Quick Summary
- Thermodynamics examines energy flow between systems and surroundings
- Three main properties: enthalpy (heat content), entropy (disorder), and free energy (useful work)
- Systems can be open, closed, or isolated based on energy and matter exchange
- First Law states energy cannot be created or destroyed, only transformed
- Second Law states natural processes increase total entropy (disorder)
Understanding Chemical Thermodynamics
When you burn firewood, you feel heat. When ice melts in your hand, it feels cold. These everyday experiences involve energy changes that thermodynamics helps us understand and measure.
Chemical thermodynamics gives us tools to answer important questions: Will this reaction happen on its own? How much heat will it produce? Can we get useful work from it? These questions matter in industries across Nigeria, from cement factories in Calabar to oil refineries in Warri.
Basic Thermodynamic Terms
System, Surroundings, and Boundary
The system is the specific part we want to study. It could be chemicals in a beaker, fuel in an engine, or food cooking in a pot. Everything else is the surroundings. The boundary separates them.
Think of boiling water in a covered pot. The water is the system. The air in your kitchen, the stove, and everything around the pot are the surroundings. The pot itself is the boundary.
Types of Systems
| System Type | Energy Exchange | Matter Exchange | Example |
|---|---|---|---|
| Open System | Yes | Yes | Boiling water in open pot (steam escapes, heat transfers) |
| Closed System | Yes | No | Sealed bottle of soft drink (heat moves in/out, but liquid stays inside) |
| Isolated System | No | No | Perfectly insulated thermos flask (ideal case, rare in practice) |
Key Thermodynamic Properties
Enthalpy (ΔH)
Enthalpy measures the total heat content of a system. The symbol ΔH (delta H) represents the change in enthalpy during a reaction.
Exothermic reactions release heat. ΔH is negative. Examples include burning petrol, respiration, and neutralization. When PHCN generators burn diesel, they release heat (exothermic).
Endothermic reactions absorb heat. ΔH is positive. Examples include photosynthesis, melting ice, and cooking food. When you cook beans, the pot absorbs heat from the fire (endothermic).
Entropy (ΔS)
Entropy measures disorder or randomness in a system. Nature prefers disorder. A neat room becomes messy on its own, but never the reverse without effort.
Gases have high entropy (molecules move freely). Liquids have medium entropy. Solids have low entropy (particles are fixed). When ice melts to water, entropy increases. When water boils to steam, entropy increases more.
The symbol ΔS represents entropy change. Positive ΔS means more disorder. Negative ΔS means more order.
Free Energy (ΔG)
Free energy (also called Gibbs free energy) measures energy available for useful work. It combines enthalpy and entropy using this equation:
ΔG = ΔH – TΔS
Where T is temperature in Kelvin. This equation tells us if a reaction will happen spontaneously:
- ΔG negative: Reaction happens spontaneously (like iron rusting)
- ΔG positive: Reaction needs energy input (like splitting water into hydrogen and oxygen)
- ΔG = 0: Reaction is at equilibrium (forward and reverse rates are equal)
Laws of Thermodynamics
First Law: Conservation of Energy
Energy cannot be created or destroyed. It only changes form. When you charge your phone, electrical energy becomes chemical energy in the battery. When you use the phone, chemical energy becomes electrical energy again.
In chemistry, this means: Energy of reactants + energy absorbed = Energy of products + energy released
Second Law: Entropy Always Increases
The total entropy of the universe always increases for spontaneous processes. Things naturally move toward disorder. This is why LASTMA must keep directing traffic – cars do not organize themselves.
This law explains why some reactions happen on their own while others need energy input. Reactions that increase total entropy (system plus surroundings) happen spontaneously.
Third Law: Zero Entropy at Absolute Zero
A perfect crystal at absolute zero (0 Kelvin or -273°C) has zero entropy. This is a theoretical reference point. In practice, we cannot reach absolute zero.
State Functions vs Path Functions
State functions depend only on the current state, not how you got there. Examples: enthalpy, entropy, temperature, pressure. Going from Lagos to Abuja, your final position (Abuja) does not depend on whether you flew, drove, or took a train.
Path functions depend on the route taken. Examples: heat and work. The fuel you use getting from Lagos to Abuja depends on your route – flying uses different fuel than driving.
Practical Applications in Nigeria
Thermodynamics has many applications:
- Petroleum industry: Predicting whether crude oil reactions will occur in refineries
- Power generation: Calculating efficiency of PHCN turbines and generators
- Food preservation: Understanding why refrigeration slows spoilage
- Manufacturing: Optimizing cement production in Dangote factories
- Environmental science: Studying pollution reactions in Lagos lagoon
Common Exam Mistakes
WAEC Chief Examiner Reports
- Confusing ΔH and ΔG: Enthalpy is heat content, free energy is useful work available. They are not the same.
- Wrong signs: Exothermic has negative ΔH (releases heat), not positive. Many students reverse this.
- Forgetting temperature in ΔG equation: The equation is ΔG = ΔH – TΔS. Many students write ΔG = ΔH – ΔS (missing T).
- Mixing up system types: Closed systems allow energy exchange but not matter exchange. Open systems allow both.
- Cannot explain spontaneity: Negative ΔG means spontaneous. Positive ΔG means non-spontaneous. Know this relationship.
- Poor understanding of entropy: Gas has higher entropy than liquid, which has higher entropy than solid. Many students reverse this order.
Practice Questions
Multiple Choice Questions
1. Which of the following processes has a negative entropy change?
(a) Ice melting to water
(b) Water boiling to steam
(c) Steam condensing to water ✓
(d) Solid sugar dissolving in water
2. A reaction has ΔH = -50 kJ and ΔS = +100 J/K at 300K. What is ΔG?
(a) +20 kJ
(b) -20 kJ
(c) -80 kJ ✓
(d) +80 kJ
Hint: ΔG = ΔH – TΔS = -50 – (300 × 0.1) = -50 – 30 = -80 kJ
3. Which type of system allows both energy and matter exchange with surroundings?
(a) Closed system
(b) Isolated system
(c) Open system ✓
(d) Adiabatic system
4. A reaction with ΔG = +25 kJ is best described as:
(a) Spontaneous at all temperatures
(b) Non-spontaneous at all temperatures ✓
(c) At equilibrium
(d) Exothermic
Essay Questions
1. (a) Define the following terms: (i) Enthalpy (ii) Entropy (iii) Free energy [6 marks]
(b) State the first and second laws of thermodynamics. [4 marks]
Examiner’s Tip: For part (a), give clear definitions in your own words. Do not just copy symbols. For part (b), state the laws simply, then explain what they mean.
2. (a) Distinguish between open, closed, and isolated systems. Give one example of each. [6 marks]
(b) Explain why reactions with negative ΔG occur spontaneously. [4 marks]
Mark Allocation Tip: Each system type plus example = 2 marks. Use clear everyday examples like those in this article.
3. A reaction has ΔH = -100 kJ/mol and ΔS = -200 J/K/mol.
(a) Calculate ΔG at 300K. [3 marks]
(b) Will this reaction be spontaneous at 300K? Explain. [2 marks]
(c) At what temperature will this reaction be at equilibrium? [5 marks]
Calculation Steps: (a) Convert J to kJ, substitute in ΔG = ΔH – TΔS. (b) Check sign of ΔG. (c) Set ΔG = 0 and solve for T.
Memory Aids
Mnemonics to Remember
For System Types (OCI):
Open – Over-sharing (shares both energy and matter)
Closed – Container (keeps matter in, but energy escapes)
Isolated – Independent (shares nothing)
For Entropy Order (GAS):
Gas (highest entropy)
Aqua/liquid (medium)
Solid (lowest)
For ΔG Equation:
“Good Heat Takes away Stress”
G = H – T × S
For Spontaneity:
“Negative Goes” (Negative ΔG = reaction goes spontaneously)
Related Topics
- Heat of Reaction and Enthalpy Changes
- Hess’s Law of Constant Heat Summation
- Spontaneous and Non-Spontaneous Reactions
- Chemical Equilibrium and Le Chatelier’s Principle
- Energy Changes in Chemical Reactions