AP®︎ Chemistry: Unit 5 Flashcards
Master key terms and definitions for Unit 5 of AP Chemistry – Kinetics to help you prep for quizzes and the AP exam.
Chemical Kinetics
The study of how fast reactions occur and what factors affect their speed.
The study of how fast reactions occur and what factors affect their speed.
The change in reactant or product concentration per unit time.
Rate = -Δ[reactant]/Δt or Δ[product]/Δt, with the negative sign for reactant loss.
Usually mol L^-1 s^-1 or M/s, showing concentration change over time.
The change in concentration over a finite time interval between two measured points.
The rate at a specific moment, equal to the slope of the tangent line.
Find the slope of the curve; secant slope gives average rate, tangent slope gives instantaneous rate.
Balanced coefficients relate species rates, so concentration changes occur in fixed mole ratios.
Higher reactant concentration usually increases rate by causing more frequent effective collisions.
Higher temperature usually increases rate because particles have greater average kinetic energy.
Greater exposed solid area usually increases rate by allowing more collisions at the surface.
Higher gas pressure usually increases rate by increasing particle concentration in a given volume.
Reactants decrease while products increase, so their concentration changes have opposite signs.
Each exponent gives order in one reactant; their sum gives the overall order.
The proportionality constant in a rate law, whose value depends on temperature.
Units vary with overall order: zero M/s, first s^-1, second M^-1 s^-1.
Compare trials where one reactant concentration changes and others stay constant to find exponents.
Reaction orders must be found from measured rate data, not from coefficients in the balanced equation.
If concentration changes by a factor x, rate changes by x raised to that reactant's order.
Substitute one trial's rate and concentrations into the rate law, then isolate k.
Measure reactant disappearance or product formation over time to determine reaction rate.
An equation showing reaction rate as k times reactant concentrations raised to powers.
The exponent of a reactant concentration in the rate law.
Zeroth: [A]t - [A]0 = -kt; first: ln[A]t - ln[A]0 = -kt; second: 1/[A]t - 1/[A]0 = kt.
The time for concentration to halve, constant for first-order reactions, with t1/2 = 0.693/k.
Unstable nuclei decay at a rate proportional to the amount present, giving a constant half-life.
[A] vs t, ln[A] vs t, and 1/[A] vs t identify order and slope gives k.
Three-particle simultaneous collisions are uncommon because exact same-time, same-place encounters are unlikely.
For a single-step reaction, reactant coefficients become exponents in the rate law.
The number of reacting particles involved in an elementary step, such as unimolecular, bimolecular, or termolecular.
Particles react only when they collide with enough energy and the correct orientation.
Reacting particles must collide in an arrangement that allows the necessary bonds to rearrange.
Higher temperature increases average kinetic energy, causing more frequent and more energetic collisions.
Only a small fraction of collisions form products because few have both enough energy and correct orientation.
Product formation occurs only when collisions have enough energy and the correct orientation.
A curve showing particle energies; higher temperature increases the fraction exceeding activation energy.
Left side starting substances, peak unstable arrangement, and right side final substances.
Endothermic: products higher than reactants; exothermic: products lower than reactants.
Some bonds must be broken and new bonds formed during a single-step molecular event.
Higher temperature increases the fraction of collisions energetic enough to reach the peak.
An equation relating a reaction's rate constant to temperature and activation energy.
The minimum energy required for reactants to reach the transition state; on an energy profile, it is the energy difference from the reactants or current species to the next transition-state peak.
The x-axis on an energy profile showing reaction progress from reactants through intermediates to products as atoms rearrange.
The potential energy difference between products and reactants, often written as ΔE or ΔH, indicating whether energy is absorbed or released.
Intermediates are produced then consumed; catalysts are consumed then regenerated across the steps.
A mechanism may include reactants, intermediates, products, and catalysts across its steps.
A species formed in one step of a mechanism and consumed in a later step; it appears in the mechanism but cancels from the overall reaction.
A sequence of elementary steps that shows how reactants become products.
Add all elementary steps and cancel repeated species such as intermediates to obtain the net balanced reaction.
The slowest step in a mechanism that controls the overall reaction rate.
Use reactant coefficients in the slow step as exponents in the rate law.
Overall rate laws are determined experimentally and used to test whether a proposed mechanism is valid.
A catalyst is consumed then regenerated; an intermediate is formed then later consumed.
A method used when a fast reversible first step precedes a slow step.
Use the fast-step equilibrium expression to replace intermediates in the slow-step rate law.
A reaction that occurs through two or more elementary steps instead of one collision.
The y-axis quantity showing the relative energy of reactants, intermediates, transition states, and products.
Plot reactants, add peaks for each step, valleys for intermediates, then label each Ea and total ΔE.
The slowest step, identified by the largest activation energy barrier between adjacent minima and maxima.
A profile where products are lower in potential energy than reactants.
A profile where products are higher in potential energy than reactants.
A single mechanistic step that occurs in one molecular event, with its own reactants, products, and activation energy.
The highest-energy, unstable arrangement of atoms that reactants must reach at the top of the activation energy barrier before forming products.
A graph of potential energy versus reaction progress that shows reactants, products, transition states, intermediates, energy barriers, and overall energy change.
The catalyzed pathway uses different elementary steps and a lower activation energy than the original pathway.
They increase effective collisions and/or provide an alternate pathway with lower activation energy.
Temporary binding can orient reactants favorably or lower activation energy by forming a new intermediate.
A temporary species formed when a catalyst is attached to one or more reactants.
Some catalysts form temporary covalent bonds with reactants, creating alternate intermediates and reaction steps.
Biological catalysts that bind substrates, orient them, and lower activation energy for reaction.
It is consumed in one step and regenerated later, so its overall amount stays constant.
Catalysis by proton transfer that forms new intermediates and elementary steps with lower activation energy.
Catalysis on a solid surface where reactants form bound intermediates and react through new steps.
A substance that speeds a reaction by providing an alternative mechanism or pathway with lower activation energy and is regenerated overall rather than consumed.
