1. Reaction rate
Reaction rate is how fast reactants are consumed or products are formed. Some reactions are essentially instantaneous (acid-base neutralization, explosions). Others are agonizingly slow (rusting, fossilization). The reason is collision theory.
2. Collision theory
For a reaction to occur, reactant particles must collide. But not every collision results in a reaction. A successful (productive) collision requires two things:
- Sufficient energy: the particles must collide with at least the activation energy (Ea), enough to break existing bonds.
- Proper orientation: particles must be lined up correctly so the right atoms can swap.
If either condition fails, the particles just bounce off and no reaction happens. Anything that increases the frequency or energy of collisions speeds up the reaction.
3. Factors that affect reaction rate
| Factor | Effect | Why (collision theory) |
|---|---|---|
| Temperature ↑ | Rate ↑ (a lot) | Higher KE means more collisions AND a greater fraction with enough energy |
| Concentration ↑ (or pressure ↑ for gases) | Rate ↑ | More particles in the same volume means more collisions per second |
| Surface area ↑ (for solids) | Rate ↑ | More particles exposed to reactant. Powder reacts faster than a chunk. |
| Catalyst added | Rate ↑ | Provides an alternative pathway with lower activation energy |
| Nature of reactants | Varies | Ionic reactions in solution are fast. Covalent bond breaking is slow. |
How a catalyst works
A catalyst is a substance that speeds up a reaction without being consumed. It works by providing a new reaction pathway with a lower activation energy. More particles have enough energy to "make it over the hill," so more collisions are productive. Importantly, a catalyst does not change the energy of the reactants or products, only the height of the activation energy barrier.
4. Potential energy diagrams
| Quantity | Symbol | How to read it off the PE diagram | Exothermic | Endothermic |
|---|---|---|---|---|
| Activation energy (forward) | Ea | Peak − reactant level | > 0 | > 0 |
| Activation energy (reverse) | Ea' | Peak − product level | larger than forward | smaller than forward |
| Heat of reaction | ΔH | Product − reactant level | ΔH < 0 (released) | ΔH > 0 (absorbed) |
| Catalyst effect | — | Lowers the peak (both Eₐ shrink) | ΔH unchanged | ΔH unchanged |
| Activated complex | — | Top of the hump (highest PE) | — | — |
A potential energy (PE) diagram shows how energy changes during a reaction. The x-axis is "reaction coordinate" (progress of reaction); the y-axis is potential energy.
Key features
Reactants (left plateau)
Starting energy of the system.
Products (right plateau)
Final energy of the system.
Activated complex (peak)
The unstable transition state at the top of the hill.
Activation energy (Ea)
Height from reactants up to the peak.
ΔH (heat of reaction)
Energy of products minus energy of reactants. Negative for exothermic, positive for endothermic.
Reverse Ea
Height from products up to the peak (going backward).
Exothermic vs endothermic shape
- Exothermic: products are lower than reactants. ΔH is negative. Energy released to surroundings. Think: combustion, neutralization. The curve goes up to the peak, then down to a lower plateau.
- Endothermic: products are higher than reactants. ΔH is positive. Energy absorbed from surroundings. Think: photosynthesis, melting ice, dissolving NH₄NO₃. The curve goes up to the peak, then down to a higher plateau.
Effect of a catalyst on a PE diagram
A catalyst lowers the activation energy peak only. The reactant level, product level, and ΔH all stay exactly the same. On the diagram, the same start and end points but a shorter hump in the middle.
5. Enthalpy and entropy
Enthalpy (ΔH) — energy
ΔH is the heat absorbed or released by a reaction at constant pressure. Reference Table I lists ΔH values for common reactions.
- ΔH negative = exothermic, energy released, products lower in energy
- ΔH positive = endothermic, energy absorbed, products higher in energy
Entropy (S) — disorder
Entropy is the measure of disorder, or the number of ways a system can be arranged. Higher entropy = more disorder. Reactions tend to proceed toward higher entropy (the second law of thermodynamics in plain English).
| Change | Entropy |
|---|---|
| Solid → liquid → gas | Increases |
| Few particles → many particles | Increases |
| Cold → hot | Increases |
| Dissolving a solid in water | Usually increases |
The pull of nature
Nature prefers two things: lower energy (exothermic, ΔH negative) and higher entropy (more disorder, ΔS positive). When a reaction has both, it is always spontaneous. When it has neither, it is never spontaneous. When it has one but not the other, temperature decides which wins.
6. Spontaneous reactions
A spontaneous reaction proceeds on its own once started, without continuous input. "Spontaneous" does not mean "fast." Rust is spontaneous but takes years. Diamond converting to graphite is also spontaneous but takes essentially forever.
| ΔH | ΔS | Spontaneous? |
|---|---|---|
| − (exothermic) | + (more disorder) | Always |
| + (endothermic) | − (less disorder) | Never |
| − (exothermic) | − (less disorder) | Yes at low T |
| + (endothermic) | + (more disorder) | Yes at high T |
7. Dynamic equilibrium
Many reactions are reversible, going forward and backward at the same time. When the forward and reverse reactions happen at equal rates, the system is at dynamic equilibrium. Concentrations of reactants and products stop changing, but the reactions are still occurring.
Three things at equilibrium
- The forward and reverse reaction rates are equal.
- The concentrations of reactants and products are constant (not necessarily equal).
- The system must be closed (no reactants or products entering or leaving).
Types of equilibrium
Phase equilibrium
Water in a sealed bottle: rate of evaporation = rate of condensation. The water level stays constant.
Solution equilibrium
Saturated solution: rate of dissolving = rate of recrystallizing. The undissolved solid stays the same size.
Chemical equilibrium
Reversible reaction like N₂ + 3H₂ ⇌ 2NH₃: forward and reverse reactions at equal rates.
8. Le Chatelier's Principle
Le Chatelier's Principle is the most testable idea in this unit. When a stress is applied to a system at equilibrium, the system shifts in the direction that relieves the stress.
Consider the equilibrium: N₂(g) + 3H₂(g) ⇌ 2NH₃(g) + heat
Stress 1: Change concentration
| Stress | Shift | Why |
|---|---|---|
| Add N₂ (reactant) | Right (toward products) | System "uses up" the added N₂ |
| Remove NH₃ (product) | Right (toward products) | System replaces what was removed |
| Add NH₃ (product) | Left (toward reactants) | System reduces the excess product |
| Remove H₂ (reactant) | Left (toward reactants) | System replaces what was removed |
Stress 2: Change temperature
Treat heat as a reactant or product depending on whether the forward reaction is exothermic or endothermic. Here, the forward reaction is exothermic (heat is on the product side).
- Increase T (add heat): shifts away from heat, so left (toward reactants). Less ammonia.
- Decrease T (remove heat): shifts toward heat, so right (toward products). More ammonia.
Stress 3: Change pressure (gases only)
Count moles of gas on each side. Increasing pressure shifts toward the side with fewer moles of gas. Decreasing pressure shifts toward the side with more moles of gas.
In our example, reactant side has 4 mol of gas (1 N₂ + 3 H₂), product side has 2 mol of gas (2 NH₃). Increasing pressure shifts right (fewer moles relieves the squeeze).
Stress 4: Add a catalyst
A catalyst speeds up the forward and reverse reactions equally. It helps the system reach equilibrium faster, but it does not shift the equilibrium position. The final concentrations are unchanged.
Inert gas trap
Adding an inert (non-reacting) gas like argon to a gas-phase equilibrium at constant volume has no effect on the equilibrium position. The partial pressures of the reacting gases don't change. The exam loves this one.
Mnemonic: "system fights back"
Whatever you do to an equilibrium, it tries to undo. Add stuff → it consumes it. Remove stuff → it makes more. Heat it → it cools itself by absorbing heat. Cool it → it warms itself by releasing heat. Squeeze it → it shrinks. Expand it → it puffs up. Always fighting back.
Key terms
| Reaction rate | How fast reactants are consumed or products formed. |
| Activation energy (Ea) | Minimum energy needed for a productive collision. |
| Activated complex | Unstable transition state at the peak of the PE diagram. |
| Catalyst | Substance that speeds up a reaction by lowering activation energy. Not consumed. |
| Exothermic | Releases energy. ΔH negative. Products lower than reactants. |
| Endothermic | Absorbs energy. ΔH positive. Products higher than reactants. |
| Enthalpy (ΔH) | Heat of reaction. Listed for common reactions on Table I. |
| Entropy (S) | Measure of disorder or randomness in a system. |
| Dynamic equilibrium | State where forward and reverse rates are equal; concentrations constant. |
| Le Chatelier's Principle | An equilibrium responds to a stress by shifting to oppose it. |
| Reversible reaction | Reaction that can proceed in both forward and reverse directions. |
| Spontaneous reaction | Proceeds on its own without continuous external input. Not necessarily fast. |
Practice questions
Q1. Which factor will NOT increase the rate of a chemical reaction?
Answer: (4) decreasing the concentration of reactants. Fewer reactant particles per unit volume means fewer collisions per second, so the rate decreases. The other three all increase collision frequency or the fraction of effective collisions.
Q2. A catalyst increases the rate of a reaction by:
Answer: (3) providing an alternative pathway with a lower activation energy. Catalysts work by offering a different reaction route over a shorter energy "hill." More particles have enough energy to make it over, so more collisions are productive. A catalyst does NOT change ΔH, temperature, or concentration.
Q3. Given the reaction at equilibrium: N₂(g) + 3H₂(g) ⇌ 2NH₃(g) + heat. What happens to the concentration of NH₃ when the temperature is increased at constant volume?
Answer: (2) It decreases. The forward reaction is exothermic (heat on the product side). Increasing temperature is the stress of adding heat. By Le Chatelier, the system shifts away from the added heat, toward the reactants (left). This consumes NH₃, so its concentration decreases.
Q4. (Part B-2) Sketch and label a potential energy diagram for an exothermic reaction. Label the following: reactants, products, activated complex, activation energy (Ea) for the forward reaction, and ΔH. Explain how the diagram would change if a catalyst were added.
Sample full-credit response: The PE diagram for an
exothermic reaction starts at the reactants' energy level on the left, rises to a peak
(the activated complex), then drops down to a lower energy level on the right (the
products). The activation energy (Ea for the forward reaction) is the
vertical distance from the reactant level up to the peak. The ΔH (heat of reaction) is
the difference between product energy and reactant energy; for an exothermic reaction,
ΔH is negative because the products are lower than the reactants.
Adding a catalyst lowers the activation energy peak, making it a shorter hump. However, the
reactant level, product level, and ΔH are all unchanged. The catalyst only provides a new
pathway with a lower energy barrier; it does not change the overall energy difference
between reactants and products.
Q5. (Part C) The Haber process is used industrially to make ammonia: N₂(g) + 3H₂(g) ⇌ 2NH₃(g) + 91.8 kJ. Explain how each of the following changes would affect the yield (amount produced) of NH₃. Justify each answer using Le Chatelier's Principle. (a) Increasing the concentration of N₂. (b) Increasing the temperature. (c) Increasing the pressure. (d) Adding a catalyst.
Sample full-credit response:
(a) Increasing [N₂]: The system is stressed by extra reactant. By Le
Chatelier, the equilibrium shifts to the right to consume the added N₂. This increases
the yield of NH₃.
(b) Increasing temperature: The forward reaction is exothermic, so heat
is effectively a product. Adding heat stresses the product side, so the equilibrium shifts
to the left to absorb the added heat. This decreases the yield of NH₃.
(c) Increasing pressure: The reactant side has 4 moles of gas (1 N₂ + 3 H₂)
and the product side has 2 moles of gas (2 NH₃). Increasing pressure stresses the system,
and by Le Chatelier, the equilibrium shifts toward the side with fewer moles of gas
(the product side) to relieve the pressure. This increases the yield of NH₃.
(d) Adding a catalyst: A catalyst speeds up both the forward and reverse
reactions equally. It helps the system reach equilibrium faster but does not shift the
equilibrium position. The final yield of NH₃ is unchanged, just achieved more quickly.
(Industrially, the Haber process compromises: it uses moderate temperatures and very high
pressure, plus an iron catalyst, to optimize both rate and yield.)