Calorimetry
Calorimetry is the experimental science of measuring the heat exchanged during chemical reactions, physical changes, or heat capacity determinations using an instrument called a calorimeter. The fundamental principle is conservation of energy: heat released by the reaction equals heat absorbed by the calorimeter and its contents. Two main types are constant-pressure calorimetry (coffee-cup calorimeter) and constant-volume calorimetry (bomb calorimeter), each suited to different experimental conditions.
Key Formula
q = m × c × ΔT
LaTeX: q = mc\Delta T
| Symbol | Meaning | Unit |
|---|---|---|
| q | Heat transferred | J or kJ |
| m | Mass of the substance | g |
| c | Specific heat capacity | J/(g·°C) |
| ΔT | Change in temperature (T_final − T_initial) | °C or K |
Worked Example
Problem
When 4.00 g of NaOH dissolves in 50.0 g of water in a coffee-cup calorimeter, the temperature rises from 22.0 °C to 37.5 °C. Calculate the heat released per mole of NaOH dissolved. Specific heat of solution = 4.18 J/(g·°C). Molar mass of NaOH = 40.0 g/mol.
Solution
Step 1: Total mass of solution = 4.00 + 50.0 = 54.0 g. Step 2: ΔT = 37.5 − 22.0 = 15.5 °C. Step 3: q_solution = m × c × ΔT = 54.0 × 4.18 × 15.5 = 3498.7 J = 3.499 kJ. Step 4: Since temperature increased, heat was released by reaction: q_rxn = −3.499 kJ. Step 5: Moles of NaOH = 4.00 / 40.0 = 0.100 mol. Step 6: ΔH_solution = −3.499 / 0.100 = −35.0 kJ/mol.
Answer
ΔH_solution = −35.0 kJ/mol
Comparison of Coffee-Cup and Bomb Calorimeters
| Feature | Coffee-Cup Calorimeter | Bomb Calorimeter |
|---|---|---|
| Pressure condition | Constant pressure (open) | Constant volume (closed) |
| Measures | ΔH (enthalpy) | ΔU (internal energy) |
| Typical use | Aqueous reactions, neutralization | Combustion reactions |
| Construction | Styrofoam cup, thermometer | Steel vessel, water bath |
| Accuracy | Moderate | High |
| Correction needed | None (qp = ΔH) | ΔH = ΔU + ΔngRT |
Interactive Tools
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Related Terms
Specific Heat (chemistry)
Specific heat capacity (c) is the amount of heat energy required to raise the temperature of one gram of a substance by one degree Celsius (or one Kelvin). It is an intrinsic physical property of a material that reflects how resistant a substance is to temperature change, with water having the exceptionally high value of 4.184 J/(g·°C). Specific heat is fundamental to calorimetry calculations and explains phenomena such as coastal climate moderation due to the ocean's high heat capacity.
Standard Enthalpy
Standard enthalpy refers to the enthalpy change of a process measured under standard conditions: 298.15 K (25 °C) and 1 bar (approximately 1 atm) pressure, with all substances in their standard states. Standard enthalpy values provide a universal reference for comparing thermochemical data across different reactions and sources. The symbol ΔH° (read "delta H naught" or "delta H standard") denotes a standard enthalpy change.
Hess's Law
Hess's Law states that the total enthalpy change of a chemical reaction is independent of the pathway taken, depending only on the initial and final states of the system. This principle allows chemists to calculate enthalpy changes for reactions that are difficult or impossible to measure directly by combining known thermochemical equations. It is a direct consequence of the First Law of Thermodynamics, making enthalpy a state function.
From Latin 'calor' (heat) and Greek 'metron' (measure). The word 'calorimeter' was coined in the late 18th century; Antoine Lavoisier and Pierre-Simon Laplace performed early calorimetric experiments around 1780.