Standard Enthalpy Change (ΔH⦵) and Calorimetry
Quick Notes:
- ΔH⦵ is the standard enthalpy change: heat transferred under standard conditions (298 K, 100 kPa, 1 mol dm⁻³, substances in their standard states).
- Units: kJ mol⁻¹
- Use two key equations:
- Q = mcΔT
to find heat energy (Q) transferred - ΔH = −Q/n
to calculate enthalpy change per mole of substance reacted
- Q = mcΔT
- Specific heat capacity of water, c = 4.18 J g⁻¹ K⁻¹ (given in data booklet).
Full Notes:
Note – Enthalpy (H) is the total heat energy of a chemical system at constant pressure. ΔH is the change in enthalpy when a chemical reaction occurs.
What Is ΔH⦵?
The standard enthalpy change of reaction, ΔH⦵, is the heat energy transferred during a reaction under standard conditions, with all substances in their standard states.
Standard conditions:
- Temperature: 298 K (25 °C)
- Pressure: 100 kPa
- Concentration (for solutions): 1 mol dm⁻³
- Substances in their standard physical states
Calorimetry and the Equation Q = mcΔT
Calorimetry is an experimental technique used to measure enthalpy changes.
The key equation used is:
Q = mcΔT
- Q = heat energy (J)
- m = mass of the substance being heated (usually water or solution), in grams
- c = specific heat capacity (for water, 4.18 J g⁻¹ K⁻¹)
- ΔT = temperature change (in K or °C – same difference)
This gives the total heat change in the experiment.
From Heat to Enthalpy: ΔH = −Q/n
Once you’ve calculated Q, we can find ΔH using:
ΔH = −Q / n
- ΔH = enthalpy change per mole (kJ mol⁻¹)
- Q = heat energy (convert to kJ if needed)
- n = number of moles of limiting reagent
The negative sign indicates whether the reaction is exothermic (−ΔH) or endothermic (+ΔH).
A student mixes 50.0 g of water with a salt, and the temperature rises by 6.0 °C. Calculate ΔH⦵ if 0.020 mol of salt reacted.
- Step 1: Calculate Q
Q = mcΔT = 50.0 × 4.18 × 6.0 = 1254 J = 1.254 kJ - Step 2: Calculate ΔH
ΔH = −Q / n = −1.254 / 0.020 = −62.7 kJ mol⁻¹
Summary Table
| Equation | Use | Units | Notes |
|---|---|---|---|
| Q = mcΔT | Find heat energy transferred | Q in joules (J) | Use c = 4.18 J g⁻¹ K⁻¹ for water |
| ΔH = −Q / n | Find enthalpy change per mole | ΔH in kJ mol⁻¹ | Convert Q to kJ; use limiting reagent |
How can the enthalpy change for combustion reactions, such as for alcohols or food, be investigated experimentally?
The enthalpy of combustion is the energy change when one mole of a substance burns completely in oxygen.
Method:
- Measure a known volume of water in a calorimeter (beaker or copper can).
- Record the starting temperature of the water.
- Weigh the spirit burner containing the fuel.
- Light the burner and allow it to heat the water.
- Stir and measure the final temperature of the water.
- Reweigh the burner to determine mass of fuel burned.
- Calculate q using Q = mcΔT, then use ΔH = q / n.
Sources of Error:
- Heat loss to surroundings (e.g., air, beaker).
- Incomplete combustion (producing CO instead of CO2).
- Evaporation of fuel from the wick.
Why do calorimetry experiments typically measure a smaller temperature change than expected?
Calorimetry often gives lower temperature changes than theoretical values because of heat losses to the surroundings, incomplete combustion, or absorption of heat by the apparatus. These experimental limitations mean not all the energy released is transferred to the water or measured system.
Summary
- ΔH⦵ is the standard enthalpy change under standard conditions (298 K, 100 kPa, 1 mol dm⁻³, substances in their standard states).
- Calorimetry uses Q = mcΔT to calculate heat transferred.
- ΔH = −Q / n converts heat to enthalpy change per mole.
- Experimental values may be lower than theoretical due to heat losses and incomplete combustion.