Q1. Identify the FALSE statement about the law of conservation of energy.

Background

Topic: Thermodynamics – Law of Conservation of Energy

This question tests your understanding of the law of conservation of energy, which is a fundamental principle in chemistry and physics.

Key Terms:

• System and Surroundings: The part of the universe being studied (system) and everything else (surroundings).

• Conservation of Energy: Energy cannot be created or destroyed, only transferred or transformed.

• Kinetic and Potential Energy: Forms of energy that can be interconverted.

Step-by-Step Guidance

1. Read each statement carefully and recall the law of conservation of energy.

2. Identify which statements correctly describe energy transfer, transformation, or conservation.

3. Look for the statement that contradicts the law (e.g., claims energy cannot be transferred or can be destroyed/created).

4. Narrow down your choices by eliminating statements that are consistent with the law.

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Q2. Which signs on q and w represent the surroundings doing work on the system and losing heat to the surroundings?

Background

Topic: Thermodynamics – First Law Sign Conventions

This question tests your understanding of the sign conventions for heat (q) and work (w) in thermodynamics.

Key Terms:

• q (heat): Positive if heat is absorbed by the system, negative if released.

• w (work): Positive if work is done on the system, negative if done by the system.

Step-by-Step Guidance

1. Recall the sign conventions: q is positive when the system gains heat, negative when it loses heat; w is positive when work is done on the system, negative when the system does work on the surroundings.

2. "Surroundings doing work on the system" means w is positive.

3. "Losing heat to the surroundings" means q is negative (system is losing heat).

4. Match these sign conventions to the answer choices.

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Q3. Calculate the change in internal energy (ΔE) for a system that gives off 25.0 kJ of heat and changes from 12.00 L to 6.00 L in volume at 1.50 atm pressure. (101.3 J = 1 L·atm)

Background

Topic: First Law of Thermodynamics – Internal Energy Change

This question tests your ability to calculate the change in internal energy using heat and work, including unit conversions.

Key Formula:

Where:

• = heat (negative if given off by the system)

• = work (negative if work is done by the system)

• = pressure (in atm)

• = change in volume (final - initial, in L)

Step-by-Step Guidance

1. Assign the correct sign to q: since the system gives off heat, q is negative ( kJ).

2. Calculate : .

3. Calculate work: (make sure to use the correct sign for ).

4. Convert work from L·atm to kJ using the conversion factor (101.3 J = 1 L·atm; 1000 J = 1 kJ).

5. Plug q and w (both in kJ) into to set up the final calculation.

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Q4. A 2.49 g sample of aniline (C6H5NH2, molar mass = 93.13 g/mol) is combusted in a bomb calorimeter with a heat capacity of 4.25 kJ/°C. If the temperature rises from 29.5°C to 69.8°C, determine the value of ΔH°comb for aniline.

Background

Topic: Calorimetry – Bomb Calorimeter and Enthalpy of Combustion

This question tests your ability to use calorimetry data to calculate the enthalpy of combustion per mole of a substance.

Key Formula:

Where:

• = heat capacity of the calorimeter (kJ/°C)

• = change in temperature (°C)

• = heat absorbed by the calorimeter (kJ)

• Moles of aniline = mass / molar mass

Step-by-Step Guidance

1. Calculate .

2. Calculate using .

3. Determine the number of moles of aniline combusted: .

4. Calculate per mole: .

5. Remember to assign the correct sign to (combustion is exothermic).

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Q5. For a reaction with positive ΔH and positive ΔS, which statement is TRUE?

Background

Topic: Thermodynamics – Spontaneity and Gibbs Free Energy

This question tests your understanding of how enthalpy and entropy changes affect the spontaneity of a reaction.

Key Formula:

Where:

• = Gibbs free energy change

• = enthalpy change (positive in this case)

• = entropy change (positive in this case)

• = temperature in Kelvin

Step-by-Step Guidance

1. Recall that a reaction is spontaneous when .

2. With both and positive, consider how increasing temperature affects .

3. Analyze the formula: as increases, the term becomes more negative.

4. Determine at what temperature range (high or low) will be negative (spontaneous).

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Q6. Above what temperature does the following reaction become nonspontaneous? 2 H₂S(g) + 3 O₂(g) → 2 SO₂(g) + 2 H₂O(g); ΔH = -1036 kJ; ΔS = -153.2 J/K

Background

Topic: Thermodynamics – Gibbs Free Energy and Spontaneity

This question tests your ability to determine the temperature at which a reaction changes from spontaneous to nonspontaneous using ΔH and ΔS.

Key Formula:

Set to find the temperature where spontaneity changes.

Where:

• = enthalpy change (in kJ, convert to J if needed)

• = entropy change (in J/K)

• = temperature in K

Step-by-Step Guidance

1. Set and solve for : .

2. Rearrange to .

3. Convert to J if necessary (1 kJ = 1000 J).

4. Plug in the values for and (with correct signs).

5. Calculate (make sure to use absolute values as needed for temperature).

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Q7. Use Hess's law to calculate ΔG°rxn for: NO(g) + O(g) → NO₂(g), given the following reactions and their ΔG°rxn values.

• 2 O₃(g) → 3 O₂(g); ΔG°rxn = +489.6 kJ

• O₂(g) → 2 O(g); ΔG°rxn = +463.4 kJ

• NO(g) + O₃(g) → NO₂(g) + O₂(g); ΔG°rxn = -199.5 kJ

Background

Topic: Thermodynamics – Hess's Law for Gibbs Free Energy

This question tests your ability to use Hess's law to combine reactions and calculate the ΔG°rxn for a target reaction.

Key Concepts:

• Hess's Law: The overall ΔG°rxn is the sum of the ΔG°rxn values for the steps that add up to the target reaction.

• Reverse or multiply reactions as needed, adjusting ΔG°rxn accordingly.

Step-by-Step Guidance

1. Write out the target reaction and the given reactions.

2. Determine how to combine the given reactions (possibly reversing or scaling) to obtain the target reaction.

3. Adjust the ΔG°rxn values for any reversed or scaled reactions (reverse sign if reversed, multiply if scaled).

4. Add the adjusted ΔG°rxn values to set up the calculation for the target reaction.

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Q8. Calculate ΔGrxn at 298 K for 2 Hg(g) + O₂(g) → 2 HgO(s), given ΔG° = -180.8 kJ and P(Hg) = 0.025 atm, P(O₂) = 0.037 atm.

Background

Topic: Thermodynamics – Nonstandard Gibbs Free Energy

This question tests your ability to calculate ΔG under nonstandard conditions using partial pressures.

Key Formula:

Where:

• = standard Gibbs free energy change (kJ)

• = gas constant ( J/mol·K)

• = temperature in K

• = reaction quotient (use partial pressures for gases)

Step-by-Step Guidance

1. Write the expression for for the reaction: (since HgO is a solid, it is omitted).

2. Plug in the given partial pressures for Hg and O₂.

3. Calculate numerically.

4. Plug , , , and into the formula (convert $\Delta G^\circ$ to J if needed).

5. Set up the calculation for .

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Q9. Write a balanced reaction for which the following rate relationships are true: Rate = = = -

Background

Topic: Chemical Kinetics – Rate Expressions and Stoichiometry

This question tests your ability to relate the rate of change of concentration of reactants and products to the stoichiometry of a balanced chemical equation.

Key Concepts:

• The rate of a reaction is related to the change in concentration of each species, divided by its stoichiometric coefficient.

• For a general reaction ,

Step-by-Step Guidance

1. Examine the given rate relationships and identify the stoichiometric coefficients implied by the rate expressions.

2. Match the coefficients to the balanced chemical equations provided in the answer choices.

3. Look for the equation where the rate of disappearance and appearance of each species matches the given relationships.

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Q10. Determine the rate law and the value of k for the reaction NO₂(g) + O₃(g) → NO₃(g) + O₂(g) using the data provided.

Background

Topic: Chemical Kinetics – Rate Laws and Determining Reaction Order

This question tests your ability to determine the rate law and rate constant from experimental data.

Key Formula:

General rate law:

Where:

• = rate constant

• = order with respect to NO₂

• = order with respect to O₃

Step-by-Step Guidance

1. Compare experiments where only one reactant concentration changes to determine the order with respect to that reactant.

2. Use the ratio of rates and concentrations to solve for the exponents and .

3. Once the orders are known, use any set of data to solve for by plugging values into the rate law.

4. Check units for based on the overall order of the reaction.

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