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Chemistry Chapter 13
Chemical Kinetics
| Term | Definition |
|---|---|
| Photochemical Smog | Forms when solar radiation interacts with the gasses formed within the internal combustion chambers of cars/trucks. (13.1) |
| Free Radicals | Atoms or molecules with an odd number of electrons, very volatile and reactive. Typically a danger to the environment. (13.1) (8.6) |
| Rate of Reaction | The rates at which reactants are consumed and products are formed. Concentration acts as a function of time. (13.1) |
| Chemical Kinetics | The study of the rates of chemical reactions and the factors that influence those rates. Chemists use this knowledge to determine how reactions happen at the molecular level—that is, the mechanisms of reactions. With this knowledge, chemists can manipulate reaction conditions to increase reaction rates and, ultimately, product yields. (13.1) |
| Relationship between Kinetic Energy and Collision Frequency | Faster speeds mean more frequent and forceful collisions between molecules. (13.2) |
| Four Factors Influence Reaction Rates | 1) Physical States of Reactants. 2) Concentration of Reactants. 3) Temperature. 4) Catalysts. (13.2) |
| Physical States of Reactants | Differences in particle motion and location are one reason why reactions in the solid phase tend to be much slower than reactions in liquids and gases. (13.2) |
| Concentration | Greater concentrations often produce more rapid reaction rates. (13.2) |
| Temperature | As temperature increases, the rates of chemical reactions tend to increase. The average kinetic energy (KE) of particles increases as temperature rises, and as their KE increases, the number of collisions between the particles increases, as does the probability that those collisions will lead to reactions between them. (13.2) |
| Catalysts | Accelerate a reaction without they themselves being consumed in the process. Typically provide a surface for reactants to bind to, creating a complex that lowers activation energy. (13.2) |
| Concentration Notation | Brackets around a [reactant] or [product] indicate the amount of substance in moles per liter. Useful in equations. (13.2) |
| Coefficient Pattern | The coefficient of each species in the balanced chemical equation appears in the denominator of its term in the relative rate expression if the numerator values are all equal to one. (13.2) |
| Reaction Rate Units | Rates of product formation or reactant consumption are ratios of changes in concentration divided by changes in time, so they have units of concentration per unit time, such as molarity per second (M/s). Can only be determined experimentally. Once the rate of a reaction is known, it can be used along with the coefficients in the balanced equation describing the reaction to calculate the rate of change in the concentration of any reactant or product. (13.2) |
| Expressing Rate of Change when no Coefficients Equal One | If none of the coefficients is 1, divide the rate of change in the concentration of any of the products by its coefficient to get the reaction rate. (13.2) |
| Instantaneous Reaction Rate | Use calculus to solve for instantaneous reaction rates. The instant ROC is equal to the slope of a tangent line at any point on the graph. Remember to keep in mind reactants are being consumed, meaning their slope will be negative. (13.2) |
| Calculating Initial Reaction Rate | The slope of the tangent line immediately as the reaction starts will be equal to the instantaneous initial reaction rate. (13.3) |
| As Reaction Ends | When the slope of the tangent to the curve reaches zero, there are no more changes in the concentrations of the product(s) or any remaining reactant(s). (13.3) |
| Reaction Order | How reaction rate depends on reactant concentrations. (13.3) |
| Second Order with respect to Reactant | The rate of reaction is proportional to the concentration of A squared. (r = k[A]²) In other words, when [A] is doubled, the rate will quadruple. (13.3) |
| First Order with respect to Reactant | The rate of reaction is proportional to the concentration of A. (r = k[A]¹) In other words, when [A] is doubled, the rate will double as well. (13.3) |
| Zero Order with respect to Reactant | The rate of reaction is not proportional to the concentration of the reactant. (r = k[A]⁰) If [A] changes, the rate will be unaffected. (13.3) |
| Overall Order of Reaction | Equal to the sum of (m + n + p) if (r = k[A]ᵐ[B]ⁿ[C]ᵖ). In other words, the sum of all orders with respect to individual reactants. (13.3) |
| Rate Law | An equation that relates reaction rate to reactant concentrations. (ex. r = k[A]ᵐ[B]ⁿ). (13.3) |
| Using Two Rates and Two Concentrations to Find Order of One Reactant | n(or m) = (log(rate₁/rate₂))÷(log([B]₁/[B]₂)). (13.3) |
| Proportionality Constant, k | Reaction rate depends on the concentration of the reactants, but the rate constant does not. The rate constant changes only with changing temperature or in the presence of a catalyst. (13.3) |
| Solving for k | Insert any set of data into a rate law equation to solve for k. For example (rate) = k(conc. A)ᵐ(conc. B)ⁿ. This data can be found on any of the experimental charts used to calculate rate law. Use algebra to solve and treat order exponents as regular exponents. (13.3) |
| Zero Order Units of k | k = mol/L ⋅ s or M/s or M ⋅ s⁻¹ (isolate k and then simplify units on right side). (13.3) |
| First Order Units of k | k = s⁻¹ (isolate k and then simplify units on right side). (13.3) |
| Second Order Units of k | k = M⁻¹ ⋅ s⁻¹ (isolate k and then simplify units on right side). (13.3) |
| Integrated Rate Law | Integral Calculus is used to derive these equations. (13.3) |
| First Order Integrated Rate Law | ln([X]÷[X]₀) = -kt can be rearranged to ln[X] = -kt + ln[X]₀ which is a linear form of the equation. If this produces a linear equation, it means that the reaction is first order with respect to [X]. (13.3) |
| Half Life (t_½) | The interval during which the concentration of a reactant decreases by half. Half-life is inversely related to the rate constant of a reaction: the higher the reaction rate, the shorter the half-life. The half-life of a first-order reaction is constant throughout the reaction and independent of concentration: no matter the initial concentration of the reactant, half of it is consumed in one half-life. (13.3) |
| Natural Logarithm Relationship | ln(a/b) = ln(a) - ln(b) |
| Calculating Half Life | ln(1/2) = -kt_½. (13.3) |
| Second Order Integrated Rate Law | 1/[X] = kt + 1/[X]₀. If the equation is linear, the reaction is second order with respect to X. This applies to any second order reaction with one reactant. (13.3) |
| Zero Order Integrated Rate Law | [X] = -kt + [X]₀. If the equation is linear, the reaction is zero order with respect to X. (13.3) |
| Activation Energy | The minimum amount of energy that enables reactions to happen is called the activation energy (Eₐ). (13.4) |
| Arrhenius Equation | The mathematical connection between temperature, the rate constant k for a reaction, and its activation energy is given by the Arrhenius equation. (k = Ae^(-Ea/RT)) (13.4) |
| Frequency Factor | Represented by A. The product of the collision frequency and a term that accounts for the fact that not every collision results in a chemical reaction. Some collisions do not create products because the colliding molecules are not oriented properly. (13.4) |
| Activated Complex | A transient species with both proper orientation and enough energy to react with one another. Activated complexes have extremely brief lifetimes and fall apart rapidly, either forming products or reforming reactants. Activated complexes are formed by reacting species that have acquired enough energy to react with each other. (13.4) |
| Transition State | The internal energy of an activated complex represents a high-energy transition state of the reaction. Represented by the top of curves on potential energy diagrams. (13.4) |
| Reaction Mechanism | The stepwise manner in which the bonds in reactant molecules break and the bonds in product molecules form. A combination of elementary steps. (13.4) |
| Intermediate | Produced in one step and consumed in the next. Intermediates are not considered reactants or products and do not appear in the equation describing a reaction. (13.4) |
| Unimolecular | An elementary step that involves a single molecule. (13.4) |
| Bimolecular | A collision between two molecules. (13.4) |
| Termolecular | Three-molecule elementary steps because the chance of three molecules colliding at exactly the same time in the proper orientation is much smaller. (13.4) |
| Multistep Reaction Mechanisms | The slower step determines rate. When the slower step is second, find the reaction mechanism's law using the following two rules: 1. The equations for the elementary steps in the mechanism must add up to the overall equation for the reaction. 2. The mechanism must be consistent with the experimental rate law. (13.4) |
| Forward RXT Activation Energy | The forward reaction sometimes has a lower activation energy than the reverse reaction. (13.4) |
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