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Unit 5
7–9% of examKinetics is about how fast reactions happen and why. You'll use experimental data to write rate laws, read concentration-versus-time graphs, explain rates with collisions and activation energy, and connect step-by-step mechanisms to the rate law you measure.
Longer videos that cover the whole unit. Good for a first pass or a final review.
Reaction rate is how fast a reactant's concentration falls or a product's rises over time, and the coefficients link the rates of different species: in 2NO₂ → 2NO + O₂, NO₂ is used up twice as fast as O₂ forms. Rates depend on concentration, temperature, surface area and catalysts.
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A rate law has the form rate = k[A]ᵐ[B]ⁿ, where the orders m and n are found by experiment, often by comparing initial rates while changing one concentration at a time. If doubling [A] doubles the rate, the reaction is first order in A; if it quadruples the rate, it's second order. The rate constant k changes with temperature, and its units depend on the overall order.
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Graphs tell you the order: [A] versus time is a straight line for zero order, ln[A] versus time for first order, and 1/[A] versus time for second order. The slope gives k: it equals −k for zero and first order and +k for second order. First-order reactions, like radioactive decay, have a constant half-life: t½ = 0.693/k.
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An elementary reaction happens in a single collision step, so its rate law comes straight from its coefficients: for A + B → products, rate = k[A][B]. Steps that need three particles to collide at the same moment are rare.
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Particles react only if they collide with enough energy to get over the activation energy and with the right orientation, so most collisions don't make products. A Maxwell-Boltzmann distribution shows that at a higher temperature a larger fraction of particles have enough energy, which is why reactions speed up.
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An energy profile plots energy along the reaction coordinate, from reactants up to the transition state and down to products. The rise from reactants to the peak is the activation energy, and the gap between reactants and products is the overall energy change. The Arrhenius equation connects rate, temperature and activation energy, but you won't be asked to calculate with it.
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A mechanism is the series of elementary steps that add up to the overall reaction. An intermediate is made in one step and used up in a later one, so it doesn't appear in the overall equation; a catalyst is used up in one step and made again in a later one.
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The slowest step, called the rate-determining or rate-limiting step, sets the overall rate. When the first step is the slow one, the rate law for the whole reaction is simply the rate law of that first step.
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When a fast, reversible step comes before the slow step, the slow step's rate law may include an intermediate. You assume the fast step stays at equilibrium (its forward and reverse rates are equal) and use that to rewrite the intermediate's concentration in terms of reactants.
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A multistep reaction's energy profile has one hump for each elementary step, with intermediates sitting in the dips between them. The step with the largest activation energy is usually the slowest, rate-determining step.
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A catalyst speeds up a reaction by providing a different pathway with a lower activation energy, and because it's regenerated it isn't used up overall; it doesn't change the reaction's overall ΔH. Examples include enzymes that bind reactants, acid-base catalysis (a proton is gained or lost) and surface catalysis on a solid such as a metal.
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