1. What is thermodynamics?
Thermodynamics studies energy, heat, work, temperature and their relationships. In chemistry, it is especially useful for analysing reactions, phase changes and energy transfer.
System
The part of the universe being studied, such as the chemicals reacting inside a flask.
Surroundings
Everything outside the system that may exchange energy or matter with it.
Boundary
The real or imaginary surface separating the system from its surroundings.
State variables
Properties such as pressure, volume, temperature and internal energy that describe the state of a system.
Types of systems
| System | Matter exchange | Energy exchange |
|---|---|---|
| Open | Yes | Yes |
| Closed | No | Yes |
| Isolated | No | No |
2. Heat and work
Heat and work are two ways energy can be transferred across a system's boundary. They are not properties stored in a system; instead, they describe energy transfer during a process.
Heat (q)
Energy transferred because of a temperature difference. Heat flows spontaneously from a hotter body to a colder one.
Work (w)
Energy transferred when a force acts through a displacement or through another organised process, such as electrical work.
Sign convention used here
- q > 0: heat enters the system.
- q < 0: heat leaves the system.
- w > 0: work is done on the system.
- w < 0: work is done by the system.
Important: Some textbooks use a different convention for work. Always check the convention before applying an equation. The following lessons use the chemistry convention, where ΔU = q + w.
3. Internal energy and the first law
Internal energy, U, is the total microscopic energy of a system, including contributions from particle motion and interactions. The change in internal energy depends on the initial and final states.
First law of thermodynamics
The change in internal energy equals heat supplied to the system plus work done on the system.
Example
Suppose a system absorbs 500 J of heat and does 200 J of work on its surroundings.
- Heat absorbed: q = +500 J.
- Work done by the system: w = −200 J.
- Change in internal energy: ΔU = 500 + (−200) = +300 J.
The internal energy increases by 300 J.
Pressure-volume work
For expansion or compression against a constant external pressure, Pₑₓₜ is the external pressure and ΔV = V₂ − V₁. Use compatible units.
- Expansion: ΔV is positive, so w is negative.
- Compression: ΔV is negative, so w is positive.
- At constant volume, pressure-volume work is zero.
4. Enthalpy and heat of reaction
Enthalpy, H, is a thermodynamic state function defined as internal energy plus pressure multiplied by volume.
At constant pressure, when only pressure-volume work is involved, the heat exchanged is equal to the enthalpy change: qₚ = ΔH.
Exothermic and endothermic reactions
Exothermic
Releases heat to the surroundings. At constant pressure, ΔH is negative.
Endothermic
Absorbs heat from the surroundings. At constant pressure, ΔH is positive.
Enthalpy change for gaseous reactions
For an ideal-gas reaction at the same temperature, Δn(g) is the total stoichiometric amount of gaseous products minus that of gaseous reactants. This relation assumes the usual ideal-gas treatment.
Watch the sign: A negative ΔH indicates an exothermic reaction under the stated conditions; a positive ΔH indicates an endothermic reaction.
5. Calorimetry and heat capacity
Calorimetry measures heat transferred during a physical or chemical process. In a simple temperature-change calculation, heat depends on mass, specific heat capacity and temperature change.
Heat from a temperature change
- q = heat transferred
- m = mass
- c = specific heat capacity
- ΔT = T₂ − T₁
Worked example
A 100 g sample of water warms from 20°C to 25°C. Take the specific heat capacity of water as 4.18 J g⁻¹ °C⁻¹.
- Mass: m = 100 g.
- Temperature change: ΔT = 25 − 20 = 5°C.
- Heat absorbed: q = 100 × 4.18 × 5 = 2090 J.
The water absorbs approximately 2.09 kJ of heat.
Heat capacity
Heat capacity is the heat required to raise the temperature of an entire object or system by one degree. Unlike specific heat capacity, it depends on the amount of material.
In an ideal insulated calorimeter, heat lost by one part equals heat gained by the other. Real experiments may also need to account for heat absorbed by the calorimeter and heat exchanged with the surroundings.
6. Hess's law
Hess's law states that the total enthalpy change for a reaction depends only on the initial and final states, not on the pathway taken. This works because enthalpy is a state function.
If a reaction can be represented as a series of reactions, their enthalpy changes can be added to find the overall change.
Rules for using Hess's law
- Reverse a reaction and its enthalpy change changes sign.
- Multiply a reaction equation by a factor and multiply its ΔH by the same factor.
- Add reaction equations and cancel species that appear on both sides.
- Add the adjusted enthalpy changes to obtain the overall ΔH.
Using standard enthalpies of formation
Here, ν represents each species' stoichiometric coefficient. Standard enthalpies of formation must be multiplied by their coefficients before the sums are calculated.
7. Entropy and spontaneity
Entropy, S, is a state function related to how energy and matter can be distributed among the possible microscopic arrangements of a system. It helps explain why some processes occur spontaneously under particular conditions.
Entropy change
Entropy often increases when a substance becomes more dispersed, such as when a liquid evaporates or a gas expands. The exact change depends on the process.
Second law
For a spontaneous process, the total entropy change of the system and surroundings is positive. For a reversible process, the total entropy change is zero.
Gibbs free energy
At constant temperature and pressure, Gibbs free energy helps assess spontaneity when the relevant conditions are specified. Use absolute temperature in kelvin and consistent units.
| Gibbs free-energy change | Interpretation at constant T and P |
|---|---|
| ΔG < 0 | Forward process is thermodynamically spontaneous. |
| ΔG > 0 | Forward process is non-spontaneous under those conditions. |
| ΔG = 0 | System is at equilibrium. |
Spontaneous does not mean fast. Thermodynamics indicates whether a process is energetically favourable under specified conditions; kinetics describes how quickly it occurs.
8. Key ideas to remember
- A system is the part being studied; everything else is its surroundings.
- Heat and work describe energy transfer across the system boundary.
- Using the chemistry sign convention, ΔU = q + w.
- At constant pressure with only pressure-volume work, qₚ = ΔH.
- Use q = mcΔT for simple temperature-change calculations.
- Hess's law follows from enthalpy being a state function.
- Entropy and Gibbs free energy help analyse spontaneity and equilibrium.
- A spontaneous reaction is not necessarily a fast reaction.