Unit 9: Kinetics and Equilibrium

Unit 9: Kinetics and Equilibrium Chapter 13 Pre-AP Chemistry I Edmond North High School

Collision Theory • The following three statements summarize the collision theory. • 1. Particles must collide in order to react. • 2. The particles must collide with the correct orientation. • 3. The particles must collide with enough energy to form an activated complex, which is an intermediate particle made up of the joined reactants. • An effective collision is one that results in a reaction; new products are formed.

Activated Complex • The Activated Complex is an intermediate particle formed when reactants collide and stick together. • Old bonds are breaking while new bonds are forming. • The minimum amount of energy colliding particles must have in order to form an activated complex is called the activation energy. • Particles that collide with less than the activation energy cannot form an activated complex.

Exothermic Reactions Review • An exothermic reaction releases heat, and an endothermic reaction absorbs heat.

Endothermic Reactions Review • The endothermic reaction absorbs heat because the products are at a higher energy level than the reactants.

Kinetics • Kinetics is the study of reaction rates and the factors that effect them. • Why is it important? • When we understand reaction rates we can control chemical reactions and use them for specific purposes.

Kinetics • As a reaction occurs: • concentration of reactants decreases • concentration of products increases • Kinetics studies the rate at which a chemical process occurs. • Besides information about the speed at which reactions occur, kinetics also sheds light on the reaction mechanism (exactly how the reaction occurs).

Reaction Rate • Reaction Rate is the number of atoms, ions or molecules that react in a given time to form products • We measure either the rate of disappearance of the reactants or the rate of appearance of one of the products • In simpler terms: • The average rate is the change in a given quantity during a specific period of time.

Measuring Reaction Rates • Change in Electrical Conductivity • Change in Color • Change in Pressure • Change in Volume

Reaction Rate Graph • For the reaction H2 + I2 → 2HI, • What is happening to the concentration of the reactants? • What is happening to the concentration of the products?

Reaction Rates & Stoichiometry • We use coefficients to give us a ratio of rates between species in a reaction. • In the reaction 3H2 + N2 2NH3, since N2 has a coefficient of 1 and NH3 has a coefficient of 2, the rate N2 is used = 2(rate NH3 is created) • Likewise, since H2 has a coefficient of 3 and N2 has a coefficient of one, it is used 3 times as fast as N2.

Factors Affecting Reaction Rates • The reaction rate for almost any chemical reaction can be modified by varying the conditions of the reaction. • An important factor that affects the rate of a chemical reaction is the reactive nature of the reactants. As you know, some substances react more readily than others. • The more reactive a substance is, the faster the reaction rate.

Factors That Affect Reaction Rates • Physical State of the Reactants • In order to react, molecules must come in contact with each other. • The more homogeneous the mixture of reactants, the faster the molecules can react. • Concentration of Reactants • As the concentration of reactants increases, so does the likelihood that reactant molecules will collide.

Factors that Effect Reaction Rate • Particle Size (surface area) • For solids, breaking up big pieces increases surface area, increasing rate by having more places for the molecules to interact • Ionic compounds have more surface area than covalent

Factors that Effect Reaction Rate • Pressure – Gases only! • As pressure increases the concentration increases, so you will have more collisions

Factors That Affect Reaction Rates • Temperature • At higher temperatures, reactant molecules have more kinetic energy, move faster, and collide more often and with greater energy. • Presence of a Catalyst • Catalysts speed up reactions by changing the mechanism of the reaction. • Catalysts are not consumed during the course of the reaction.

Catalyst

Catalysts • Catalysts increase the rate of a reaction by decreasing the activation energy of the reaction. • Catalysts change the mechanism by which the process occurs. • One way a catalyst can speed up a reaction is by holding the reactants together and helping bonds to break.

How Do Catalysts Work? • Instead of increasing KE, the catalyst lowers the Ea • It creates a new pathway with a lower activation energy • It is never consumed in the reaction so is not considered a reactant • Enzymes are biological catalysts. • Homogenous Catalysts are in the same phase as the reactants. • Heterogeneous Catalysts are in a different phase as the reactants.

Factors that Effect Reaction Rate • Inhibitors • Compounds added to a reaction that slow it down • Examples: Lead in diesel, preservatives in food.

Reaction Rate Laws • Rate Laws relate the reaction rate and reactant concentration • We write the rate laws for elementary reactions from the equation. • They include only the gaseous or aqueous reactants! • Solids & liquids are not included because they have a constant concentration that does not increase or decrease

Rate Law • A Rate Law is an expression that shows how the rate depends on concentration of reactants • Example: • Rate laws use reactants only • k is proportionality constant (called the rate constant) • n is the order (exponent) of the reactant determined experimentally • n is generally the coefficient of the chemical in the balanced equation HOWEVER experimental data overrides the coefficient. • Reaction is xth order in A • Reaction is yth order in B • Reaction is (x +y)th order overall aA + bB cC + dD Rate = k [A]x[B]y

Order of Reaction • Order of Reaction • For the reaction with the rate = k[A]2[B] • Rate is 2nd order with respect to A • Rate is 1st order with respect to B • Rate is 3rd order overall (add exponents) • How does concentration affect reaction rate? • If [A] doubles what happens to the rate? • If [B] triples, what happens to the rate?

Initial Rate Method • Reactions are reversible. • As products accumulate they can begin to turn back into reactants. • Early on the rate will depend on only the amount of reactants present. • We want to measure the reactants as soon as they are mixed. • This is called the Initial Rate Method.

Method of Initial Rates • Used to find the form of the rate law • Choose one reactant to start • Find two experiments where the concentration of that reactant changes but all other reactants stay the same • Write the rate laws for both experiments • Divide the two rate laws • Use log rules to solve for the order • Follow the same technique for other reactants

Determining Experimental Rate Law • What is the order with respect to A? • What is the order with respect to B? • What is the overall order of the reaction? • What is the value for k?

Determining Experimental Rate Law • What is the order with respect to Cl2? • What is the order with respect to NO? • What is the overall order of the reaction? • What is the value for k?

Remember! • Two key points • The concentration of the products do not appear in the rate law because this is an initial rate. • The order must be determined experimentally, can’t be obtained from the equation

Temperature and Rate • Generally, as temperature increases, so does the reaction rate. • This is because k is temperature dependent.

Reaction Mechanisms • In a complex reaction, a certain series of steps must occur with correct fit and in an exact sequence in order to yield the reaction products. This series of steps is called a reaction mechanism. • When we write a chemical equation we show the overall reaction (reactants → products) • We call the compounds that form/decompose during the steps but are not in the overall reaction Intermediates.

Reaction Mechanisms • Elementary Reactions, are the simplest one-step reactions in a reaction mechanism. • Example: • HBr + O2 → HOOBr (slow) • HOOBr + HBr → 2HOBr (fast) • 2HOBr + 2HBr → 2H2O + 2Br2 (fast) • The Rate-determining step is the slowest step in the chemical reaction. • Which is the rate-determining step for the reaction above?

Multistep Mechanisms • There is an activation energy for each elementary step. • In a multistep process, one of the steps will be slower than all others. • Slowest step (rate determining) must have the highest activation energy. • The overall reaction cannot occur faster than this slowest, rate-determining step.

Intermediates 2NO2 + F2 2NO2F • Rate = k[NO2][F2] • The proposed mechanism is • NO2 + F2 NO2F + F (slow) • F + NO2 NO2F (fast) • F is called an intermediate. It is formed then consumed in the reaction.

Reaction Intermediates • Example: • 2 Steps: NO2(g) + NO2 (g)  NO3 (g) + NO(g) • NO3(g) + CO(g)  NO2 (g) +CO2 (g) • Overall: NO2(g) + CO(g)  NO(g) + CO2(g)

Slow Initial Step • A proposed mechanism for this reaction is • Step 1: NO2 + NO2 NO3 + NO (slow) • Step 2: NO3 + CO  NO2 + CO2 (fast) • The NO3 intermediate is consumed in the second step. • As CO is not involved in the slow, rate-determining step, it does not appear in the rate law: Rate = k [NO2]2 • CO is necessary for this reaction to occur, but the rate of the reaction does not depend on its concentration. NO2(g) + CO (g) NO (g) + CO2(g)

Fast Reactions • 2 IBr I2+ Br2 • Mechanism • IBr I + Br (fast) • IBr + Br  I + Br2 (slow) • I + I  I2 (fast) • Rate = k[IBr][Br] but [Br]= [IBr] • Rate = k[IBr][IBr] = k[IBr]2

Indicators of Completed Reactions • So how do you know when a reaction has gone to completion? • A gas is produced and escapes (open container) • A precipitate forms from two aqueous solutions. • A covalent product is formed (usually water)

What is Equilibrium? • When a reaction results in almost complete conversion of reactants to products the reaction goes to completion. • Most reactions, however, do not go to completion. They appear to stop. • The reason is that these reactions are reversible. • A reversible reaction is one that can occur in both the forward and the reverse directions.

Chemical Equilibrium • Chemical equilibrium occurs when two opposite reactions occurring at the same time and rate A + B ↔ AB • Forward & Reverse reactions don’t stop at equilibrium, just looks that way because concentration remains constant • Rateforward reaction = Ratereverse reaction

Equilibrium Example • You have a bridge between 2 cities. The number of cars going in either direction on the bridge are equal (at equilibrium) • The populations of the cities on either side of the bridge do not have to be equal!

Equilibrium Position • The equilibrium position of a reaction is determined by: • Initial concentrations • Energy of reactants and products • Degree of organization of reactants and products • GOAL of reactions at equilibrium is to: • Lower energy

Progression of Equilibrium • As a reaction progresses: • [A] decreases to a constant, • [B] increases from zero to a constant. • When [A] and [B] are constant, equilibrium is achieved. • In a system at equilibrium, both the forward and reverse reactions are being carried out; as a result, we write its equation with a double arrow.

A System at Equilibrium • Once equilibrium is achieved, the amount of each reactant and product remains constant. • Rates are equal; concentrations are not. • The concentrations do not change at equilibrium.

Equilibrium Positions • If equilibrium lies “to the left”: • There are more reactants and less products. • If equilibrium lies “to the right”: • There are less reactants and more products. • If reactants are mixed and concentrations do not change: • The reaction could already be at equilibrium. • Reaction rates are so slow that change is too difficult to detect.

Equilibrium Expressions • Equilibrium expressions relate concentrations of reactants to those of the products • These can be written from balanced equations • They look similar to rate laws! BE CAREFUL! • Uses both reactants & products • Include only the gaseous or aqueous phases as with rate laws

[NO2]2 [N2O4] Keq = kf kr = Deriving an Equilibrium Expression • Forward reaction: N2O4 (g)  2 NO2 (g) • Rate law: Rate = kf [N2O4] • Reverse reaction: 2 NO2 (g)  N2O4 (g) • Rate law: Rate = kr [NO2]2 • Therefore, at equilibrium: Ratef = Rater • kf [N2O4] = kr [NO2]2 • Rewriting this, it becomes

General Equilibrium Expressions • A (s) + 2B (g) ↔ 2C (g) + 3D (g) Keq = [C]2[D]3 [B]2 • Keq = equilibrium constant (capital K) • Numerical value of the ratio of product concentrations to reactant concentrations • [ ] = concentration in M (mol/L) • Order = coefficient becomes exponent

PbCl2 (s) Pb2+ (aq) + 2 Cl−(aq) Solids and Liquids Are Constant • Both can be obtained by dividing the density of the substance by its molar mass—and both of these are constants at constant temperature. • Therefore, the concentrations of solids and liquids do not appear in the equilibrium expression. Kc = [Pb2+] [Cl−]2

Equilibrium Constant, K • If we know the value of K, we can predict: • The tendency of a reaction to occur. • If a set of concentrations could be at equilibrium. • The equilibrium position, given initial concentrations. • K will always have the same value at a certain temperature. • No matter what amounts are added, the ratio at equilibrium will always be same

Equilibrium Constant (Keq) • Tells you whether the products or reactants are favored • Keq > 1 • Products are favored (forward rxn); lots of product is made • Keq < 1 • Reactants are favored (reverse rxn); not much product made • The size of K and time needed to reach equilibrium are NOT related • This is determined by reaction rate (Ea)

Unit 9: Kinetics and Equilibrium