Questions: Thermochemistry: Enthalpy and Heat of Reaction

Question 1

When ammonium nitrate dissolves in water, the solution becomes noticeably cold. What does this tell you about the enthalpy change, and what is happening energetically?

• A. ΔH < 0 — the reaction is exothermic and releases heat, cooling the surroundings
• B. ΔH > 0 — the reaction is endothermic and absorbs heat from the surroundings, making them feel cold
• C. ΔH = 0 — the temperature change is a physical change, not a chemical one, so enthalpy is unchanged
• D. ΔH < 0 — the solution cools because dissolved ions have less energy than solid ammonium nitrate

Answer: (B)

When the solution gets cold, it is losing heat to the reaction — the reaction is absorbing thermal energy from the surroundings. From the system's perspective (the dissolving process), energy flows in, so ΔH > 0 (endothermic). This is the basis of instant cold packs. Option A is the classic reversal error: exothermic reactions release heat and warm the surroundings, not cool them. The sign convention is always from the system's viewpoint: positive ΔH means the system absorbed energy, leaving the surroundings cooler.

Question 2

You want ΔH for: C(s) + O₂(g) → CO₂(g). You have: (1) C(s) + ½O₂(g) → CO(g), ΔH₁ = −110 kJ; (2) CO(g) + ½O₂(g) → CO₂(g), ΔH₂ = −283 kJ. What is ΔH for the target reaction?

• A. −173 kJ (ΔH₂ − ΔH₁)
• B. −393 kJ (ΔH₁ + ΔH₂)
• C. +393 kJ (reversing the sum)
• D. −283 kJ (only the final step matters)

Answer: (B)

Adding reactions (1) and (2) gives: C(s) + ½O₂ + CO(g) + ½O₂ → CO(g) + CO₂(g). The CO(g) cancels, leaving C(s) + O₂(g) → CO₂(g) — exactly the target. Because enthalpy is a state function, ΔH adds: −110 + (−283) = −393 kJ. This is Hess's Law in action: the path doesn't matter, only the initial and final states. Option A is the subtraction error; option D ignores the first step entirely.

Question 3

Enthalpy is a state function, meaning the value of ΔH for a reaction is the same whether it occurs in one step or through a series of intermediate reactions.

• T. True
• F. False

Answer: True

This is precisely what makes Hess's Law work. Because enthalpy depends only on the current state of the system (pressure, temperature, composition) and not on how the system arrived at that state, ΔH between two states is path-independent. Whether combustion of carbon proceeds directly to CO₂ or goes through CO first, the total enthalpy change is the same. This allows chemists to calculate ΔH for reactions that are too slow, dangerous, or complex to measure directly.

Question 4

A reaction with ΔH = −500 kJ will necessarily proceed faster than a reaction with ΔH = −50 kJ.

• T. True
• F. False

Answer: False

Enthalpy change (ΔH) describes thermodynamics — the difference in energy between products and reactants — not kinetics (how fast the reaction proceeds). Reaction rate is determined by the activation energy (the energy barrier between reactants and the transition state), not by ΔH. Many reactions with very large negative ΔH are extremely slow at room temperature (e.g., iron rusting, wood spontaneously combusting in air). A catalyst speeds up a reaction without changing ΔH. Thermodynamics and kinetics are independent.

Question 5

Explain the difference between heat and temperature, and describe how a calorimetry experiment uses q = mcΔT to measure the enthalpy change of a reaction.

Model answer: Heat (q) is energy transferred between objects due to a temperature difference, measured in joules. Temperature is the average kinetic energy of particles in an object, measured in °C or K. In calorimetry, a reaction occurs in a known mass (m) of water or solution with known specific heat capacity (c). The temperature change (ΔT) of the water is measured. q = mcΔT gives the heat absorbed or released by the water, which equals (with sign reversed) the heat released or absorbed by the reaction. This connects the abstract concept of ΔH to a measurable physical change.

The key distinction is that heat is a process (energy in transit) while temperature is a property (a state of matter). A large swimming pool at 20°C contains far more thermal energy than a cup of boiling water, but the cup is at higher temperature. Calorimetry exploits q = mcΔT to make ΔH measurable: the reaction transfers heat to or from the calorimeter, and we observe the temperature change that results. The water acts as a thermometric gauge — its temperature change is the observable proxy for the invisible enthalpy change of the chemical reaction.