General Chemistry Study Guide: Thermochemistry, Quantum Chemistry, and Periodic Trends

Q1. What happens to the internal energy of the system and surroundings in various scenarios?

Background

Topic: Thermodynamics – Internal Energy

This question tests your understanding of energy transfer between a system and its surroundings, specifically how internal energy changes during processes.

Key Terms:

• System: The part of the universe being studied.

• Surroundings: Everything outside the system.

• Internal Energy ($U$): The total energy contained within the system.

Step-by-Step Guidance

1. Recall that energy can be transferred as heat or work between the system and surroundings.

2. Consider the law of conservation of energy: energy lost by the system is gained by the surroundings, and vice versa.

3. Think about whether the process described is endothermic (system absorbs energy) or exothermic (system releases energy).

4. Analyze each answer choice to determine if it correctly describes the energy changes for both system and surroundings.

Try solving on your own before revealing the answer!

Q2. A system releases 622 kJ of heat and does 105 kJ of work on the surroundings. What is the change in internal energy of the system?

Background

Topic: First Law of Thermodynamics

This question tests your ability to apply the first law of thermodynamics to calculate the change in internal energy ($\Delta U$) of a system.

Key formula:

$\Delta U = q + w$

• $q$ = heat absorbed by the system (negative if released)

• $w$ = work done on the system (negative if done by the system)

Step-by-Step Guidance

1. Identify the values: $q = -622$ kJ (heat released), $w = -105$ kJ (work done by the system).

2. Plug these values into the formula: $\Delta U = q + w$.

3. Set up the calculation: $\Delta U = (-622) + (-105)$ kJ.

4. Combine the values to find the change in internal energy, but stop before calculating the final sum.

Try solving on your own before revealing the answer!

Q3. The gas in a piston warms and absorbs 655 J of heat. The expansion performs 344 J of work on the surroundings. What is the change in internal energy for the system?

Background

Topic: First Law of Thermodynamics

This question tests your understanding of how heat and work affect the internal energy of a system.

Key formula:

$\Delta U = q + w$

• $q$ = heat absorbed by the system (positive if absorbed)

• $w$ = work done on the system (negative if done by the system)

Step-by-Step Guidance

1. Identify the values: $q = +655$ J (heat absorbed), $w = -344$ J (work done by the system).

2. Plug these values into the formula: $\Delta U = q + w$.

3. Set up the calculation: $\Delta U = 655 + (-344)$ J.

4. Combine the values to find the change in internal energy, but stop before calculating the final sum.

Try solving on your own before revealing the answer!

Q4. How much heat is required to warm 1.50 L of water from 25.0°C to 100.0°C? (Assume a density of 1.0 g/mL for the water.)

Background

Topic: Calorimetry – Heat Capacity

This question tests your ability to calculate the heat required to change the temperature of a substance using its mass, specific heat, and temperature change.

Key formula:

$q = m c \Delta T$

• $q$ = heat absorbed (in J)

• $m$ = mass (in g)

• $c$ = specific heat capacity (for water, $4.18$ J/g·°C)

• $\Delta T$ = change in temperature ($T_{final} - T_{initial}$)

Step-by-Step Guidance

1. Convert the volume of water to mass using the density: $1.50$ L $\times$ $1000$mL/L$\times$$1.0$ g/mL.

2. Calculate the temperature change: $\Delta T = 100.0 - 25.0$ °C.

3. Plug the mass, specific heat, and $\Delta T$ into the formula $q = m c \Delta T$.

4. Set up the calculation, but stop before multiplying out the final value.

Try solving on your own before revealing the answer!

Q5. Suppose that 25 g of each substance is initially at 27.0°C. What is the final temperature of each substance upon absorbing 2.35 kJ of heat?

Background

Topic: Calorimetry – Specific Heat

This question tests your ability to calculate the final temperature of a substance after absorbing a certain amount of heat.

Key formula:

$q = m c \Delta T$

• $q$ = heat absorbed (in J)

• $m$ = mass (in g)

• $c$ = specific heat capacity (varies for gold, silver, aluminum, water)

• $\Delta T$ = change in temperature ($T_{final} - T_{initial}$)

Step-by-Step Guidance

1. Convert 2.35 kJ to J: $2.35 \times 10^3$ J.

2. Rearrange the formula to solve for $T_{final}$: $T_{final} = T_{initial} + \frac{q}{m c}$.

3. Plug in the values for each substance (use the specific heat for gold, silver, aluminum, and water).

4. Set up the calculation for each, but stop before computing the final temperatures.

Try solving on your own before revealing the answer!

Q6. When 1 mol of a fuel burns at constant pressure, it produces 3452 kJ of heat and does 11 kJ of work. What are ΔE and ΔH for the combustion of the fuel?

Background

Topic: Thermochemistry – Internal Energy and Enthalpy

This question tests your understanding of the relationship between internal energy ($\Delta E$) and enthalpy ($\Delta H$) for a reaction.

Key formulas:

$\Delta E = q + w$

$\Delta H = q_{p}$ (heat at constant pressure)

• $q$ = heat absorbed/released

• $w$ = work done

Step-by-Step Guidance

1. Identify the values: $q_p = -3452$ kJ (heat released), $w = -11$ kJ (work done by the system).

2. Calculate $\Delta E$ using $\Delta E = q + w$.

3. Recall that $\Delta H = q_p$ at constant pressure.

4. Set up the calculations for both $\Delta E$ and $\Delta H$, but stop before finding the final values.

Try solving on your own before revealing the answer!

Q7. Determine whether each process is exothermic or endothermic and indicate the sign of ΔH.

Background

Topic: Thermochemistry – Exothermic and Endothermic Processes

This question tests your ability to classify processes as exothermic or endothermic and assign the correct sign to enthalpy change ($\Delta H$).

Key Terms:

• Exothermic: Releases heat ($\Delta H$ is negative)

• Endothermic: Absorbs heat ($\Delta H$ is positive)

Step-by-Step Guidance

1. For each process, decide if heat is absorbed or released.

2. Assign the correct sign to $\Delta H$ based on whether the process is exothermic or endothermic.

3. Write out your reasoning for each process, but stop before stating the final classification and sign.

Try solving on your own before revealing the answer!

Q8. Consider the thermochemical equation for the combustion of acetone. If a bottle of nail polish remover contains 177 mL of acetone, how much heat is released by its complete combustion? The density of acetone is 0.788 g/mL.

Background

Topic: Thermochemistry – Stoichiometry and Heat Calculations

This question tests your ability to use stoichiometry and thermochemical equations to calculate heat released in a reaction.

Key formula:

$q = n \Delta H_{rxn}$

• $n$ = moles of acetone

• $\Delta H_{rxn}$ = enthalpy change per mole

Step-by-Step Guidance

1. Convert the volume of acetone to mass: $177$mL$\times$$0.788$ g/mL.

2. Convert mass to moles using the molar mass of acetone ($C_3H_6O$).

3. Multiply moles by $\Delta H_{rxn}$ to find the total heat released.

4. Set up the calculation, but stop before multiplying out the final value.

Try solving on your own before revealing the answer!

Final Answer: See below for each question's solution.

Each question's final answer involves plugging in the values and completing the calculations as set up above. Review your steps and check your work for accuracy!