CHEM10007 · Fundamentals Of Chemistry
Thermochemistry, Calorimetry & Hess's Law
Energy bookkeeping for chemical reactions. You define heat, work and enthalpy, classify reactions as exothermic (ΔH < 0) or endothermic (ΔH > 0), and use the first law ΔU = q + w. The measurable skills are calorimetry (q = mcΔT and q = CcalΔT, linking measured heat to moles reacted and hence ΔH per mole) and Hess's Law (combining step enthalpies because ΔH is a path-independent state function).
What this chapter covers
- 01Kinetic vs potential energy; heat (q) vs temperature; work (w); system vs surroundings
- 02Exothermic (ΔH < 0, heat released) vs endothermic (ΔH > 0, heat absorbed)
- 03First law of thermodynamics: ΔU = q + w
- 04Enthalpy change ΔH of reaction
- 05Calorimetry: q = mcΔT for a solution and q = CcalΔT for a calorimeter
- 06Specific heat of water = 4.18 J g−1 °C−1
- 07Relating measured q to moles reacted to obtain ΔH per mole
- 08Hess's Law: reverse (flip sign) and scale (× n) component reactions; standard enthalpies of formation
Calorimetry — heat released raises water temperature
- 1 mark — heat from moles × molar enthalpyHeat released: q = 0.0150 mol × 2220 kJ mol−1 = 33.3 kJ = 3.33 × 104 J.
- 1 mark — correct rearrangementRearrange q = mcΔT for ΔT: ΔT = q ÷ (mc) = 3.33 × 104 ÷ (400 × 4.18).
- 1 mark — temperature changeΔT = 3.33 × 104 ÷ 1672 = 19.9 °C.
- 1 mark — final temperature and sanity checkTfinal = 20.0 + 19.9 = 39.9 °C — a reasonable rise for combustion heating water.
Key terms
- Enthalpy (ΔH)
- The heat exchanged by a reaction at constant pressure. Negative for exothermic reactions (heat released), positive for endothermic ones.
- Specific heat capacity (c)
- The energy needed to raise 1 g of a substance by 1 °C; for water, 4.18 J g−1 °C−1. Used in q = mcΔT.
- Calorimetry
- Measuring heat flow from a temperature change, using q = mcΔT for the solution or q = CcalΔT for the calorimeter, to find a reaction's ΔH.
- Hess's Law
- Because enthalpy is a state function, the ΔH of an overall reaction equals the sum of the ΔH values of any set of steps that add to it.
- State function
- A property (like enthalpy) whose value depends only on the current state, not the path taken — which is why Hess's Law works.
Thermochemistry, Calorimetry & Hess's Law FAQ
What is the difference between heat and temperature?
Temperature measures the average kinetic energy of particles; heat (q) is the energy transferred because of a temperature difference. A small mass at high temperature can carry less heat than a large mass at lower temperature.
How do I use Hess's Law?
Manipulate the given reactions so they add up to the target equation: reverse a reaction (flip the sign of its ΔH) and scale it (multiply ΔH by the same factor). Sum the adjusted ΔH values to get the overall enthalpy change.
Why must I convert kJ to J in calorimetry?
The specific heat of water is given in J g−1 °C−1, so q in q = mcΔT comes out in joules. Mixing kJ and J is a common error that throws the temperature change off by a factor of 1000.
Exam move
Keep your sign convention straight: a reaction releasing heat (ΔH < 0) warms the surroundings, so the water gains that energy. Practise the two-way link between measured q and ΔH per mole, since the calorimetry archetype recurs and connects to the F4 thermodynamics practical. For Hess's Law, set up a clean table of reversed/scaled reactions and their ΔH so the bookkeeping is transparent and easy to mark.
