2 Chemical KineticsIn learning chemical kinetics, you will learn how to:Predict whether or not a reaction will take place.Once started, determine how fast a reaction will proceed.Learn how far a reaction will go before it stops.
3 Rate of a Reaction Thermodynamics- Does a reaction take place? Kinetics- How fast does a reaction proceed?Chemical Kinetics- the area of chemistry concerned with the speeds or rates at which a chemical reaction occurs.Reaction Rate- the change in the concentration of a reactant or product with time. (M/s)
4 Rate of a Reaction Why do we need to know the rate of a reaction? Practical knowledge is always usefulPreparation of drugsFood processingHome repairDrugs- drug interactions, making pharmaceuticalsFood Processing- making glazes, combining foods, flavorsHome repair- grout, cement, etc.
5 Rate of a Reaction General equation for a reaction: A → BReactant → ProductIn order to monitor a reaction’s speed or rate, we can look at one of two things:Decrease in [ reactant ]Increase in [ product ]Can be represented as:rate = - Δ [A] / Δ t orrate = Δ [B] / Δ tChange in concentration (M) over time (t)
6 Rate of a ReactionProgression of a to b over time.
8 Rate of a Reaction How do we measure this experimentally? For reactions in solution:Changes in concentration can be measured spectroscopicallyFor reactions involving gases:Changes in pressure can be measuredFor reactions in solution with ions present:Change in concentrations can be measured through electrical conductanceBy definition, we know that in order to measure the rate of a reaction, we have to measure the concentration of the reactants and products….but how do we do this?
9 Rate of a ReactionSo if we have an aqueous solution of molecular bromine and formic acid, how do we determine the reaction rate?Br2(aq)+HCOOH(aq) → 2Br–(aq)+2H+(aq)+CO2 (g)time
10 Rate of a Reaction Look for color changes Molecular bromine is usually reddish-brown in color. Formic acid is colorless.As the reaction progresses, the color of the solution changes.It fades until it becomes colorless.What does this mean?If the color is fading, then the concentration of bromine is decreasing.How can you double check that this is true? Spectrophotometer. Plot wavelength vs. absorption.
11 Rate of a ReactionWavelength vs. absorption graph
12 Rate Calculations How do we calculate the rate of a reaction? We first need this information:Time (s)[reactant]Now if you are asked to perform these calculations on a test, then you will have a rate chart provided.
13 Rate CalculationsBr2 (aq) + HCOOH (aq) → 2Br– (aq) + 2H+ (aq) + CO2 (g)Table rate of reactions
14 Rate CalculationsInstantaneous Rate– rate of a reaction for a specific point in time.Average Rate vs. Instantaneous rateExamples????Average rate = meanInstantaneous rate = rate at a specific point in time
15 Rate Calculations Average Rate = -Δ [Br2] / Δt = - [Br2]final – [Br2]initial / [t]final – [t]initialInstantaneous Rate =rate for specific instance in time[Br2] / t
16 Rate CalculationsUsing this information, calculate the average rate of the bromine reaction over the first 50s of the reaction.
18 Average RateDifferent rates can be accounted for by mistakes in experiments. The ratio never changes, but the rate value may slightly change.
19 Reaction Rates and Stoichiometry For reactions more complex than A → B we cannot use the rate expression initially described.Example:2A → BDisappearance of A is twice as fast formation of BRate = - ½ Δ[A] /Δt
20 Reaction Rates and Stoichiometry In general, for the reactionaA + bB → cC + dDRate =- 1/a Δ[A] /Δt = - 1/b Δ[B] /Δt = 1/c Δ[C] /Δt = 1/d Δ[D] /Δt
21 Reaction Stoichiometry Write the rate expression for the following reaction:CH4 (g) + 2O2 (g) CO2 (g) + 2H2O (g)rate = -D[CH4]Dt= -D[O2]Dt12=D[CO2]Dt=D[H2O]Dt12
25 Rate Constant Look back to molecular bromine chart. What is k? K- the rate constant. A constant of proportionality between the reaction rate and the concentration of the reactant.K may change slightly over time.K is represented as:K = rate/ [reactant]K is not affected by the [reactant] or rate alone, since it is a ratio of these two. At any given point on a graph, k should be similar in value to it’s value at other points in the same graph.
26 Rate ConstantSmall deviations in k are only due to experimental deviations or changes in temperature.Rate vs. reactant concentration graph.K increases with concentration, but the ratio stays the same.“straight line” means direct proportionality.
27 The Rate LawRate Law- expresses the relationship of the rate of a reaction to the rate constant and the concentrations of the reactants raised to some power.Using the general reaction:aA + bB → cC + dDRate Law is:rate = k [A]x[B]yX and y are numbers that must be determined experimentally.X and y are not equal to the stoichiometric coefficients a and b.
28 The Rate Law aA + bB cC + dD Rate = k [A]x[B]y reaction is xth order in Areaction is yth order in Breaction is (x + y)th order overall
29 Reaction OrderReaction Order- the sum of the powers to which all reactant concentrations appearing in the rate law are raised.Reaction order is always defined in terms of reactant concentration.Overall reaction order- x + yExample:Rate = k [F2] [ClO2]Reaction order = firstOverall reaction order = second
30 Reaction OrderWhat is the rate expression for aA + bB → cC + dD where x=1 and y=2?Rate = k[A][B]2What is the reaction order?First in A, second in BOverall reaction order?2 +1 = 3
31 Reaction Order F2 (g) + 2ClO2 (g) 2FClO2 (g) rate = k [F2]x[ClO2]y Rate expression when x=1 and y=2?rate = k [F2]x[ClO2]y
32 Reaction OrderIf initially [F2] = 1.0M and [ClO2]=1.0M, what will happen to the reaction rate if F2 is doubled?Rate1 = k(1.0M)(1.0M)2Rate1 = k(1.0M3) [F2 ] = 1.0MRate2 = k(2.0M)(1.0M)2Rate2 = k(2.0M3) [F2 ] = 2.0MRate2 = 2 x Rate1Rate doubles
33 Reaction OrderWhat will happen in the same reaction if the [ClO2] is doubled?Rate1 = k(1.0M)(1.0M)2Rate1 = k(1.0M3) [ClO2 ] = 1.0MRate2 = k(1.0M)(2.0M)2Rate2 = k(4.0M3) [ClO2 ] = 2.0MRate2 = 4 x Rate1
34 Determination of Rate Law F2 (g) + 2ClO2 (g) FClO2 (g)Experiment[F2][ClO2]Rate (M/s)10.040.031.0x10-222.0x10-230.0240.06
35 Determination of Rate Law Experiments 1 & 4As [F2] doubles, so does the rateExperiments 2 & 3As [ClO2] doubles, so does the rate2:2 ratio…..1:1 ratiox = 1 and y = 1Rate = k [F2] [ClO2]
36 Rate law/Expression Calculations Determine the rate law and calculate the rate constant for the following reaction from the following data:S2O82- (aq) + 3I- (aq) SO42- (aq) + I3- (aq)Experiment[S2O82-][I-]Initial Rate (M/s)10.080.0342.2 x 10-420.0171.1 x 10-430.16Double [I-], rate doubles (experiment 1 & 2)Double [S2O82-], rate doubles (experiment 2 & 3)rate = k [S2O82-]x[I-]yy = 1x = 1rate = k [S2O82-][I-]k =rate[S2O82-][I-]=2.2 x 10-4 M/s(0.08 M)(0.034 M)= 0.08/M•s
37 Rate Law/Reaction Order Rate laws are always determined experimentallyReaction order is always defined in terms of reactantReactant order is not related to the stoichiomteric coefficient in the overall reaction.F2 (g) + 2ClO2 (g) FClO2 (g)rate = k [F2][ClO2]
38 Relation between Reactant Concentration and Time First Order Reaction- a reaction whose rate depends on the reactant concentration raised to the first power.Reaction Type: A→BRate of: -Δ [A]/Δt or k[A]Combining and simplifying these equations brings us to the following rate equation:ln[A]t = -kt + ln[A0]Ln equation has the form of the linear equation y=mx+b. ln At= y. k is the slope of the line. t=0 is start time. T=t is time chosen. Ln A0= y intercept.
39 Relation between Reactant Concentration and Time Decrease in reactant concentration with time. (b) will allow you to determine –k….slope and ln[A0]…..y-intercept.
40 Reaction TimeThe reaction 2A B is first order in A with a rate constant of 2.8 x 10-2 s-1 at 800C. How long will it take for A to decrease from 0.88 M to 0.14 M ?[A]0 = 0.88 Mln[A] = ln[A]0 - kt[A] = 0.14 Mkt = ln[A]0 – ln[A]ln0.88 M0.14 M2.8 x 10-2 s-1=ln[A]0[A]k=ln[A]0 – ln[A]k= 66 st =
49 Reaction Half-lifeAs a reaction proceeds, the concentrations of the reactants decreases.Another way to measure [reactant] over time is to use the half-life.Half-life, t1/2 – the time required for the concentration of a reactant to decrease to half of its initial concentration.
50 Reaction Half-lifeExpression for half-life of a first order reaction is:t1/2 = ln2/kort1/2 = 0.693/kWhat can you tell about the half-life of a reaction from this equation?The half-life of a first order reaction is independent of the concentration of the reactant.Measuring the half-life of a reaction is one way to determine the rate constant of a first order reaction.
51 Reaction Half-lifeCollege student’s four years of undergraduate work. Independent of how many students are present.
52 Reaction Half-lifeWhat is the half-life of N2O5 if it decomposes with a rate constant of 5.7 x 10-4 s-1?t½ln2k=0.6935.7 x 10-4 s-1=t½t½= 1200 st½= 20 minutes
54 Second-Order Reactions Second-order reaction- a reaction whose rate depends on the concentration of one reactant raised to the second power OR on the concentrations of two different reactants, each raised to the first power.Simple Type: A→Brate = k[A]2Complex Type: A + B→Crate = k[A][B]
55 Second-order Reactions For A→B, the following expression is used:1[A]=[A]0+ ktINTEGRATED RATE LAW has the form of a linear equation.Plot of 1/[A]t vs. t gives a staright line with a slope of k and a y-intercept of 1/[A]0.
56 Half-life of a Second-order Reaction Equation for half-lifeWhat is the difference between this equation and the equation for half-life of first-order reactions?t½ =1k[A]0This equation relies on the concentration of the reactants. The half-life is inversely proportional to the initial reaction concentration.Makes sense because in the first part of the reaction, the half-life should be shorter in the early stages. This is when more reactant molecules are present to collide with one another.If you didn’t know if you had a first or second order reaction, how could you tell mathematically? Measure the half-life at different initial concentrations. Look for variations.
61 Activation Energy and Temperature Dependence of Rate Constants Reactions usually take place at a raster rate when temperature is increased. Ex: cooking spaghetti in water at 80 degrees vs. 100 degrees.Why do people put food in freezers? To preserve food and stop bacterial decay.
62 The Collision Theory of Chemical Kinetics Gas molecules frequently collide with one anotherExpect that the rate of a reaction is equivalent to the number of collisionsReaction rate is dependent on concentrationBillions of collisions every second in the gas phase of a reaction. Not every collision causes a reaction. If that happened, most reactions would be instantaneous.Molecules need to have enough kinetic energy upon collision to break apart to cause vibration and break apart compounds.If not enough kinetic energy, molecules remain intact and no change results from the collisions.
63 The Collision Theory of Chemical Kinetics Activation Energy (Ea)- the minimum amount of energy required to initiate a chemical reaction.Activated Complex (Transition State)- a temporary species formed by the reactant molecules as a result of the collision before they form the product.Exothermic. Products are more stable than reactants and after product formation, there is release of heat.endothermic.
64 The Collision Theory of Chemical Kinetics What does this have to do with temperature?High energy moleculesHigh temperaturesIncreased product formation
65 The Collision Theory of Chemical Kinetics Factors that affect rate1.2.3.Collision frequency, temperature, activation energy.
66 The Arrhenius Equation Relation between activation energy and temperature.lnk = (Ea/R) x (1/T) + lnAShows dependence or rate on temperature.Plot of lnk vs. 1/t gives a straight line. Can use to determine activation energy of a reaction.
72 Activation Energy, Reaction Rates and Temperature As stated earlier, for a reaction to take place, molecules must posses enough kinetic energy.Kinetic energy must be higher than Ea.Each reaction takes place at a specific temperature……but what happens if we adjust this temp.?
73 Activation Energy, Reaction Rates and Temperature Increasing Temperature leads to:Molecules reach high ke fasterNumber of molecules with high enough ke increasesReaction rate increasesRate of a reaction doubles for every 10 degrees reaised.
74 CatalystsA catalyst is defined by the ability of a substance to do each of the following:Catalysts increase the rate of reaction.Catalysts are not consumed by the reaction.A small quantity of catalyst should be able to affect the rate of reaction for a large amount of reactant.Catalysts do not change the equilibrium constant for the reaction.The first criterion provides the basis for defining a catalyst as something that increases the rate of a reaction. The second reflects the fact that anything consumed in the reaction is a reactant, not a catalyst. The third criterion is a consequence of the second; because catalysts are not consumed in the reaction, they can catalyze the reaction over and over again. The fourth criterion results from the fact that catalysts speed up the rates of the forward and reverse reactions equally, so the equilibrium constant for the reaction remains the same.
75 CatalystsHeterogeneous catalyst- the reactants and the catalyst are in different phases. catalyst = solid reactants = liquid/gasHomogeneous catalyst- catalyst and reactants are in the same phase, usually liquid.
76 CatalystsCatalysts lower the Ea, so that more molecules can reach the ke and proceed to product.More collisions occur and reaction rate increases.
77 Enzyme Catalysts Biological catalyst. Lock and key method. Normally substrate converts to products slowly.With enzyme, reaction speeds up drastically.