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

Background

Topic: Thermodynamics – Conservation of Energy

This question tests your understanding of the law of conservation of energy, which is a fundamental concept in thermodynamics. It asks you to identify which statement does NOT correctly describe this law.

Key Terms:

• Law of Conservation of Energy: States that energy cannot be created or destroyed, only transformed or transferred.

• System and Surroundings: The system is the part of the universe you are studying; the surroundings are everything else.

• Potential and Kinetic Energy: Energy can exist in different forms, such as stored (potential) or motion (kinetic).

Step-by-Step Guidance

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

2. For each statement, ask yourself: Does this statement agree with the idea that energy can only be transferred or transformed, not created or destroyed?

3. Identify which statement contradicts the law (for example, by saying energy cannot be transferred or can be created/destroyed).

4. Eliminate the statements that are consistent with the law, and focus on the one that is inconsistent.

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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, specifically for the system and surroundings.

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.

• System vs. Surroundings: The direction of energy flow determines the sign.

Step-by-Step Guidance

1. Recall that if the surroundings do work on the system, w is positive for the system.

2. If the system loses heat to the surroundings, q is negative for the system.

3. Match these sign conventions to the answer choices provided.

4. Eliminate options that do not fit both criteria.

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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 apply the first law of thermodynamics to calculate the change in internal energy, considering both heat and work.

Key Formula:

Where:

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

• = work done on or by the system

Work done by the system at constant pressure:

Unit conversion:

Step-by-Step Guidance

1. Identify the values: (since heat is given off), , , .

2. Calculate .

3. Calculate (be careful with the sign and units).

4. Convert from L·atm to kJ using the conversion factor.

5. Set up with your calculated values, but do not compute the final sum yet.

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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 – Enthalpy of Combustion

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

Key Formulas:

• Heat absorbed by calorimeter:

• For a bomb calorimeter,

• Convert grams of aniline to moles:

• Calculate per mole:

Step-by-Step Guidance

1. Calculate the temperature change: .

2. Calculate the heat absorbed by the calorimeter: .

3. Determine (the heat released by the reaction): .

4. Convert the mass of aniline to moles using its molar mass.

5. Set up the calculation for per mole, but do not compute the final value yet.

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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 (ΔH) and entropy (ΔS) affect the spontaneity of a reaction, as described by Gibbs free energy.

Key Formula:

• If , the reaction is spontaneous.

• If , the reaction is nonspontaneous.

Step-by-Step Guidance

1. Recall that both and are positive.

2. Consider how increasing temperature () affects in the equation .

3. Think about what happens to at low and high temperatures.

4. Use this reasoning to determine at which temperatures the reaction is spontaneous or nonspontaneous.

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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:

• The reaction becomes nonspontaneous when .

• Solve for when .

Step-by-Step Guidance

1. Set and rearrange the equation to solve for .

2. Plug in the given values for and (be careful with units: convert kJ to J or vice versa).

3. Set up the calculation for , but do not compute the final value yet.

4. Check the sign of and to confirm the direction of spontaneity at different temperatures.

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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 standard Gibbs free energy change for a target reaction.

Key Concepts:

• Hess's Law: The total ΔG° for a reaction can be found by adding the ΔG° values of steps that sum to the overall reaction.

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

Step-by-Step Guidance

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

2. Determine how to combine or reverse the given reactions to obtain the target reaction.

3. Adjust the ΔG° values for any reactions you reverse (change the sign) or multiply (multiply ΔG° by the same factor).

4. Set up the sum of ΔG° values for the sequence, but do not add them up yet.

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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 the Gibbs free energy change under nonstandard conditions using partial pressures.

Key Formula:

• is the reaction quotient, calculated from partial pressures.

Step-by-Step Guidance

1. Write the expression for for the reaction using the given partial pressures.

2. Plug the values for , , , and into the formula .

3. Be careful with units: convert to J if necessary.

4. Set up the calculation for , but do not compute the final value yet.

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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 interpret rate expressions and relate them to the stoichiometry of a balanced chemical equation.

Key Concepts:

• Rate expressions relate the change in concentration of reactants and products to the overall reaction rate.

• The coefficients in the balanced equation determine the relative rates.

Step-by-Step Guidance

1. Recall that the rate of disappearance of reactants and appearance of products is related by their stoichiometric coefficients.

2. Match the given rate relationships to the possible balanced equations provided.

3. Look for the equation where the coefficients match the rate relationships (e.g., if the rate of disappearance of N₂O is twice the rate of appearance of O₂, the coefficient for N₂O should be 2, and for O₂ should be 1).

4. Eliminate options that do not fit the rate 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 Rate Constants

This question tests your ability to determine the rate law (orders of reaction with respect to each reactant) and calculate the rate constant using experimental data.

Key Formulas:

• General rate law:

• Use ratios of experiments to determine the order with respect to each reactant.

• Once orders are known, use any data set to solve for .

Step-by-Step Guidance

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

2. Repeat for the other reactant.

3. Write the full rate law with the determined orders.

4. Plug in the concentrations and rate from any experiment to solve for , but do not compute the final value yet.

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