Kinetics

This is Unit Nine of Chemistry, Part Two: collision theory, mechanisms and catalysis. Part One measured how fast reactions go and established that a rate law's exponents come from experiment rather than from a balanced equation. This part explains why — and the explanation turns out to be that the balanced equation is not what actually happens. Real reactions proceed through a sequence of individual collision events, and the equation is only their sum. By the end you will be able to say why most collisions do not produce a reaction; tell a transition state from a reaction intermediate; say which step sets a reaction's overall rate; say when a coefficient can legitimately be used as a rate-law exponent; and say what a catalyst does not do.

Four entries. The idea that a reaction rate is really a collision rate, the two separate conditions a collision has to satisfy before anything happens, and the fleeting species that exists at the moment it does.

Rate from collisions.

Chem 1683 Collision Theory's Rate-Collision Proportionality

Collision theory holds that a reaction's rate is proportional to the rate at which its reactant particles collide. What does not follow is that every collision reacts — if it did, rates would equal collision frequency, and they are far lower. Most collisions fail, from bad orientation or insufficient energy, and the next two entries take those in turn.

The right way round.

Chem 1686 Collision Orientation Requirement

A collision reacts only if the colliding particles meet in an orientation that brings the atoms which will bond into contact. A collision between the right particles at the right concentration still fails when their geometry is wrong at the moment of impact — the molecules glance off unchanged.

Enough energy.

Chem 1689 Activation Energy

Activation energy is the minimum energy a properly oriented collision must supply for reactants to form products. It is unrelated to whether a reaction is overall exothermic or endothermic — a reaction that releases energy overall still needs energy to get started, which is why fuels do not ignite on their own.

At the peak.

Chem 1692 The Transition State

A transition state, or activated complex, is the short-lived, high-energy species existing at a collision's peak, which may go on to products or fall back to reactants. It is not an isolable intermediate — it exists only at the instant of peak energy, which is what distinguishes it from the intermediates this part covers later.

Two entries. Part One established that raising temperature raises rate but that the per-degree response varies by reaction. This is the equation that makes that precise, and the procedure for getting an activation energy out of two measurements.

Rate constant and temperature.

Chem 1695 The Arrhenius Equation

The Arrhenius equation relates a rate constant to activation energy and absolute temperature — k equals A times e to the minus activation energy over R T, with R the gas constant. Temperature must enter in kelvin, the same requirement already flagged for the ideal gas law and for Clausius-Clapeyron.

Two measurements.

Chem 1698 Determining Activation Energy From Rate Constants at Two Temperatures

Determining an activation energy from rate constants measured at two temperatures means substituting both temperature and rate-constant pairs into the two-point form of the Arrhenius equation, then solving for the activation energy term. Using Celsius for either temperature, or swapping which pair is point one, produces an activation energy far from the correct magnitude or of the wrong sign entirely.

Four entries. What a balanced equation conceals, the number of particles involved in a single real step, why three-particle steps are rare, and a species that exists partway through and never appears in the equation at all.

What actually happens.

Chem 1703 Elementary Reactions and Reaction Mechanisms

A reaction mechanism is the sequence of elementary reactions — each occurring exactly as written — that sum to the overall equation. A mechanism's steps are literal bond-breaking and bond-forming events, unlike an overall equation, which shows only the net substances going in and coming out.

How many particles.

Chem 1706 Molecularity: Unimolecular and Bimolecular Steps

An elementary step's molecularity is the number of particles involved in it: a unimolecular step involves one, a bimolecular step two. Molecularity counts particles in one elementary step — it describes an overall reaction only in the special case where that reaction happens in a single step.

Three at once.

Chem 1709 Termolecular Elementary Reactions

A termolecular elementary reaction requires three particles to collide simultaneously — a rare event, far less probable than an ordinary two-particle collision. Seeing three reactant particles in an overall equation does not mean the mechanism contains a termolecular step; a sequence of bimolecular steps is much likelier.

Appears, then goes.

Chem 1712 Reaction Intermediates

A reaction intermediate is produced in one mechanism step and fully consumed in a later one, so it never appears in the overall equation. Unlike a transition state, an intermediate is a real if short-lived species that can be detected between steps — which is the distinction between the two easiest to lose.

Three entries, and this is where Part One's central correction gets its explanation. Which step governs the overall rate, the one situation in which coefficients legitimately become exponents, and the derivation that combines both.

The bottleneck.

Chem 1715 The Rate-Determining Step

A mechanism's rate-determining step is its slowest elementary step, and the overall reaction proceeds no faster than that step allows. Assuming the first written step is always rate-determining is wrong — any step in the sequence can be the slowest, regardless of where it appears.

Here it does work.

Chem 1718 An Elementary Step's Rate Law Follows From Its Molecularity

Unlike a reaction's overall rate law, an elementary step's rate law can be written straight from its equation, because in a single collision event the coefficients genuinely count the colliding particles. This works for a single elementary step only — applying it to a multistep overall equation is exactly the error Part One named.

Put it together.

Chem 1721 Deriving a Rate Law from a Multistep Mechanism

Deriving a rate law from a multistep mechanism means writing the rate law for the rate-determining step, then, if that law contains an intermediate's concentration, using an earlier fast reversible step at equilibrium to re-express the intermediate in terms of the original reactants. Leaving an intermediate in the final rate law is an incomplete derivation — a correct rate law names only species appearing in the overall balanced equation.

Six entries closing the unit. What a catalyst does, the thing it is most often wrongly believed to do, how it actually achieves its effect, the two kinds, and then the capstone joining rate law, mechanism and catalysis on one reaction.

A lower path.

Chem 1726 Catalysts Lower Activation Energy Without Being Consumed

A catalyst speeds a reaction by providing a pathway with lower activation energy, and though it takes part in the mechanism it is regenerated unchanged by the end. A catalyst appearing inside a mechanism's steps is easy to mistake for a reactant — the test is whether it reappears unchanged once the reaction completes.

What it does not do.

Chem 1729 The "Catalysts Shift Equilibrium" Misreading

A catalyst speeds how fast a reaction reaches equilibrium, but does not shift where that equilibrium lies, because it lowers the forward and reverse activation energies equally. "Catalysts speed up reactions" is easy to over-extend into "catalysts favour products" — only the speed changes, not the position, which Unit Ten builds in full.

A different route.

Chem 1732 A Catalyst's Alternative, Lower-Energy Mechanism

A catalysed reaction proceeds through a different mechanism from its uncatalysed form, one whose rate-determining step has a lower activation energy. A catalysed mechanism is not merely the same pathway running faster — it is frequently an entirely different route to the same products.

Same phase.

Chem 1735 Homogeneous Catalysis

A homogeneous catalyst exists in the same phase as the reactants, forming an intermediate that later decomposes to regenerate the catalyst. It must still be regenerated by the end — a species appearing once with nothing reproducing it is a reactant, whatever else it does.

Different phase.

Chem 1738 Heterogeneous Catalysis

A heterogeneous catalyst exists in a different phase from the reactants, usually a solid, furnishing a surface where reactants adsorb, react, and then desorb. Its action happens at that surface only — so surface area rather than total mass is what determines how much it speeds the reaction, the same distinction Part One drew for heterogeneous reactions generally.

One reaction, three views.

Chem 1740 Capstone: From Experimental Rate Law to Catalyzed Mechanism

The capstone carries one reaction from its experimentally determined rate law, through collision theory's activation-energy picture, to a proposed mechanism consistent with that rate law, and then to the catalysed pathway that lowers the barrier. The point is that the three are not separate topics: a rate law is evidence about a mechanism, a mechanism explains why the rate law has the form it does, and a catalyst changes the rate law by changing the mechanism.

Five things to carry out of this part, and out of the unit. Most collisions do not react, which is why a reaction's rate is far below its collision frequency: a collision must bring the right atoms into contact — the orientation requirement — and it must supply at least the activation energy, and failing either leaves the molecules to glance off unchanged. A transition state is the high-energy species at a collision's peak and cannot be isolated, existing only at that instant; a reaction intermediate is produced in one step and consumed in a later one, is a real if short-lived species, and can be detected between steps. A reaction's overall rate is set by its rate-determining step, the slowest elementary step in the mechanism — and that is not necessarily the first one written. A coefficient can be used as a rate-law exponent in exactly one case: an elementary step, where the coefficients genuinely count the particles colliding. Applying it to an overall multistep equation is the error Part One named, and this part explains why — the overall equation is only a sum of steps and does not describe any actual collision. And what a catalyst does not do is shift where equilibrium lies: it lowers the forward and reverse activation energies equally, so it changes only how quickly equilibrium is reached. Unit Ten takes up equilibrium itself.