Which of the following is the correct definition of formation enthalpy?

A. The enthalpy change when one mole of a compound is formed from its elements in their standard states.

B. The enthalpy change when one mole of a compound is burned completely in oxygen.

C. The enthalpy change when a solution is formed from a solute and a solvent.

D. The enthalpy change when an acid reacts with a base to form one mole of water.

What does combustion enthalpy refer to?

A. The energy required to break one mole of bonds in a compound.

B. The energy change when one mole of a compound is dissolved in water.

C. The energy change when one mole of a compound is burned completely in oxygen.

D. The energy change when an acid reacts with a base.

Which of the following best describes Hess's Law?

A. The total enthalpy change of a reaction is independent of the route taken.

B. The enthalpy change of a reaction is directly proportional to the amount of reactants used.

C. The enthalpy change of a reaction is always negative.

D. The enthalpy change of a reaction is equal to the sum of the enthalpies of the products minus the enthalpies of the reactants.

What is the significance of standard conditions when discussing enthalpy changes?

A. They refer to conditions at 0°C and 1 atm pressure.

B. They ensure that enthalpy changes are measured under consistent conditions for comparison.

C. They are the conditions under which all reactions occur.

D. They are the conditions at which a reaction is most spontaneous.

Which of the following reactions is likely to have a negative solution enthalpy?

A. A reaction in which a gas is produced.

B. A reaction in which a solid is formed from a solution.

C. A reaction in which a solution gets colder upon dissolving a solute.

D. A reaction in which a solution gets warmer upon dissolving a solute.

a) Define formation enthalpy and explain its significance in thermochemistry. [3]

b) Given that the formation enthalpy of water is -285.8 kJ/mol, determine the energy change when 2 moles of water are formed from its elements. [2]

a) Describe the concept of combustion enthalpy and how it differs from formation enthalpy. [3]

b) Why is the combustion enthalpy of methane always negative? [2]

a) Explain the principle behind Hess's Law and its applications in determining enthalpy changes. [3]

b) If the enthalpy change for the combustion of carbon to carbon dioxide is -393.5 kJ/mol and that of the combustion of hydrogen to water is -285.8 kJ/mol, calculate the enthalpy change for the combustion of methane (CH4). [2]

a) Discuss the concept of entropy and its relation to the disorder of a system. [3]

b) Calculate the change in entropy when 2 moles of a gas at 25°C and 1 atm pressure expand irreversibly from a volume of 5 L to 10 L. (Assume ideal behaviour) [4]

c) Explain how the change in entropy is related to the spontaneity of a process. [2]

a) Define standard conditions in thermochemistry and explain why they are important when calculating enthalpy changes. [3]

b) Calculate the enthalpy change for the combustion of 1 mole of benzene (C6H6) to form carbon dioxide and water vapour. Given: C6H6(l) + 15/2 O2(g) → 6 CO2(g) + 3 H2O(g) ΔH = -3267 kJ/mol [4]

c) Discuss the factors that can affect the magnitude of the entropy change in a chemical reaction. [3]

In an energy diagram, what does the height of the peak represent?

A. The change in enthalpy for the reaction.

B. The activation energy for the forward reaction.

C. The energy of the products.

D. The energy of the reactants.

Which of the following is true about neutralisation enthalpy?

A. It is always positive.

B. It refers to the energy change when an acid reacts with a metal.

C. It is the energy change when an acid reacts with a base to form one mole of water.

D. It is the energy change when a base reacts with water.

What is the importance of enthalpy changes in chemical reactions?

A. They determine the speed of a reaction.

B. They indicate whether a reaction is spontaneous or not.

C. They provide information about the bond strengths in reactants and products.

D. They determine the equilibrium position of a reaction.

For a given reaction, if the formation enthalpy of the products is greater than the formation enthalpy of the reactants, the reaction is:

A. Exothermic

B. Endothermic

C. At equilibrium

D. Impossible

Which of the following best describes the combustion enthalpy of methane?

A. The energy change when one mole of methane is formed from carbon and hydrogen.

B. The energy change when one mole of methane reacts with oxygen to form carbon dioxide and water.

C. The energy change when one mole of methane is dissolved in water.

D. The energy change when methane reacts with an acid.

a) Explain the concept of microstates in relation to entropy and the second law of thermodynamics. [3]

b) Calculate the change in entropy when 1 mole of ice at -10°C is heated to form 1 mole of liquid water at 20°C. (Assume ideal behavior) [4]

c) Discuss the implications of a negative ΔG value for a chemical reaction in terms of spontaneity. [3]

a) Define Gibbs free energy (G) and explain its significance in determining spontaneity. [3]

b) Calculate the change in Gibbs free energy (ΔG) for a reaction where ΔH = -150 kJ and ΔS = 100 J/K at 25°C. [4]

c) How does the temperature affect the spontaneity of a reaction, as indicated by the sign of ΔG? [3]

a) Explain the concept of standard enthalpy of formation (ΔHf°) and how it is determined. [4]

b) Calculate the standard enthalpy change (ΔH°) for the combustion of 2 moles of methane (CH4) gas at 25°C and 1 atm pressure to form carbon dioxide (CO2) and water (H2O) vapour. Given: ΔHf° (CH4) = -74.8 kJ/mol ΔHf° (CO2) = -393.5 kJ/mol ΔHf° (H2O) = -285.8 kJ/mol [5]

c) Discuss the significance of Hess's Law in determining enthalpy changes and its practical applications in thermochemistry. [3]

d) Calculate the standard entropy change (ΔS°) for the combustion of 1 mole of ethene (C2H4) gas at 25°C and 1 atm pressure to form carbon dioxide (CO2) and water (H2O) vapour. Given: ΔS° (C2H4) = 219.6 J/(mol·K) ΔS° (CO2) = 213.8 J/(mol·K) ΔS° (H2O) = 188.8 J/(mol·K) [4]

a) Define the term "spontaneous process" in thermodynamics and explain how it relates to the sign of ΔG. [4]

b) Calculate the standard Gibbs free energy change (ΔG°) for the reaction: 2 SO2(g) + O2(g) → 2 SO3(g) Given the following data: ΔH° = -198.2 kJ/mol ΔS° = -160.2 J/(mol·K) T = 298 K [5]

c) Describe the concept of dynamic equilibrium and how it is related to reversible processes in thermodynamics. [3]

d) Calculate the standard Gibbs free energy change (ΔG°) for the dissolution of 1 mole of calcium chloride (CaCl2) in water at 25°C. Given: ΔH° = -82.8 kJ/mol ΔS° = -161.9 J/(mol·K) [4]

a) Explain the concept of entropy (S) in thermodynamics and how it is related to the degree of disorder in a system. [4]

b) Calculate the standard Gibbs free energy change (ΔG°) for the vaporisation of 1 mole of water (H2O) at its boiling point (100°C). Given: ΔHvap° = 40.7 kJ/mol ΔSvap° = 109.0 J/(mol·K) T = 373 K [5]

c) Discuss the effect of concentration on the spontaneity of a reaction, with reference to Le Chatelier's principle. [3]

d) Calculate the standard Gibbs free energy change (ΔG°) for the combustion of 1 mole of glucose (C6H12O6) to form carbon dioxide (CO2) and water (H2O) vapor at 25°C. Given: ΔH° = -2800 kJ/mol ΔS° = -909 J/(mol·K) [4]