Thermochemistry & Thermodynamics Review

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Question-and-answer flashcards covering key concepts from thermochemistry, bond energy, heat/work, enthalpy, and the three laws of thermodynamics.

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26 Terms

1
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What does thermochemistry study?

The energy changes—especially heat—during chemical reactions and physical transformations.

2
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How is energy defined in thermochemistry?

The capacity to do work or transfer heat.

3
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Name the two main forms of energy discussed in thermochemistry.

Kinetic energy and potential energy.

4
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In molecules, what motions contribute to kinetic energy?

Translational, rotational, and vibrational motion.

5
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How is potential energy stored in chemical systems?

In chemical bonds; it decreases when bonds form and increases when bonds break.

6
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Does breaking a chemical bond absorb or release energy?

It absorbs energy, raising the potential energy of the system.

7
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Give the symbolic equation that represents breaking an O–H bond in water.

H₂O → H⁺ + OH⁻

8
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Does forming a chemical bond absorb or release energy?

It releases energy, lowering the system’s potential energy.

9
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Give the symbolic equation that represents forming an O–H bond in water.

H⁺ + OH⁻ → H₂O

10
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What is internal energy (U)?

The total energy (kinetic + potential) of a system.

11
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What is heat (q) in the context of energy transfer?

The flow of energy due to a temperature difference, always from hot to cold.

12
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What is work (w) in thermochemistry?

Energy transfer that results from a force acting over a distance.

13
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Write the formula that relates change in internal energy to heat and work.

ΔU = q + w

14
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According to sign conventions, what does q > 0 indicate?

Heat is absorbed by the system (endothermic process).

15
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According to sign conventions, what does w < 0 indicate for a gas?

Work is done by the system on the surroundings, as in gas expansion.

16
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Provide the equation for pressure–volume work in a gas reaction.

w = −PΔV

17
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Define enthalpy (H).

A thermodynamic property representing the heat content of a system at constant pressure (H = U + PV).

18
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How is the enthalpy change (ΔH) of a reaction at constant pressure calculated?

ΔH = Hproducts − Hreactants

19
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What sign of ΔH characterizes an exothermic reaction?

ΔH < 0 (heat is released).

20
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Give an example of an exothermic reaction and its ΔH value.

Combustion of methane: CH₄ + 2O₂ → CO₂ + 2H₂O, ΔH = −890 kJ.

21
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What sign of ΔH characterizes an endothermic reaction?

ΔH > 0 (heat is absorbed).

22
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Give an example of an endothermic process noted in the lecture.

Photosynthesis: 6CO₂ + 6H₂O + light → C₆H₁₂O₆ + 6O₂.

23
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State the First Law of Thermodynamics.

Energy is conserved; ΔU = q + w, and energy cannot be created or destroyed, only transferred or converted.

24
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State the Second Law of Thermodynamics in terms of entropy.

For any spontaneous process, the total entropy of the universe increases (ΔS_universe > 0).

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Provide an everyday example illustrating the Second Law.

Ice melting at room temperature: the organized solid becomes a more disordered liquid, increasing entropy.

26
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State the Third Law of Thermodynamics.

The entropy of a perfect crystal at absolute zero (0 K) is zero.