Chemical kinetics
56 flashcards covering Chemical kinetics for the MCAT Chem / Phys / Psych / Soc section.
Chemical kinetics is the study of how quickly chemical reactions occur and what influences their speed. It examines factors like temperature, concentration of reactants, and the presence of catalysts, helping us understand why some reactions happen in seconds while others take years. This branch of chemistry is essential for predicting and controlling reaction rates in real-world applications, from drug synthesis to industrial processes.
On the MCAT, chemical kinetics appears in the Chemical and Physical Foundations section, often through multiple-choice questions that test your ability to interpret rate laws, determine reaction orders, and analyze factors like activation energy. Common traps include confusing kinetics with equilibrium or misreading graphs of concentration versus time, so focus on mastering calculations involving rate equations and understanding reaction mechanisms. A solid grasp here can boost your score by applying these concepts to biological contexts, like enzyme activity.
Practice sketching reaction coordinate diagrams to visualize energy barriers.
Terms (56)
- 01
Rate of reaction
The rate of reaction is the change in concentration of a reactant or product per unit time, indicating how fast a chemical reaction occurs.
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Instantaneous rate
The instantaneous rate is the rate of reaction at a specific moment, determined from the slope of the tangent to the concentration-time curve at that point.
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Average rate
The average rate is the change in concentration of a reactant or product over a specific time interval, calculated by dividing the concentration change by the time elapsed.
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Rate law
The rate law expresses the relationship between the reaction rate and the concentrations of reactants, typically in the form rate = k [A]^m [B]^n, where k is the rate constant and m and n are the orders.
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Order of reaction
The order of reaction with respect to a reactant is the exponent of its concentration in the rate law, indicating how the rate depends on that reactant's concentration.
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Overall reaction order
The overall reaction order is the sum of the exponents in the rate law, determining the total dependence of the rate on all reactant concentrations.
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Zero-order reaction
A zero-order reaction has a rate that is independent of reactant concentrations, so the rate is constant and equal to the rate constant k.
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First-order reaction
A first-order reaction has a rate directly proportional to the concentration of one reactant, following the rate law rate = k [A].
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Second-order reaction
A second-order reaction has a rate proportional to the square of one reactant's concentration or the product of two reactants' concentrations, such as rate = k [A]^2 or rate = k [A][B].
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Method of initial rates
The method of initial rates determines the order of reaction by measuring the initial reaction rates at different initial concentrations and analyzing how the rate changes.
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Rate constant
The rate constant is the proportionality constant in the rate law that relates the reaction rate to reactant concentrations, with its value depending on temperature and the reaction.
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Units of rate constant
The units of the rate constant depend on the overall reaction order, such as mol/L/s for first-order, L/mol/s for second-order, and mol/L/s for zero-order reactions.
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Integrated rate law for first-order
The integrated rate law for a first-order reaction is ln([A]t / [A]0) = -kt, where [A]t is the concentration at time t, [A]0 is the initial concentration, and k is the rate constant.
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Integrated rate law for zero-order
The integrated rate law for a zero-order reaction is [A]t = [A]0 - kt, allowing calculation of concentration at time t from the initial concentration and rate constant.
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Integrated rate law for second-order
The integrated rate law for a second-order reaction is 1/[A]t = 1/[A]0 + kt, used to find concentration over time for reactions like rate = k [A]^2.
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Half-life
Half-life is the time required for the concentration of a reactant to decrease to half its initial value, varying by reaction order.
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Half-life for first-order reactions
For first-order reactions, half-life is constant and given by t1/2 = ln(2)/k, meaning it does not depend on initial concentration.
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Half-life for zero-order reactions
For zero-order reactions, half-life is t1/2 = [A]0 / (2k), so it depends on the initial concentration of the reactant.
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Half-life for second-order reactions
For second-order reactions, half-life is t1/2 = 1 / (k [A]0), increasing as the initial concentration decreases.
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Activation energy
Activation energy is the minimum energy required for reactant molecules to collide effectively and form products, determining how temperature affects reaction rate.
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Arrhenius equation
The Arrhenius equation, k = A e^(-Ea/RT), relates the rate constant k to temperature, where A is the pre-exponential factor, Ea is activation energy, R is the gas constant, and T is temperature in Kelvin.
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Pre-exponential factor
The pre-exponential factor in the Arrhenius equation represents the frequency of collisions and the orientation factor, indicating how often molecules collide with proper orientation.
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Collision theory
Collision theory states that for a reaction to occur, molecules must collide with sufficient energy and proper orientation, explaining the dependence of rate on concentration and temperature.
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Effective collisions
Effective collisions are those that result in product formation, requiring molecules to have energy at least equal to the activation energy and the correct orientation.
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Orientation factor
The orientation factor is the fraction of collisions with the proper geometry for reaction, a component of the pre-exponential factor in the Arrhenius equation.
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Transition state
The transition state is the high-energy arrangement of atoms at the peak of the reaction coordinate, representing the point of no return in a chemical reaction.
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Reaction mechanism
A reaction mechanism is the step-by-step sequence of elementary reactions that describe how overall reactants transform into products.
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Elementary step
An elementary step is a single reaction event in a mechanism, where the molecularity equals the number of molecules involved, and its rate law can be written directly from the stoichiometry.
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Rate-determining step
The rate-determining step is the slowest step in a reaction mechanism, controlling the overall reaction rate and determining the form of the rate law.
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Molecularity
Molecularity is the number of molecules or ions that participate in an elementary reaction step, such as unimolecular or bimolecular.
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Catalyst
A catalyst is a substance that increases the reaction rate by providing an alternative pathway with lower activation energy, without being consumed in the process.
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Homogeneous catalyst
A homogeneous catalyst is in the same phase as the reactants, such as a dissolved acid speeding up a reaction in solution.
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Heterogeneous catalyst
A heterogeneous catalyst is in a different phase from the reactants, like a solid metal surface facilitating gas-phase reactions.
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Enzyme as catalyst
Enzymes are biological catalysts that speed up reactions in living organisms by lowering activation energy, often through specific active site binding.
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Inhibitor
An inhibitor is a substance that decreases the reaction rate by interfering with the catalyst or reactants, such as by blocking the active site of an enzyme.
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Temperature effect on rate
Increasing temperature generally increases reaction rate by providing more kinetic energy to molecules, leading to more frequent and energetic collisions.
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Maxwell-Boltzmann distribution
The Maxwell-Boltzmann distribution describes the distribution of molecular speeds in a gas, explaining how temperature affects the fraction of molecules with sufficient energy to react.
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Pseudo-first-order reaction
A pseudo-first-order reaction occurs when one reactant's concentration is held constant, making the rate law appear first-order overall.
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Reversible reactions
Reversible reactions can proceed in both forward and reverse directions, and their rates depend on the concentrations of both reactants and products.
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Effect of concentration on rate
Higher reactant concentrations typically increase the reaction rate by increasing the frequency of collisions between molecules.
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Strategy for determining order graphically
To determine reaction order graphically, plot concentration versus time for zero-order, ln(concentration) versus time for first-order, or 1/concentration versus time for second-order, and identify the straight line.
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Radioactive decay as first-order
Radioactive decay follows first-order kinetics, where the rate is proportional to the amount of radioactive substance present, with a constant half-life.
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Beer-Lambert law in kinetics
The Beer-Lambert law relates absorbance to concentration, allowing monitoring of reaction rates spectrophotometrically by measuring changes in absorbance over time.
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Differential rate laws
Differential rate laws express the instantaneous rate of change of concentration with respect to time, such as d[A]/dt = -k [A] for first-order reactions.
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Gibbs free energy of activation
The Gibbs free energy of activation is the energy barrier for the reaction, related to both enthalpy and entropy changes in the transition state.
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Enthalpy of activation
The enthalpy of activation is the energy difference between the transition state and reactants, contributing to the overall activation energy.
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Entropy of activation
The entropy of activation reflects the change in disorder from reactants to the transition state, affecting the pre-exponential factor in reaction rates.
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Michaelis-Menten kinetics
Michaelis-Menten kinetics describes enzyme-catalyzed reactions, where the rate depends on substrate concentration and reaches a maximum at high substrate levels.
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Lineweaver-Burk plot
The Lineweaver-Burk plot is a double-reciprocal graph of 1/rate versus 1/substrate concentration, used to determine kinetic parameters like Km and Vmax for enzyme reactions.
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Autocatalysis
Autocatalysis occurs when a product of the reaction acts as a catalyst, leading to an initial slow rate that accelerates as more product forms.
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Induction period
The induction period is the initial time in a reaction where the rate is slow due to the buildup of intermediates before the main reaction accelerates.
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Chain reactions
Chain reactions involve a series of steps where reactive intermediates propagate the reaction, common in free radical processes like combustion.
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Free radical mechanisms
Free radical mechanisms involve highly reactive species with unpaired electrons, initiating and propagating chain reactions in processes like halogenation.
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Example: Calculating rate constant from data
To calculate the rate constant, use initial rates and concentrations in the rate law; for a first-order reaction, plot ln([A]) vs. time and find the slope as -k.
For a reaction with [A] decreasing from 1.0 M to 0.5 M in 100 seconds, k = ln(1.0/0.5)/100 = 0.00693 s^-1.
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Example: Using Arrhenius equation
The Arrhenius equation calculates how k changes with temperature; if two k values at different T are known, solve for Ea.
If k1 at T1 and k2 at T2 are given, Ea = R ln(k2/k1) / (1/T1 - 1/T2).
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Strategy for identifying rate-determining step
To identify the rate-determining step, compare the rates of individual steps in a mechanism and assume the slowest one dictates the overall rate law.