?
A) [Ne] 3s² 3p⁶ 4s²
B) [Ar] 4s²
C) [He] 2s² 2p⁶ 3s² 3p⁶ 4s²
D) [Kr] 5s²
Answer: B) [Ar] 4s²
10. Which element has the following electronic configuration: 1s² 2s² 2p⁶ 3s² 3p¹?
A) Sodium (Na)
B) Magnesium (Mg)
C) Aluminum (Al)
D) Silicon (Si)
Answer: C) Aluminum (Al)

Determination of Quantum Numbers

11. The principal quantum number (n) describes:

A) The energy level of an electron
B) The shape of an orbital
 C) The spin of an electron
D) The orientation of an orbital
Answer: A) The energy level of an electron
12. For an electron in the 3p orbital, what is the value of the azimuthal quantum
number (l)?
A) 0
B) 1
C) 2
D) 3
Answer: B) 1
13. What is the possible range of values for the magnetic quantum number (mₗ) if l = 2?
A) -2 to +2
B) -1 to +1
C) -3 to +3
D) 0 to +2
Answer: A) -2 to +2
14. If an electron has quantum numbers n = 4 and l = 2, which orbital does it occupy?
A) 4s
B) 4p
C) 4d
D) 4f
Answer: C) 4d
15. How many possible orbitals exist for n = 3?
A) 3
B) 9
C) 4
D) 6
Answer: B) 9

Electronic Configurations of Common Atoms

16. What is the electronic configuration of oxygen (O, Z = 8)?

A) 1s² 2s² 2p⁴
B) 1s² 2s² 2p²
C) 1s² 2s² 2p⁶
D) 1s² 2s¹
Answer: A) 1s² 2s² 2p⁴
17. What is the electronic configuration of neon (Ne, Z = 10)?
A) 1s² 2s² 2p⁶
B) 1s² 2s² 2p²
C) 1s² 2s² 2p⁴
D) 1s² 2s²
Answer: A) 1s² 2s² 2p⁶
 18. Which of the following elements has the valence shell configuration 3s² 3p⁵?
A) Fluorine
B) Chlorine
C) Bromine
D) Neon
Answer: B) Chlorine
19. Which element has the electron configuration [Kr] 5s² 4d¹?
A) Yttrium (Y)
B) Zirconium (Zr)
C) Strontium (Sr)
D) Rubidium (Rb)
Answer: A) Yttrium (Y)
20. Which of the following is the correct electronic configuration of iron (Fe, Z = 26)?
A) [Ar] 4s² 3d⁶
B) [Ar] 3d⁶ 4s²
C) [Ar] 4s² 3d⁴
D) [Ne] 3s² 3p⁶ 4s² 3d⁶
Answer: A) [Ar] 4s² 3d⁶

Miscellaneous Questions

21. How many electrons can a d-orbital hold?

A) 2
B) 6
C) 10
D) 14
Answer: C) 10
22. What is the maximum number of electrons in the second energy level (n = 2)?
A) 2
B) 8
C) 18
D) 32
Answer: B) 8
23. Which rule states that orbitals of equal energy must be occupied singly before
pairing occurs?
A) Aufbau Principle
B) Hund’s Rule
C) Pauli Exclusion Principle
D) Heisenberg Uncertainty Principle
Answer: B) Hund’s Rule
24. Which element has the highest atomic number with a noble gas configuration?
A) Krypton
B) Radon
C) Xenon
 D) Argon
Answer: B) Radon
25. Which block of the periodic table do transition metals belong to?
A) s-block
B) p-block
C) d-block
D) f-block
Answer: C) d-block

Structure of solids

Introduction to Solids

 Solids have a fixed shape and volume due to closely packed particles.
 They exhibit high density, rigidity, and strong intermolecular forces.
 Solids are classified into Amorphous solids and Crystalline solids.

Types of Solids

1. Amorphous Solids
o Lack an orderly atomic arrangement.
o Examples: Glass, plastics, rubber.
o No sharp melting point, gradually soften over a temperature range.
o Isotropic properties (same properties in all directions).
2. Crystalline Solids
o Have a well-defined geometric structure.
o Examples: Salt, sugar, diamond.
o Exhibit anisotropy (properties vary with direction).
o Sharp melting point.
o Show cleavage along specific planes.

Types of Crystalline Solids

1. Ionic Solids
o Composed of oppositely charged ions.
o High melting points (300 - 1000°C).
o Hard, brittle, non-conductive in solid state but conductive in solution.
o Example: Sodium chloride (NaCl).
2. Molecular Solids
o Held by Van der Waals forces.
o Low melting points, soft.
o Can be polar (e.g., sugar) or non-polar (e.g., iodine, fullerenes).
o Non-conductive.
3. Network Covalent Solids
o Atoms connected by covalent bonds in a continuous network.
 o Very hard, high melting points, non-conductive.
o Examples: Diamond, quartz.
4. Metallic Solids
o Composed of metal atoms with a "sea of electrons".
o Malleable, ductile, and good conductors.
o Examples: Iron, copper, alloys (bronze, steel).

Key Properties of Crystalline Solids

 Orderly arrangement: Atoms repeat in a geometric pattern.

 Plane faces and fixed angles: Crystal faces meet at definite angles.
 Anisotropy: Properties depend on direction.
 Cleavage: Some crystals split along specific planes.
 Definite melting point: Sharp transition from solid to liquid.
 Heat of fusion: Definite amount of energy required to melt.

The seven system of crystals

Cubic (Isometric) System

 Axial Lengths: a = b = c
 Axial Angles: α = β = γ = 90°
 Example: Sodium chloride (NaCl), Diamond, Copper

2. Tetragonal System

 Axial Lengths: a = b ≠ c
 Axial Angles: α = β = γ = 90°
 Example: Tin (β-Sn), Zircon, Rutile (TiO₂)

3. Orthorhombic System

 Axial Lengths: a ≠ b ≠ c
 Axial Angles: α = β = γ = 90°
 Example: Sulfur, Rhombic Sulfur, Topaz

4. Monoclinic System

 Axial Lengths: a ≠ b ≠ c
 Axial Angles: α = γ = 90° ≠ β
 Example: Gypsum, Sugar, Mica

5. Triclinic System

 Axial Lengths: a ≠ b ≠ c
  Axial Angles: α ≠ β ≠ γ ≠ 90°
 Example: Kyanite, Turquoise, Feldspar

6. Hexagonal System

 Axial Lengths: a = b ≠ c
 Axial Angles: α = β = 90°, γ = 120°
 Example: Graphite, Quartz, Beryllium (Be)

7. Trigonal (Rhombohedral) System

 Axial Lengths: a = b = c
 Axial Angles: α = β = γ ≠ 90°
 Example: Calcite, Quartz (α-Quartz), Cinnabar

Questions

1. Which of the following is a characteristic of solids?

a) Indefinite shape and volume
b) Low density and weak intermolecular forces
c) Fixed shape and volume
d) Free movement of particles
o Answer: c) Fixed shape and volume
2. Why do solids have high density?
a) Particles are loosely packed
b) Strong intermolecular forces hold particles closely
c) Solids have weak bonds
d) Solids expand freely
o Answer: b) Strong intermolecular forces hold particles closely
3. Which of the following states of matter has the strongest intermolecular forces?
a) Gas
b) Liquid
c) Plasma
d) Solid
o Answer: d) Solid
4. How are solids classified?
a) Crystalline and amorphous
b) Ionic and metallic
c) Hard and soft
d) Polar and non-polar
o Answer: a) Crystalline and amorphous

2. Amorphous Solids
 5. Which of the following is NOT a characteristic of amorphous solids?
a) Lack of orderly atomic arrangement
b) Sharp melting point
c) Isotropic properties
d) Gradual softening over a temperature range
o Answer: b) Sharp melting point
6. Which of these is an example of an amorphous solid?
a) Diamond
b) Salt
c) Glass
d) Copper
o Answer: c) Glass
7. Why do amorphous solids lack a definite melting point?
a) They have weak bonds
b) They contain liquid molecules
c) Their bonds break over a range of temperatures
d) They have a fixed crystal structure
o Answer: c) Their bonds break over a range of temperatures
8. Which property of amorphous solids makes them different from crystalline solids?
a) Anisotropy
b) Isotropy
c) High melting point
d) Definite shape
o Answer: b) Isotropy

3. Crystalline Solids

9. Which of the following is NOT a property of crystalline solids?

a) Sharp melting point
b) Definite geometric structure
c) Isotropic behavior
d) Anisotropy
o Answer: c) Isotropic behavior
10. Which of the following is an example of a crystalline solid?
a) Rubber
b) Sugar
c) Plastic
d) Wax

 Answer: b) Sugar

11. What is the key difference between amorphous and crystalline solids?
a) Crystalline solids have a defined internal structure
b) Amorphous solids conduct electricity
 c) Crystalline solids lack a definite shape
d) Amorphous solids have high melting points

 Answer: a) Crystalline solids have a defined internal structure

12. Crystalline solids exhibit which of the following properties?

a) Cleavage along specific planes
b) No definite shape
c) Gradual melting
d) Lack of order

 Answer: a) Cleavage along specific planes

4. Types of Crystalline Solids

Ionic Solids

13. What type of bonding is present in ionic solids?

a) Covalent bonding
b) Ionic bonding
c) Metallic bonding
d) Hydrogen bonding

 Answer: b) Ionic bonding

14. Which of the following is an example of an ionic solid?

a) Diamond
b) Sodium chloride (NaCl)
c) Copper
d) Graphite

 Answer: b) Sodium chloride (NaCl)

15. Why are ionic solids brittle?

a) Weak intermolecular forces
b) Strong but directional ionic bonds
c) Free movement of electrons
d) High electrical conductivity in the solid state

 Answer: b) Strong but directional ionic bonds

Molecular Solids
 16. Which force holds molecular solids together?
a) Ionic bonds
b) Covalent bonds
c) Van der Waals forces
d) Metallic bonds

 Answer: c) Van der Waals forces

17. Which of the following is a molecular solid?

a) Ice
b) Iron
c) Diamond
d) NaCl

 Answer: a) Ice

18. Why do molecular solids have low melting points?

a) They have weak intermolecular forces
b) They are composed of ions
c) They conduct electricity
d) They have metallic bonding

 Answer: a) They have weak intermolecular forces

Network Covalent Solids

19. Which of the following is NOT a network covalent solid?

a) Graphite
b) Quartz
c) Sodium chloride
d) Diamond

 Answer: c) Sodium chloride

20. What type of bonding is present in network covalent solids?

a) Ionic bonding
b) Covalent bonding
c) Metallic bonding
d) Van der Waals forces

 Answer: b) Covalent bonding

21. Which property is NOT exhibited by network covalent solids?

a) High melting points
b) Hardness
 c) Electrical conductivity
d) Covalent bonding

 Answer: c) Electrical conductivity

Metallic Solids

22. What is the key feature of metallic solids?

a) Strong ionic bonds
b) Delocalized electrons
c) High brittleness
d) Weak intermolecular forces

 Answer: b) Delocalized electrons

23. Which of the following is NOT a characteristic of metallic solids?

a) Ductility
b) High electrical conductivity
c) High brittleness
d) Malleability

 Answer: c) High brittleness

24. Why do metallic solids conduct electricity?

a) They contain ions
b) They have freely moving electrons
c) They are composed of neutral molecules
d) They have covalent bonds

 Answer: b) They have freely moving electrons

25. Which of the following is an example of a metallic solid?

a) Ice
b) Diamond
c) Copper
d) NaCl

 Answer: c) Copper

Kinetic theory of gases

The Kinetic Theory of Gases explains the macroscopic properties of gases (such as pressure,
temperature, and volume) in terms of the microscopic motion of gas molecules. The theory is
based on the following assumptions:
Gas molecules are in constant, random motion.

The size of gas molecules is negligible compared to the space they occupy.

Collisions between gas molecules and with the container walls are perfectly elastic, meaning no
energy is lost.

There are no intermolecular forces between gas molecules except during collisions.

The average kinetic energy of gas molecules is proportional to the absolute temperature (Kelvin

Laws That Make Up the Ideal Gas Equation

The ideal gas equation is:

PV=nRTPV=nRT

where:

 PP = Pressure (atm, Pa)

 VV = Volume (L, m³)
 nn = Number of moles (mol)
 RR = Universal gas constant (8.314 J/mol·K or 0.0821 L·atm/mol·K)
 TT = Temperature (Kelvin)

This equation is derived from four fundamental gas laws:

1. Boyle’s Law (Pressure-Volume Relationship)

P∝1V(at constant n and T)

When temperature and the number of moles are constant, pressure and volume are
o
inversely proportional.
o Example: When a gas is compressed, its pressure increases.
2. Charles’s Law (Temperature-Volume Relationship)

V∝T(at constant P and n)

At constant pressure, volume increases with temperature.

o
Example: A balloon expands when heated.
o
3. Gay-Lussac’s Law (Temperature-Pressure Relationship)

P∝T(at constant V and n)

o At constant volume, pressure increases with temperature.

o
 o Example: A closed aerosol can explode when exposed to heat.
4. Avogadro’s Law (Volume-Mole Relationship)

V∝n(at constant P and T)V∝n(at constant P and T)

o At constant pressure and temperature, volume increases with the number of gas
molecules.
o Example: Adding gas to a balloon increases its volume.
Questions

1. Which of the following is NOT an assumption of the kinetic theory of gases?

a) Gas molecules are in continuous random motion
b) Gas molecules have strong intermolecular forces
c) The volume of individual gas molecules is negligible
d) Collisions between gas molecules are elastic
o Answer: b) Gas molecules have strong intermolecular forces
2. According to the kinetic theory of gases, gas molecules move in:
a) Straight-line motion until they collide
b) Circular motion around a center
c) Random orbits
d) A fixed pattern
o Answer: a) Straight-line motion until they collide
3. Which of the following correctly describes an ideal gas according to the kinetic
theory?
a) Gas molecules occupy a fixed volume
b) Gas molecules experience no attractive or repulsive forces
c) Gas molecules lose energy after collisions
d) The motion of gas molecules is limited
o Answer: b) Gas molecules experience no attractive or repulsive forces

2. Gas Properties and Molecular Speed

4. The kinetic energy of a gas molecule is directly proportional to:

a) Its pressure
b) Its volume
c) Its absolute temperature
d) Its molecular mass
o Answer: c) Its absolute temperature
5. Which of the following increases as the temperature of a gas increases?
a) The mass of each gas molecule
b) The number of gas molecules
c) The average kinetic energy of the gas molecules
d) The size of gas molecules
o Answer: c) The average kinetic energy of the gas molecules
6. Which of the following correctly expresses the relationship between the average
kinetic energy of gas molecules and temperature?
a) KE=3/2KT
b) KE=1/2KT
c) KE=1/2mv^2
d) KE=PV
 o Answer: a) KE=3/2KT

3. Gas Laws and Their Interpretation

7. Which gas law states that the volume of a gas is directly proportional to its
temperature at constant pressure?
a) Boyle’s Law
b) Charles’ Law
c) Dalton’s Law
d) Avogadro’s Law
o Answer: b) Charles’ Law
8. Boyle’s Law states that at constant temperature:
a) Pressure is directly proportional to volume
b) Volume is directly proportional to temperature
c) Pressure is inversely proportional to volume
d) Pressure is independent of volume
o Answer: c) Pressure is inversely proportional to volume
9. The ideal gas equation is given as:
a) PV=kT
b) P+V=nRT
c) PV=nRT
d) PV=RT
o Answer: c) PV=nRT

4Which speed is the most probable speed of gas molecules?

a) Root mean square speed
b) Average speed
c) Most probable speed
d) Escape velocity

 Answer: c) Most probable speed

5. Diffusion and Effusion

13. According to Graham’s Law of Diffusion, the rate of diffusion of a gas is inversely
proportional to:
a) Its temperature
b) Its molecular mass
c) Its volume
d) Its kinetic energy
  Answer: b) Its molecular mass

14. If two gases, A and B, diffuse under identical conditions, and gas A has twice the
molar mass of gas B, what is the ratio of their diffusion rates?
a) 1:1
b) 1:2
c) 1:21:2
d) 2:12:1

 Answer: c) 1:21:2

15. Which of the following processes describes the movement of gas molecules through a
small hole without collisions?
a) Diffusion
b) Effusion
c) Conduction
d) Evaporation

 Answer: b) Effusion

Calculations

1. Calculate the pressure of 2 moles of an ideal gas occupying 10 L at a temperature of

300 K.
2. A gas sample has a volume of 15 L at a pressure of 2 atm and temperature of 350 K.
How many moles of gas are present? (
3. What is the volume occupied by 3 moles of an ideal gas at 500 K and 5 atm
pressure?
4. At what temperature will 4 moles of a gas occupy 20 L under 3 atm pressure?
5. Determine the mass of oxygen gas (O22) in a 22.4 L container at STP. (Molar mass
of O22 = 32 g/mol,
6. A gas cylinder contains 5 moles of nitrogen gas at a pressure of 8 atm and
temperature of 400 K. Find the volume of the gas.
7. If 3.2 g of methane (CH44) is kept in a 2 L container at 273 K, what is the pressure
inside the container? (Molar mass of CH44 = 16 g/mol,
8. A 5 L gas container contains 0.2 moles of helium gas at 1 atm pressure. Find the
temperature in Kelvin.

General Gas Law (P1V1T1=P2V2T2T1P1V1=T2P2V2) Based Questions

9. A gas has a volume of 6 L at 300 K and 2 atm pressure. What will be its volume at
400 K and 4 atm pressure?
10. A gas occupies 500 mL at 1.2 atm and 298 K. If the temperature increases to 350 K
and the pressure becomes 1.5 atm, find the new volume.
11. A 2 L sample of gas at 1 atm is heated from 273 K to 546 K. What is the final
pressure if the volume remains constant?
12. A gas sample at 2 atm and 400 K expands from 3 L to 6 L. What is the final
pressure if the temperature increases to 600 K?
13. A 1 L sample of gas at 400 K and 5 atm expands to 3 L at constant pressure. What is
the final temperature?
14. A 250 mL gas sample is cooled from 350 K to 175 K while keeping the pressure
constant. What is the new volume?
15. A gas at 1 atm, 500 K, and 2 L is compressed to 1 L and cooled to 250 K. What is the
new pressure?