ABRIDGED CHE1031 Lecture 10: Reaction kinetics Lecture 10 topics Brown chapter 14 10.1 Reaction rates Factors that effect reaction rates 14.1 Visualizing rates & units 14.2 Average reaction rates Instantaneous reaction rates Stoichiometry & reaction rates 10.2 Concentration & reaction rates 14.3 Rate laws Reaction orders 10.3 Change in concentration with time 14.4 First- & second-order reactions Half-life 10.4 Temperature & reaction rate 14.5 Collision, orientation & Ea 10.5 Reaction mechanisms 14.6 Elementary Multistep Rate-limiting steps 10.6 Catalysis ` 14.7 ABRIDGED

Kinetics describe the speed of reactions Four factors affect kinetics. Instantaneous and average rates can be measured. Stoichiometry affects kinetics The general process of advancing scientific knowledge by making experimental observations and by formulating hypotheses, theories, and laws. It’s a systematic problems solving process AND it’s hands-on….. Experiments must be done, data generated, conclusions made. This method is “iterative”; it requires looping back and starting over if needed. [Why do you think they call it REsearch?] Often years, decades or more of experiments are required to prove a theory. While it’s possible to prove a hypothesis wrong, it’s actually NOT possible to absolutely prove a hypothesis correct as the outcome may have had a cause that the scientist hasn’t considered.

Kinetics & relevant factors the study of the speed at which chemical reactions occur (aka study of reaction rates) Range from very rapid to extremely slow – “glacial” Factors that affect kinetics? Physical state of reactants Concentration of reactants Temperature of reaction Presence of a catalyst Physical state controls two factors: Frequency of collision – rxn rate increases with energy level, so gas > liq > solid Surface area of contact – highest for gases & liquids; can be increased in solids Increased concentration increases reaction rate Mainly because higher concentrations favor frequent collisions between reactants Increased temperature increases reaction rate Mainly because higher temperatures favor high energy physical states & rapid motion Again, increases rate of collision Catalysts: substances that increase rate of rxn without, themselves, being changed. Biological systems depend on catalysts called enzymes. Catalytic converters increase the rate at which pollutants are broken down. p. 557 - 9

Reactant concentration increases rate Steel wool burns more rapidly in pure O2 than in air. Flask full of oxygen gas p. 557 - 9

Visualizing reaction rate Imagine a reaction like this: A  B  Units or reaction rate? Usually measure as concentration of reactant or product per unit time. Rate is measured in units of M/s. p. 559 - 61

Calculating average reaction rates Again using this example: A  B Average rate of appearance of B = change in concentration of B = [B] @ t2 - [B] @ t1 = Δ [B] t2 – t1 Δt = 0.46 M – 0.00 M = 2.3x10-2 M/s 20 s – 0 s Average rate of disappearance of A = - change in concentration of A = - ([A] @ t2 – [A] @ t1) = - Δ [A] t2 – t2 Δt = 0.54 M - 1.00 M = - 2.3x10-2 M/s 20 s – 0 s Notes: Average rates can only be calculated over a defined period of time. Rates are always expressed in M/s. p. 559 - 61

Reaction rate example: Calculate the average rate at which A disappears between 20 and 40 s. Calculate the rate at which B appears between 0 and 40 s. Avg rate = - Δ [A] = - (0.30 M – 0.54 M) = -(-0.24 M) = 0.012 M/s Δt 40 s – 20 s 20 s Avg rate = Δ [B] = 0.70 M – 0 M) = 0.70 M = 0.018 M/s Δt 40 s – 0 s 40 s Is it always true that - over a given period of time – rate of loss of A = rate of gain of B? Depends on reaction stoichiometry! p. 559 - 61

Rates may change over time! Note that the rate of the chemical reaction shown here slowed over time. The curve starts steep and becomes more gradual as concentration of reactant falls. C4H9Cl + H2O  C4H9OH + HCl Instantaneous rates can be calculated by drawing a line tangent to a time point and then determining the slope of this tangent line. slope = rise/run Calculate the instantaneous rate at 600 s. Inst. Rate = Δ [C4H9Cl] Δt = - (0.017 M – 0.042 M) 800 s – 400 s = 6.3x10-5 M/s p. 561 - 62

Stoichiometry & reaction rates Stoichiometry affects reaction rate when coefficients ≠ 1. 2HI  H2 + I2 For every 2 moles of HI that disappear, 1 mole of H2 and 1 mole of I2 appear. In general, stoichiometric coefficients become fractional rate coefficients: aA + bB  cC + dD Rate = - 1 Δ[A] = - 1 Δ[B] = 1 Δ[C] = 1 Δ[D] a Δt b Δt c Δt d Δt Apply this to creating formulas for calculating rates for each reactant & products shown in the reaction of HI. Rate = - 1 Δ[HI] = Δ[H2] = Δ[I2] 2 Δt Δt Δt p. 562 - 3

Temperature & kinetics Rates ~double with 10C increase. How temperature increases rate Activation energy & transition states The general process of advancing scientific knowledge by making experimental observations and by formulating hypotheses, theories, and laws. It’s a systematic problems solving process AND it’s hands-on….. Experiments must be done, data generated, conclusions made. This method is “iterative”; it requires looping back and starting over if needed. [Why do you think they call it REsearch?] Often years, decades or more of experiments are required to prove a theory. While it’s possible to prove a hypothesis wrong, it’s actually NOT possible to absolutely prove a hypothesis correct as the outcome may have had a cause that the scientist hasn’t considered.

Temperature & rate Generally, temperature increases the rate of chemical reactions. A rough estimate is that rate doubles for every 10°C increase in temperature. p. 575 – 6

How does temperature increase rate? There are three ways in which temperature can increase kinetics: Collision Orientation 3. Activation Energy Molecules must collide in order to react; temperature speeds motion & increases chances of constructive collision. At room temp gas molecules collide 1010 s; however, only one of every 1013 collisions results in reaction. - Molecules be oriented so that atoms fated to form new bonds contact each other during collisions. Energy barrier that must be overcome in order for a chemical reaction to occur. p. 576 - 7

Activation energy Energy (Ea) must be input in order to change the conformation of the reactant from its original state, through an intermediate transition state. The final reaction can then occur. p. 577

Transition states Note that temperature is able to cause changes in molecular shape (aka conformation). p. 577 - 80

Elementary vs. multistep reactions Reaction mechanisms Elementary vs. multistep reactions Rate-limiting steps Rate-limiting step The general process of advancing scientific knowledge by making experimental observations and by formulating hypotheses, theories, and laws. It’s a systematic problems solving process AND it’s hands-on….. Experiments must be done, data generated, conclusions made. This method is “iterative”; it requires looping back and starting over if needed. [Why do you think they call it REsearch?] Often years, decades or more of experiments are required to prove a theory. While it’s possible to prove a hypothesis wrong, it’s actually NOT possible to absolutely prove a hypothesis correct as the outcome may have had a cause that the scientist hasn’t considered.

Rate-limiting steps When chemical reactions require more than one step, their overall rate is often limited by the slowest of the steps. So this slowest stop is called the rate-limiting step because it limits the overall rate of reaction. Step 1: NO2 + NO2  NO3 + NO (slow) Step 2: NO3 + CO  NO2 + CO2 (fast) Overall: NO2 + CO  NO + CO2 k1 k2 >> k1 k2 What is the rate law of the overall reaction? Because step 1 is much slower than step 2, it is rate-limiting. The rate of the overall reaction is equal to the rate of the slow step (1). Step 1 is bimolecular, so rate = k1[NO2]2 p. 584 - 5

Rate-limiting examples Nitrous oxide decomposes by a two-step mechanism. N2O  N2 + O (slow) N2O + O  N2 + O2 (fast) Write the equation for the overall reaction. Write the rate law for the overall reaction. a) 2N2O  2N2 + O2 b) Rate = k[N2O]2 Ozone reacts with nitrogen dioxide by a two-step mechanism: Step 1: O3 + NO2  NO3 + O2 Step 2: NO3 + NO2  N2O5 Overall: O3 + 2NO2  N2O5 + O2 Overall experimental rate law is: rate = k[O3][NO2]. Which step is slower? Step 1 is the slow step, since it is used in the rate law for the overall reaction. p. 584 - 5

Catalysts speed reactions Homogeneous vs. heterogeneous Catalysts lower activation energy. Enzymes are nature’s catalysts. The general process of advancing scientific knowledge by making experimental observations and by formulating hypotheses, theories, and laws. It’s a systematic problems solving process AND it’s hands-on….. Experiments must be done, data generated, conclusions made. This method is “iterative”; it requires looping back and starting over if needed. [Why do you think they call it REsearch?] Often years, decades or more of experiments are required to prove a theory. While it’s possible to prove a hypothesis wrong, it’s actually NOT possible to absolutely prove a hypothesis correct as the outcome may have had a cause that the scientist hasn’t considered.

Catalysts Catalyst – a substance that increases the speed of a chemical reaction without undergoing any permanent change itself. Catalysts are common in biochemistry (living organisms) & chem in general. Here the salt, NaBr, dissociates, becoming part of the solution & is therefore a homogeneous catalyst. p. 589 - 93

How do catalysts work? In general, catalysts decrease the Ea, activation energy, needed to cause a reaction to happen. Catalysts can do this in a number of ways: Supplies a very reactive intermediate (like Br2). Makes a reactant more “reactive”. Brings reactants into contact. Br ion forms the Br2 intermediate -very reactive – which then oxidizes H2O2 into water & O2. p. 589 - 93

Heterogeneous catalyst: increases reactivity Heterogeneous catalysts – exist in a different phase than the reactants. Here a layer of metal atoms (Ni, Pt, Pd) helps to separate the two H atoms of H2 gas, causing the H atoms to become more reactive. The H’s then add to the organic molecule ethylene, reducing it to ethane. Note that the metal surface also “concentrates” the gas-phase reactants by “pulling” them down to the metal’s surface. This brings reactants into close proximity & also increases the speed of rxn. C2H4 + H2  C2H6 Note that Pt is the catalyst here & does not appear in the equation. p. 589 - 93

Enzymes: nature’s catalysts Enzymes – biological molecules, mainly proteins, that lower reaction Ea & therefore increase reaction rate. Enzymes work by a “lock & key” mechanism which can have several aspects: Brings reactants into close physical proximity. Provides chemically reactive groups Flexes and provides motion to potentiate rxn. enzyme reactant Some toxins & poisons (like mercury & nerve gas) kill by inhibiting enzymes & preventing our body from doing necessary chemistry. p. 589 - 93

CHE1031 Lecture 10: Reaction kinetics Lecture 10 topics Brown chapter 14 1. Reaction rates Factors that effect reaction rates 14.1 Visualizing rates & units 14.2 Average reaction rates Instantaneous reaction rates Stoichiometry & reaction rates 2. Concentration & reaction rates 14.3 Rate laws Reaction orders 3. Change in concentration with time 14.4 First- & second-order reactions Half-life 4. Temperature & reaction rate 14.5 Collision, orientation & Ea 5. Reaction mechanisms 14.6 Elementary Multistep 6. Catalysis ` 14.7

Lecture 10: Terms to know Kinetics Rate Average rate Instantaneous rate Tangent Rate law Rate constant Orders of reaction Overall order of reaction First vs. second order reactions Half-life Collision Orientation Activation energy Transition state Reaction pathway Reaction mechanisms (elementary vs. multistep) Rate-limiting step Catalyst (homogeneous vs. heterogenous) Enzymes PLEASE note that the formula calculates a weighted average, so there’s not need to add and then divide the sum by the number of isotopes in the problem. This is an incredibly common student error so don’t get caught!