Kinetics

**(Chapter 14, pages: 579-616**)
 * Kinetics **

1. Understand the following terms/concepts: a. Chemical Kinetics Chemical Kinetics is the study of how fast reactions take place. It also describes the reaction mechanism, or exactly "how" the reaction takes place, by identifying any intermediary species and states in the reaction (or things that happen in between).

**b. Reaction Rate-** How fast or slow the reaction takes place. Ex. aA + bB → pP + qQ (lower case letters (a,b,p,q) = coefficients) (t = time) rate = (-1/a)( ΔA/ Δt) = (-1/b)( ΔB/ Δt) = (-1/p)( ΔP/ Δt) = (-1/q)( ΔQ/ Δt) The average rate over shorter and shorter time intervals; the rate for a specific instant in time. Equal to the slope of a tangent to the curve at any particular time. (The time it takes for the reaction is negligible.)
 * c. Instantaneous rate- **

The instantaneous rate at the beginning of the reaction. This data is used to determine the rate law of a reaction by comparing the initial rates at various concentrations of reactants.
 * d. Initial rate- **

The average rate which a reaction is undergoing its reaction process. Calculated by taking the Final value (Molarity in this case) minus the initial value, and divide it by the time elapsed.
 * e. Average rate- **

This is what is used in real world experiments because it is much easier to determine and calculate if we tried to find the instantaneous rate in a lab we'd have to know the weights of all the substances at the instants and the times would need to be very precise. That's why we use the average rate and just wait until the reaction is complete to find the rate

Ex: The reaction rate is directly proportional to the number of molecular collisions per second. describes how the rate o f a reaction is related to the number of properly-oriented collisions of the molecules involved. It must have the right orientation and enough energy so that it will be enough to overcome the activation energy and form a product, as well as to avoid repulsion.
 * f. Collision theory- **



A theory that explains how the molecular nature of gases gives rise to their macroscopic properties. It also tells us that the thermal/kinetic motion of particles can explain the behavior of a gas. The minimum amount of energy to begin a chemical reaction. The minimum energy of collision that reactants must need to have a successful activated complex, which is the most energetic point of the reaction.
 * g. Kinetic molecular theory- **
 * Increasing kinetic energy will create more collisions
 * The pressure of a gaseous system is related to the number of collisions of the molecules with the sides of the container at a specific time
 * h. Activation energy- **

something that speeds up a reaction by lowering its required activation energy. A catalyst is not consumed during the course of a reaction, I may however form an intermediate during the reaction, but it is regenerated in a later step in the reaction. "A substance that initiates or accelerates a chemical reaction without itself being affected"
 * i. Catalyst- **

The speed at which two things mix and react with each other The change in concentration of reactants or products per unit of time. Rates can be figured as average rates of change, or for a more accurate rates, the instanateous rate of change may be calucated.
 * j. Reaction rate- **

k. **Rate law** The reaction that relates the rate of the reaction to the initial concentration of the reactants. So, the rate of the reaction is equal to the product of the rate constant and the initial concentration of the reactants raised to a superscript that is experimentally derived. Zero Order: [A] t = -kt + [A] o Half life: t 1/2 = [A] o / 2k First Order: ln [A] t = -kt + ln[A] o Half life: t 1/2 = 0.693/k Second Order: 1/[A] t = kt + 1/[A] o Half life: t 1/2 = 1/k [A] o

the order of reaction with respect to a certain reactant is the power to which its concentration in the rate equation is raised to. This is a measure of how much the reaction rate will change in relation to the concentration of a particular reactant.
 * l. Order of reaction- **

m. **Transition state theory** Transition State Theory attempts to explain what happens to molecules after they collide. It is the state of the molecules before they either (a) collide with the correct orientation, and enough kinetic energy is produced from the collision to overcome the needed minimum activation energy to form the products (b) collide with the correct orientation, but are not able to form products because the energy needed to overcome the minimum activation energy was not satisfied.

**a. Physical state**- ﻿the molecules must come in contact in order for the reaction to occur. b. **Temperature**- nearly all reactions happen faster at higher temperatures. This occurs because there are more collisions between particles in the reaction. Nearly all reactions will slow down when the temperature decreases. This is because there are less collisions between particles. c. **Pressure** - An increase in pressure of the system means either (a) the concentration per volume ( volume is constant) of the system increased or (b) the volume of the system decreased. Either way, an increase in pressure translates into molecules colliding more frequently and so, there is an increased probability of a reaction occurring. If the pressure decreases, the molecules collide less often, and so, the reaction is less likely to occur. d. **Molar concentration-** ﻿An increase in molar concentration increases the chances of the reactants to collide. e. **Activation energy-** ﻿energy needed for the reaction to start. f. **Catalyst**- speeds up a reaction by lowering the activation energy.
 * 2. Factors that affect the reaction rate: **
 * The more homogeneous the mixture of the reaction, the faster the molecule can react.
 * An increase in temperature provides molecules with enough energy to overcome the activation energy

a. **Method of initial rates -** The instantaneous rate of reaction when t=0. This method measures and compares the inital rates to determine the rate law. 1st assume that the rate is based on what reactants are in the equation. Then n and m are the orders of the reactants. Then divide the 2nd rate divided by the 3rd rate. We would then get the values in a table of the 1st and 2nd experiment and add them together. Then simplify the equation. b**. Method of graphical analysis** By graphing time vs: (A), (ln(A)), or (1/A), as three seperate graphs, the one that forms a straight line is the graph which shows you the overall rate of reaction. Defined as a series of steps by which a chemical reaction occurs. The slow step is the rate determining step. These steps are descriptive of how compounds act as they proceed from reactants to product.
 * 3. Be able to calculate the rate law of the reaction by using: **
 * c. Reaction mechanisms **


 * C4H8O2 (aq) + 2 NaBH4 (g) ßà C4H10O2 (aq) + 2 NaBH3 (aq) **
 * Trail || [C4H8O2], M || [NaBH4], M || Rate, M-1s-1 ||
 * 1 || 1.033 || 1.033 || 0.0269 ||
 * 2 || 2.060 || 3.100 || 0.2421 ||
 * 3 || 1.033 || 5.200 || 0.6818 ||
 * <span style="display: block; font-family: Cambria; font-size: 11pt; margin: 0in 0in 0pt; text-align: center;">4 || <span style="display: block; font-family: Cambria; font-size: 11pt; margin: 0in 0in 0pt; text-align: center;">0.560 || <span style="display: block; font-family: Cambria; font-size: 11pt; margin: 0in 0in 0pt; text-align: center;">1.033 || <span style="display: block; font-family: Cambria; font-size: 11pt; margin: 0in 0in 0pt; text-align: center;">0.0271 ||

1. Determine the rate law for the reaction using the data presented above. (5 pts)

rate3/rate1: 0.6818/0.02609= k[1.033]^n[5.200]^m/k[1.033]^n[1.033]^m k's and [A]^n cancel out so you get 25=5^m then m=2 then you take rate4/rate1: 0.0271/0.0296=k[0.560]^n[1.033]^m/k[1.033]^n[1.033]^m k's and [B]^m cancel out so you get 1=0.5^n then n=0 using rate 1 0.0269=k[1.033]^0[1.033]^2 0.0269=k[1][1.067] 0.0269/1.067=k so k=0.0252 rate law is rate=0.0252[A]^0[B]^2

3. Calculate the rate constant value k of the reaction and express the rate law (5 pts)

3. Using your results from the previous question, identify which of the following is a possible mechanism if any for the reaction. (5 pts) // Assume: C4H8O2 = A; NaBH4 = B; C4H10O2 = C; NaBH3 = D // // Hintz: Express the calculated rate law in terms of A and B //

Mechanism 1: A ßà D (fast) B + B à D (slow) 2D à C (fast)

Mechanism 2: This is the possible mechanism for the reaction. Each elementary mechanism follows the rules, and provides a second order reaction.

B + A ß à C + E (fast) E + B ßà 2D (slow)

Mechanism 2: A + 2B à C + 2D

Mechanism 4: A <-> B + E + D (fast) 2B à C (slow) B + E à D (fast)

1. Determine the rate law for the reaction using the following experimental data:

R1/R2 = k[.24]^x[.48]^y[.36]^z / k[.122]^x[.48]^y[.36]^z = .1003/.0229 2^x = 4 x = 2
 * ** Trail ** || ** [A], M ** || ** [B], M ** || ** [C], M ** || ** Rate, M-2s ** ||
 * ** 1 ** || 0.240 || 0.480 || 0.360 || 0.1003 ||
 * ** 2 ** || 0.122 || 0.480 || 0.360 || 0.0229 ||
 * ** 3 ** || 0.240 || 0.960 || 0.090 || 0.1986 ||
 * ** 4 ** || 0.122 || 0.480 || 0.180 || 0.0246 ||
 * ** 5 ** || 0.122 || 0.240 || 0.180 || 0.0119 ||

R2/R4 = k[.122]^x[.48]^y[.36]^z / k[.122]^x[.48]^y[.18]^z = .0229 / .0246 2^z = 1 z = 0

R4/R5 = k[.122]^x[.48]^y[.18]^z / k[.122]^x[.24]^y[.18]^z = .0246 / .0119 2^y = 2 y = 1

Rate law: R = k[A]^2[B]^1[C]^0
 * R=k[A]^2[B]**
 * k=3.62**
 * R=(3.62)[A]^2[B]**

2. Identify which of the following could be a possible mechanism for the reaction, using the rate law calculated on question 1. Mechanism 1: C + B ßà F + G (fast) 2F + B à D (fast) G + A à F + E (slow)

Mechanism 2: B + B ßà H (fast) H + A à G + E (slow) G + C à D (fast)

Mechanism 3 A + 2B à C + G (slow) G + 2C à D (fast)

Mechanism 4 A + B ß à H + 2D (fast) H + B + C à K (slow) K + D à E (fast)

a**. Elementary steps** - ﻿Every step in a reaction mechanism represents an elementary reaction, it occurs in a single collision of the reactant molecules. Elementary reactions may be unimolecular, bimolecular, or termolecular (one, two or three reactants) The rate of the product formation depends on how frequently the reactants collide, which in turn depends on the concentrations of the reactants. The slowest of the Elementary steps is the determining step. In this step, for a correct mechanism, the Coefficients on the reactants should match the powers for the reactants in the reaction's rate law. b. **Rate law** - ﻿Rate laws are determined using a table of starting reactant concentrations and initial rates. The exponents in a rate law must be determined from a table of experimental data, they are not related to stoichiometric coefficients. rate=K[A]x[B]y k is the rate constant A and B are the reactants concentrations x and y serves as the order of the corresponding compounds which can only be found experimentally a. **Reactants** - Substances initially present in a chemical reaction that are consumed during the reaction to make products. ex: 2NO2 (g) + F2 (g) ---> 2NO2F (g)
 * 4. Understand reaction mechanisms, and be able identify possible mechanism by using: **
 * 5. Be able to identify all the species in a mechanism. **

b**. Products** - formed during chemical reactions where reagents are consumed. These are the materials that are produced in results of the chemical reaction. ex: 2NO2 (g) + F2 (g) ---> 2NO2F (g)

c. **Intermediates**- A chemical species that is produced in one step of a reaction mechanism and consumed in a subsequent step. d. **Catalysts-** A substance that increases the rate of a chemical reaction without itself being consumed. It increases the rate of reaction by lowering the activation energy.

Identify the species of the following mechanism (5 pts): <span style="font-family: 'Times New Roman','serif'; font-size: 12pt; line-height: 115%;">HCl (g) + HNO3 (aq) ß à <span style="font-family: 'Times New Roman','serif'; font-size: 12pt; line-height: 115%;"> H2 (g) + NO3Cl (g) <span style="font-family: 'Times New Roman','serif'; font-size: 12pt; line-height: 115%;">NO3Cl (g) + Pt (s) à <span style="font-family: 'Times New Roman','serif'; font-size: 12pt; line-height: 115%;"> NOCl (g) + PtO2 (g) <span style="font-family: 'Times New Roman','serif'; font-size: 12pt; line-height: 115%;">HCl (aq) + HCl (aq) + NO2 (aq) à <span style="font-family: 'Times New Roman','serif'; font-size: 12pt; line-height: 115%;"> H2 (g) + Cl2 (g) <span style="font-family: 'Times New Roman','serif'; font-size: 12pt; line-height: 115%;">2 H2 (g) + PtO2 (g) à <span style="font-family: 'Times New Roman','serif'; font-size: 12pt; line-height: 115%;"> 2 H2O (l) + Pt (s) + NO2 (aq)

<span style="font-family: 'Times New Roman','serif'; font-size: 12pt; line-height: 115%; margin: 0in 0in 10pt;">Compounds to identify: HCl, HNO3, H2, NO3Cl, Pt, NOCl, PtO2, NO2, H2, Cl2, H2O. Reactants: 3HCL, HNO3 __ Products: __ NOCL, Cl2, 2H2O __ Intermediates: __ NO3Cl, 2H2, PtO2 __ Catalyst: __ Pt __ Spectator Ions: __ NO2

6. Understand and be able to use the rate laws and the integrated rate laws to determine: For a zero order Reactant: [A]final = -kt + [A]initial This changes to: t = ([A]initial - [A]final) / k. For a first order Reactant: ln[A]final = -kt + ln[A]initial This equation can be manipulated into: t = (ln[A]initial - ln[A]final) / k From here we get: t = (ln[ Ainitial / Afinal]) / k So the reaction time can be found by taking the natural log of the initial concentration divided by the final concentration and dividing that by the rate constant. For a second order Reactant: 1/[A]final = kt + 1/[A]initial This changes to: (1/[A]final - 1/[A]initial) / k. For a zero order reactant: [A]final = -kt + [A]initial For a first order reactant: ln[A]final = -kt + ln[A]initial For a second order reactant: 1/[A]final = kt + 1/[A]initial as for the concentration of products: For a zero order reactant: [A]final = [A]initial + kt  For a zero order reactant: ln[A]final = ln[A]initial + kt  For a zero order reactant: 1/[A]final = 1/[A]initial - kt
 * a. Effects of concentration of reactants in the reaction rate- **R=﻿ k [A]^n [B]^m Both A and B are measured in Molarity. So, if the concentration of reactants is greater the reaction rate will be larger ( as long as n and m are positive values). If the concentrations are lower the reaction rate will be slower.
 * b. Reaction time ** - The time it takes for all or part of a reaction to occur.
 * c. Concentration of reactants and products as a function of time ** - the concentration of reactants can be found with these equations:

7. Relationship between spontaneous reactions and chemical kinetics of the reaction The point at which a reaction begins to occur (without outside influence) is when a reaction is considered spontaneous, but it does not let us know when the reaction will be completed. Chemical kinetics provides us with a rate at which this reaction will occur.

8. Be able to correlate the following concepts: a. Changes in concentration with changes in reaction rate A reaction can only occur at a rate proportionate to the reactant of a lesser value. Also, the if a catalyst is used then no matter how concentrated the solution is, the reaction can only happen at a rate proportionate to the quantity of catalyst.

b. Presence of catalyst and transition state The transition state is the point when two molecules collide with enough energy to form a new species. A catalyst lowers the activation energy required for the collision to be effective. The collisions therefore create new species (products) more easily and at a greater pace than if the catalyst was not present during the collision.