Difficulty: Intermediate | Prerequisites: General chemistry foundations, basic thermodynamics (Chapter 4 or equivalent)
Big picture: Chemical kinetics is the study of how fast reactions happen and what controls that speed. Where thermodynamics tells you whether a reaction can occur, kinetics tells you whether it will occur on any useful timescale. This chapter sits at the intersection of energy, molecular collisions and reaction design. If you have not reviewed Gibbs free energy and equilibrium basics, do that first.
Chemical kinetics is about reaction speed: how quickly reactants turn into products, and what you can do to change that speed. The rate depends on molecular collisions, activation energy, temperature, concentration and whether a catalyst is present. Thermodynamics says if a reaction is favourable; kinetics says how long you will be waiting.
Gibbs free energy (ΔG)
The thermodynamic quantity that determines whether a reaction is spontaneous. A negative ΔG means the reaction is exergonic (energy-releasing); a positive ΔG means it is endergonic (energy-absorbing). In simple terms, ΔG tells you whether a reaction wants to happen, not whether it happens quickly.
Activation energy (Eₐ)
The minimum energy that colliding molecules must possess for a reaction to proceed. Think of it as the hill reactants have to climb before they can roll down to products.
Rate constant (k)
A proportionality constant in the rate law that captures how fast a particular reaction runs at a given temperature. A larger k means a faster reaction under the same conditions.
Frequency factor (A) / pre-exponential factor
The term in the Arrhenius equation representing how often molecules collide with the correct orientation. In simple terms, it is the "best-case" rate if every collision had enough energy.
Collision theory
The model stating that a reaction occurs only when molecules collide with sufficient energy and proper orientation. Think of it as two conditions that must both be met: hard enough and aimed right.
Transition state / activated complex
A fleeting, highest-energy arrangement of atoms that exists at the top of the energy barrier between reactants and products. It cannot be isolated or observed directly.
Reaction intermediate
A species that is produced in one elementary step of a mechanism and consumed in a later step, so it never appears in the overall balanced equation. Think of it as a temporary molecule that helps get from start to finish.
Rate-determining step (RDS)
The slowest elementary step in a multi-step reaction mechanism. It acts as a bottleneck: the overall reaction can only go as fast as this step allows.
Catalyst
A substance that increases the rate of a reaction by lowering the activation energy, without being consumed in the process. It speeds things up but does not change what products form or whether the reaction is thermodynamically favourable.
Homogeneous catalysis
Catalysis where the catalyst is in the same phase as the reactants (e.g. all dissolved in solution).
Heterogeneous catalysis
Catalysis where the catalyst is in a different phase from the reactants (e.g. a solid catalyst with gaseous reactants).
Exergonic reaction
A reaction with a negative ΔG that releases free energy to its surroundings.
Endergonic reaction
A reaction with a positive ΔG that absorbs free energy from its surroundings.
ΔG is the difference between the activation energy of the forward reaction and the reverse reaction: ΔG = Eₐ(forward) − Eₐ(reverse).
A negative ΔG (exergonic) means the reaction releases energy overall.
A positive ΔG (endergonic) means the reaction absorbs energy overall.
A spontaneous reaction (negative ΔG) is not necessarily a fast reaction. Spontaneity and speed are separate questions entirely.
Enzymes serve as catalysts for most biochemical reactions, making otherwise sluggish spontaneous processes run at biologically useful speeds.
For a reaction to occur, reactant molecules must physically collide.
Two conditions must be satisfied in every productive collision:
The molecules must be oriented correctly so the right bonds can break and form.
The collision must carry at least the activation energy (Eₐ).
Most collisions fail one or both conditions, which is why reactions are slower than the raw collision frequency would suggest.
The rate can be expressed as Rate = Zf, where Z is the number of collisions per second and f is the fraction of those collisions that are effective.
k = A × e^(−Eₐ / RT), where:
k = rate constant
A = frequency factor (pre-exponential factor)
Eₐ = activation energy
R = gas constant (8.314 J/mol·K)
T = temperature in kelvin
Key relationships to remember:
Increasing A increases k (more frequent properly oriented collisions).
Increasing T increases k. The exponent −Eₐ/RT becomes less negative as T rises, so the exponential term grows.
Increasing the number of molecules in the vessel increases A (more chances for collisions).
The transition state (activated complex) sits at the peak of the energy profile between reactants and products.
It has higher energy than both reactants and products.
Once formed, it can proceed forward to products or collapse back to reactants without additional energy input.
Transition states cannot be isolated, have no measurable lifetime, and are purely theoretical constructs representing the highest-energy point along the reaction coordinate.
The free energy change of the reaction (ΔG_rxn) is the difference between the free energy of products and reactants:
Negative ΔG_rxn: exergonic, net energy released.
Positive ΔG_rxn: endergonic, net energy absorbed.
Concentration
Higher reactant concentration means more collisions per unit time.
In the Arrhenius framework, this effectively increases the frequency factor A.
Reaction rate increases with concentration for all reaction orders except zero order.
Temperature
Raising temperature increases the rate of nearly all reactions.
More molecules have kinetic energy equal to or exceeding Eₐ.
Every reaction has an optimal temperature. For human enzymes, this is around 38 to 40 °C; above that, proteins denature and activity drops sharply.
Reaction medium
The solvent or physical environment can affect rate.
Some reactions proceed faster in aqueous (water-based) environments; others prefer nonaqueous solvents.
Polar solvents tend to stabilise charged intermediates, create dipoles and increase rate.
The physical state of the medium (gas, liquid, solid interface) also matters.
A catalyst increases k by lowering Eₐ. It is not consumed in the reaction.
Catalysts work by stabilising intermediates, improving molecular orientation, increasing collision frequency, donating electron density, or reducing intramolecular bonding in reactant molecules.
Homogeneous catalysis: catalyst and reactants are in the same phase.
Heterogeneous catalysis: catalyst is in a different phase from reactants.
What catalysts do NOT do:
They do not change the free energy of reactants or products.
They do not shift the equilibrium position or change K_eq.
They speed up both the forward and reverse reactions by the same factor.
They cannot make a non-spontaneous reaction spontaneous.
Gibbs free energy relationship
ΔG = Eₐ(forward) − Eₐ(reverse)
Negative ΔG = exergonic
Positive ΔG = endergonic
Collision theory rate expression
Rate = Z × f
Z = collisions per second
f = fraction of collisions that are effective (correct orientation + sufficient energy)
Arrhenius equation
k = A × e^(−Eₐ / RT)
k = rate constant
A = frequency factor (pre-exponential factor)
Eₐ = activation energy (J/mol)
R = 8.314 J/(mol·K)
T = temperature in kelvin
Trends: increasing A increases k. Increasing T increases k (the exponent becomes less negative, pushing the exponential term closer to 1).
Catalytic converters in cars use heterogeneous catalysis (platinum, palladium, rhodium surfaces) to speed the breakdown of exhaust pollutants. The converter lowers Eₐ so toxic gases convert to less harmful products at manageable temperatures.
Enzymes in your body are biological catalysts operating under the same principles: they lower Eₐ for biochemical reactions that would otherwise take far too long to sustain life. Human enzymes are optimised for roughly 37 to 40 °C, which is why a fever can disrupt normal metabolism.
Students often confuse spontaneity with speed. A reaction can be thermodynamically spontaneous (negative ΔG) yet proceed imperceptibly slowly without a catalyst. Diamond turning into graphite is spontaneous, but you will not see it happen.
Students often think catalysts shift the equilibrium toward products. They do not. A catalyst speeds up both the forward and reverse reactions equally, leaving K_eq unchanged.
Students often assume stoichiometric coefficients from the balanced equation can be used directly as reaction orders in the rate law. They cannot, unless the reaction is a single elementary step.
Students sometimes believe raising temperature always helps. In enzyme-catalysed reactions, exceeding the optimal temperature denatures the enzyme and the rate drops sharply.
⚠️ You will almost certainly be asked to explain why a spontaneous reaction can still be slow. Have a clear one-sentence answer ready.
⚠️ The Arrhenius equation and its variable relationships (what happens to k when T increases, when Eₐ decreases) are high-frequency exam material. Know the direction of every trend.
⚠️ Catalyst questions often test whether you understand what catalysts cannot do: they cannot change ΔG, K_eq, or make a non-spontaneous reaction spontaneous.
⚠️ Distinguish between reaction intermediates (real species, produced and consumed within the mechanism) and transition states (theoretical, cannot be isolated). This distinction appears regularly in mechanism questions.
True or False: A catalyst changes the equilibrium constant of a reaction. (False. It changes k for both directions equally, leaving K_eq the same.)
Fill in the blank: In the Arrhenius equation, the two factors that determine k are the ______ and the exponential term containing activation energy and temperature. (frequency factor, A)
True or False: All molecular collisions between reactants lead to product formation. (False. Only collisions with correct orientation and sufficient energy are effective.)
Fill in the blank: The slowest step in a multi-step reaction mechanism is called the ______. (rate-determining step)
True or False: Increasing temperature decreases the rate constant k. (False. Increasing T makes the exponent less negative, increasing k.)
Q: A reaction has a negative ΔG but proceeds extremely slowly at room temperature. Explain how this is possible and name two ways to increase the rate.
A: A negative ΔG means the reaction is thermodynamically favourable, but the activation energy barrier may be very high. Most molecules at room temperature lack the kinetic energy to overcome it. Two ways to increase the rate: (1) raise the temperature so more molecules exceed Eₐ, or (2) add a catalyst to lower Eₐ.
Q: In the Arrhenius equation k = Ae^(−Eₐ/RT), what happens to the rate constant if the temperature is doubled (in kelvin)?
A: The exponent −Eₐ/RT becomes half as negative (since T is in the denominator). The exponential term increases, so k increases. The exact magnitude depends on the value of Eₐ, but k always rises when T rises.
Q: A student claims that adding a catalyst to an endergonic reaction will make it spontaneous. Is this correct?
A: No. A catalyst lowers the activation energy and increases the rate, but it does not change ΔG. The reaction remains endergonic (non-spontaneous under standard conditions) regardless of the catalyst.
Q: Distinguish between a reaction intermediate and a transition state.
A: A reaction intermediate is a real, short-lived species formed in one step and consumed in a later step of the mechanism; it sits in an energy valley between two transition states. A transition state is the highest-energy configuration along the reaction coordinate for a single step; it has no measurable lifetime and cannot be isolated.
Q: Explain why heterogeneous catalysis and homogeneous catalysis differ in practical terms, even though the underlying principle (lowering Eₐ) is the same.
A: In homogeneous catalysis the catalyst is in the same phase as reactants, which can make separation from products difficult. In heterogeneous catalysis the catalyst is in a different phase (typically a solid surface), making it easier to separate and reuse, which is why industrial processes often prefer it.
This material connects directly to chemical equilibrium: K_eq = k_forward / k_reverse, so understanding rate constants feeds straight into equilibrium calculations. It also ties into thermodynamics (Chapter 4), since ΔG determines spontaneity while Eₐ determines speed. If you move on to biochemistry, enzyme kinetics (Michaelis-Menten) builds on the zero-order and first-order rate behaviour covered here.
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