1 of 39 Boardworks Ltd 2007 What does
1 of 39 © Boardworks Ltd 2007
What does rate of reaction mean? The speed of different chemical reactions varies hugely. Some reactions are very fast and others are very slow. The speed of a reaction is called the rate of the reaction. What is the rate of these reactions? rusting baking explosion slow fast very fast 2 of 39 © Boardworks Ltd 2007
Rates of reaction Why are some reactions faster than others? 3 of 39 © Boardworks Ltd 2007
Reactions, particles and collisions Reactions take place when particles collide with a certain amount of energy. The minimum amount of energy needed for the particles to react is called the activation energy, and is different for each reaction. The rate of a reaction depends on two things: l the frequency of collisions between particles l the energy with which particles collide. If particles collide with less energy than the activation energy, they will not react. The particles will just bounce off each other. 4 of 39 © Boardworks Ltd 2007
Changing the rate of reactions Anything that increases the number of successful collisions between reactant particles will speed up a reaction. What factors affect the rate of reactions? l increased temperature l increased concentration of dissolved reactants, and increased pressure of gaseous reactants l increased surface area of solid reactants l use of a catalyst. 5 of 39 © Boardworks Ltd 2007
Slower and slower! Reactions do not proceed at a steady rate. They start off at a certain speed, then get slower and slower until they stop. As the reaction progresses, the concentration of reactants decreases. This reduces the frequency of collisions between particles and so the reaction slows down. 0% 25% reactants product 6 of 39 50% 75% 100% percentage completion of reaction © Boardworks Ltd 2007
Graphing rates of reaction 7 of 39 © Boardworks Ltd 2007
Reactant–product mix 8 of 39 © Boardworks Ltd 2007
How can rate of reaction be measured? Measuring the rate of a reaction means measuring the change in the amount of a reactant or the amount of a product. What can be measured to calculate the rate of reaction between magnesium and hydrochloric acid? magnesium + hydrochloric magnesium acid chloride + hydrogen l The amount of hydrochloric acid used up (cm 3/min). l The amount of magnesium chloride produced (g/min). l The amount of hydrogen product (cm 3/min). 9 of 39 © Boardworks Ltd 2007
Setting up rate experiments What equipment is needed to investigate the rate of hydrogen production? glass tube conical flask rubber connecter gas syringe rubber bung hydrochloric acid magnesium 10 of 39 © Boardworks Ltd 2007
Calculating rate of reaction from graphs hydrogen produced (cm 3) How can the rate of reaction be calculated from a graph? 70 60 x 50 rate of reaction = y x 40 30 y 20 10 0 0 10 20 30 time (seconds) 40 50 The gradient of the graph is equal to the initial rate of reaction at that time rate of reaction = 45 cm 3 rate of reaction = 2. 25 cm 3/s 20 s 11 of 39 © Boardworks Ltd 2007
The reactant/product mix 12 of 39 © Boardworks Ltd 2007
Collisions and reactions: summary 13 of 39 © Boardworks Ltd 2007
14 of 39 © Boardworks Ltd 2007
Temperature and collisions How does temperature affect the rate of particle collision? 15 of 39 © Boardworks Ltd 2007
Effect of temperature on rate The higher the temperature, the faster the rate of a reaction. In many reactions, a rise in temperature of 10 °C causes the rate of reaction to approximately double. Why does increased temperature increase the rate of reaction? At a higher temperature, particles have more energy. This means they move faster and are more likely to collide with other particles. When the particles collide, they do so with more energy, and so the number of successful collisions increases. 16 of 39 © Boardworks Ltd 2007
Temperature and particle collisions 17 of 39 © Boardworks Ltd 2007
Temperature and batteries Why are batteries more likely to rundown more quickly in cold weather? At low temperatures the reaction that generates the electric current proceeds more slowly than at higher temperatures. This means batteries are less likely to deliver enough current to meet demand. 18 of 39 © Boardworks Ltd 2007
How does temperature affect rate? The reaction between sodium thiosulfate and hydrochloric acid produces sulfur. sodium hydrochloric thiosulfate + acid Na 2 S 2 O 3 (aq) + 2 HCl (aq) sodium chloride 2 Na. Cl (aq) + sulfur dioxide + SO 2 (g) + sulfur + S (s) + water + H 2 O (l) Sulfur is solid and so it turns the solution cloudy. How can this fact be used to measure the effect of temperature on rate of reaction? 19 of 39 © Boardworks Ltd 2007
The effect of temperature on rate 20 of 39 © Boardworks Ltd 2007
21 of 39 © Boardworks Ltd 2007
Effect of concentration on rate of reaction The higher the concentration of a dissolved reactant, the faster the rate of a reaction. Why does increased concentration increase the rate of reaction? At a higher concentration, there are more particles in the same amount of space. This means that the particles are more likely to collide and therefore more likely to react. lower concentration 22 of 39 higher concentration © Boardworks Ltd 2007
Concentration and particle collisions 23 of 39 © Boardworks Ltd 2007
The effect of concentration on rate 24 of 39 © Boardworks Ltd 2007
Effect of pressure on rate of reaction Why does increasing the pressure of gaseous reactants increase the rate of reaction? As the pressure increases, the space in which the gas particles are moving becomes smaller. The gas particles become closer together, increasing the frequency of collisions. This means that the particles are more likely to react. lower pressure 25 of 39 higher pressure © Boardworks Ltd 2007
26 of 39 © Boardworks Ltd 2007
Effect of surface area on rate of reaction Any reaction involving a solid can only take place at the surface of the solid. If the solid is split into several pieces, the surface area increases. What effect will this have on rate of reaction? low surface area high surface area This means that there is an increased area for the reactant particles to collide with. The smaller the pieces, the larger the surface area. This means more collisions and a greater chance of reaction. 27 of 39 © Boardworks Ltd 2007
Surface area and particle collisions 28 of 39 © Boardworks Ltd 2007
Reaction between a carbonate and acid Marble chips are made of calcium carbonate. They react with hydrochloric acid to produce carbon dioxide. calcium carbonate Ca. CO 3 (aq) + + hydrochloric acid 2 HCl (aq) calcium chloride Ca. Cl 2 (aq) + water + H 2 O (aq) + carbon dioxide + CO 2 (g) The effect of increasing surface area on the rate of reaction can be measured by comparing how quickly the mass of the reactants decreases using marble chips of different sizes. 29 of 39 © Boardworks Ltd 2007
The effect of surface area on rate 30 of 39 © Boardworks Ltd 2007
What are catalysts? Catalysts are substances that change the rate of a reaction without being used up in the reaction. Catalysts never produce more product – they just produce the same amount more quickly. energy (k. J) Ea without catalyst Different catalysts work in different ways, but most lower the reaction’s activation energy (Ea). Ea with catalyst reaction (time) 31 of 39 © Boardworks Ltd 2007
Everyday catalysts Many catalysts are transition metals or their compounds. For example: l Nickel is a catalyst in the production of margarine (hydrogenation of vegetable oils). l Iron is a catalyst in the production of ammonia from nitrogen and hydrogen (the Haber process). l Platinum is a catalyst in the catalytic converters of car exhausts. It catalyzes the conversion of carbon monoxide and nitrogen oxide into the less polluting carbon dioxide and nitrogen. 32 of 39 © Boardworks Ltd 2007
Catalysts in industry Why are catalysts so important for industry? l Products can be made more quickly, saving time and money. l Catalysts reduce the need for high temperatures, saving fuel and reducing pollution. Catalysts are also essential for living cells. Biological catalysts are special types of protein called enzymes. 33 of 39 © Boardworks Ltd 2007
34 of 39 © Boardworks Ltd 2007
Glossary l activation energy – The amount of energy needed to start a reaction. l catalyst – A substance that increases the rate of a chemical reaction without being used up. l concentration – The number of molecules of a substance in a given volume. l enzyme – A biological catalyst. l rate of reaction – The change in the concentration over a certain period of time. 35 of 39 © Boardworks Ltd 2007
Anagrams 36 of 39 © Boardworks Ltd 2007
Rates of reaction: summary 37 of 39 © Boardworks Ltd 2007
Multiple-choice quiz 38 of 39 © Boardworks Ltd 2007
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). Chemical Kinetics
Outline: Kinetics Reaction Rates How we measure rates. Rate Laws How the rate depends on amounts of reactants. Integrated Rate Laws How to calc amount left or time to reach a given amount. Half-life How long it takes to react 50% of reactants. Arrhenius Equation How rate constant changes with T. Mechanisms Link between rate and molecular scale processes. Chemical Kinetics
Reaction Rates Rxn Movie Rates of reactions can be determined by monitoring the change in concentration of either reactants or products as a function of time. [A] vs t Chemical Kinetics
Reaction Rates C 4 H 9 Cl(aq) + H 2 O(l) C 4 H 9 OH(aq) + HCl(aq) [C 4 H 9 Cl] M In this reaction, the concentration of butyl chloride, C 4 H 9 Cl, was measured at various times, t. Chemical Kinetics
Reaction Rates C 4 H 9 Cl(aq) + H 2 O(l) C 4 H 9 OH(aq) + HCl(aq) Average Rate, M/s The average rate of the reaction over each interval is the change in concentration divided by the change in time: Chemical Kinetics
Reaction Rates C 4 H 9 Cl(aq) + H 2 O(l) C 4 H 9 OH(aq) + HCl(aq) • Note that the average rate decreases as the reaction proceeds. • This is because as the reaction goes forward, there are fewer collisions between reactant molecules. Chemical Kinetics
Reaction Rates C 4 H 9 Cl(aq) + H 2 O(l) C 4 H 9 OH(aq) + HCl(aq) • A plot of concentration vs. time for this reaction yields a curve like this. • The slope of a line tangent to the curve at any point is the instantaneous rate at that time. Chemical Kinetics
Reaction Rates C 4 H 9 Cl(aq) + H 2 O(l) C 4 H 9 OH(aq) + HCl(aq) • The reaction slows down with time because the concentration of the reactants decreases. Chemical Kinetics
Reaction Rates and Stoichiometry C 4 H 9 Cl(aq) + H 2 O(l) HCl(aq) • In this reaction, the ratio of C 4 H 9 Cl to C 4 H 9 OH is 1: 1. • Thus, the rate of disappearance of C 4 H 9 Cl is the same as the rate of appearance of C 4 H 9 OH. Rate = - [C 4 H 9 Cl] = t [C 4 H 9 OH] t C 4 H 9 OH(aq) + Chemical Kinetics
Reaction Rates and Stoichiometry • What if the ratio is not 1: 1? H 2(g) + I 2(g) 2 HI(g) • Only 1/2 HI is made for each H 2 used. Chemical Kinetics
Reaction Rates and Stoichiometry • To generalize, for the reaction a. A + b. B Reactants (decrease) c. C + d. D Products (increase) Chemical Kinetics
Concentration and Rate Each reaction has its own equation that gives its rate as a function of reactant concentrations. this is called its Rate Law To determine the rate law we measure the rate at different starting concentrations. Chemical Kinetics
Concentration and Rate Compare Experiments 1 and 2: when [NH 4+] doubles, the initial rate doubles. Chemical Kinetics
Concentration and Rate Likewise, compare Experiments 5 and 6: when [NO 2 -] doubles, the initial rate doubles. Chemical Kinetics
Concentration and Rate This equation is called the rate law, and k is the rate constant. Chemical Kinetics
Rate Laws • A rate law shows the relationship between the reaction rate and the concentrations of reactants. Ø For gas-phase reactants use PA instead of [A]. • k is a constant that has a specific value for each reaction. • The value of k is determined experimentally. “Constant” is relative herek is unique for each rxn k changes with T (section 14. 5) Chemical Kinetics
Rate Laws • Exponents tell the order of the reaction with respect to each reactant. • This reaction is First-order in [NH 4+] First-order in [NO 2−] • The overall reaction order can be found by adding the exponents on the reactants in the rate law. • This reaction is second-order overall. Chemical Kinetics
Integrated Rate Laws Consider a simple 1 st order rxn: A B Differential form: How much A is left after time t? Integrate: Chemical Kinetics
Integrated Rate Laws The integrated form of first order rate law: Can be rearranged to give: [A]0 is the initial concentration of A (t=0). [A]t is the concentration of A at some time, t, Chemical during the course of the reaction. Kinetics
Integrated Rate Laws Manipulating this equation produces… …which is in the form y = mx + b Chemical Kinetics
First-Order Processes If a reaction is first-order, a plot of ln [A]t vs. t will yield a straight line with a slope of -k. So, use graphs to determine rxn order. Chemical Kinetics
First-Order Processes Consider the process in which methyl isonitrile is converted to acetonitrile. CH 3 NC CH 3 CN How do we know this is a first order rxn? Chemical Kinetics
First-Order Processes CH 3 NC CH 3 CN This data was collected for this reaction at 198. 9°C. Does rate=k[CH 3 NC] for all time intervals? Chemical Kinetics
First-Order Processes • When ln P is plotted as a function of time, a straight line results. Ø The process is first-order. Ø k is the negative slope: 5. 1 10 -5 s-1. Chemical Kinetics
Second-Order Processes Similarly, integrating the rate law for a process that is second-order in reactant A: Rearrange, integrate: also in the form y = mx + b Chemical Kinetics
Second-Order Processes So if a process is second-order in A, a plot of 1/[A] vs. t will yield a straight line with a slope of k. First order: If a reaction is first-order, a plot of ln [A]t vs. t will yield a straight line with a slope of -k. Chemical Kinetics
Determining rxn order The decomposition of NO 2 at 300°C is described by the equation NO 2 (g) NO (g) + 1/2 O 2 (g) and yields these data: Time (s) 0. 0 50. 0 100. 0 [NO 2], M 0. 01000 0. 00787 0. 00649 200. 0 300. 00481 0. 00380 Chemical Kinetics
Determining rxn order Graphing ln [NO 2] vs. t yields: • The plot is not a straight line, so the process is not first-order in [A]. Time (s) 0. 0 50. 0 100. 0 [NO 2], M 0. 01000 0. 00787 0. 00649 ln [NO 2] -4. 610 -4. 845 -5. 038 200. 0 300. 00481 0. 00380 -5. 337 -5. 573 Does not fit: Chemical Kinetics
Second-Order Processes A graph of 1/[NO 2] vs. t gives this plot. Time (s) 0. 0 50. 0 100. 0 [NO 2], M 0. 01000 0. 00787 0. 00649 1/[NO 2] 100 127 154 200. 0 300. 00481 0. 00380 208 263 • This is a straight line. Therefore, the process is secondorder in [NO 2]. Chemical Kinetics
Half-Life • Half-life is defined as the time required for one-half of a reactant to react. • Because [A] at t 1/2 is one-half of the original [A], [A]t = 0. 5 [A]0. Chemical Kinetics
Half-Life For a first-order process, set [A]t=0. 5 [A]0 in integrated rate equation: NOTE: For a first-order process, the half-life does not depend on [A]0. Chemical Kinetics
Half-Life- 2 nd order For a second-order process, set [A]t=0. 5 [A]0 in 2 nd order equation. Chemical Kinetics
Outline: Kinetics First order Second order Rate Laws Integrate d Rate Laws Half-life complicated Chemical Kinetics
Temperature and Rate • Generally, as temperature increases, so does the reaction rate. • This is because k is temperature dependent. Chemical Kinetics
The Collision Model • In a chemical reaction, bonds are broken and new bonds are formed. • Molecules can only react if they collide with each other. Chemical Kinetics
The Collision Model Furthermore, molecules must collide with the correct orientation and with enough energy to cause bond breakage and formation. Chemical Kinetics
Activation Energy • In other words, there is a minimum amount of energy required for reaction: the activation energy, Ea. • Just as a ball cannot get over a hill if it does not roll up the hill with enough energy, a reaction cannot occur unless the molecules possess sufficient energy to get over the activation energy barrier. Chemical Kinetics
Reaction Coordinate Diagrams It is helpful to visualize energy changes throughout a process on a reaction coordinate diagram like this one for the rearrangement of methyl isonitrile. Chemical Kinetics
Reaction Coordinate Diagrams • It shows the energy of the reactants and products (and, therefore, E). • The high point on the diagram is the transition state. • The species present at the transition state is called the activated complex. • The energy gap between the reactants and the activated complex is the activation energy barrier. Chemical Kinetics
Maxwell–Boltzmann Distributions • Temperature is defined as a measure of the average kinetic energy of the molecules in a sample. • At any temperature there is a wide distribution of kinetic energies. Chemical Kinetics
Maxwell–Boltzmann Distributions • As the temperature increases, the curve flattens and broadens. • Thus at higher temperatures, a larger population of molecules has higher energy. Chemical Kinetics
Maxwell–Boltzmann Distributions • If the dotted line represents the activation energy, as the temperature increases, so does the fraction of molecules that can overcome the activation energy barrier. • As a result, the reaction rate increases. Chemical Kinetics
Maxwell–Boltzmann Distributions This fraction of molecules can be found through the expression: where R is the gas constant and T is the temperature in Kelvin. Chemical Kinetics
Arrhenius Equation Svante Arrhenius developed a mathematical relationship between k and Ea: where A is the frequency factor, a number that represents the likelihood that collisions would occur with the proper orientation for reaction. Chemical Kinetics
Arrhenius Equation Taking the natural logarithm of both sides, the equation becomes 1 RT y = mx + b When k is determined experimentally at several temperatures, Ea can be calculated from the slope of a plot of ln k vs. 1/T. Chemical Kinetics
Outline: Kinetics First order Second order Rate Laws Integrate d Rate Laws Half-life k(T) complicated Chemical Kinetics
Reaction Mechanisms The sequence of events that describes the actual process by which reactants become products is called the reaction mechanism. Chemical Kinetics
Reaction Mechanisms • Reactions may occur all at once or through several discrete steps. • Each of these processes is known as an elementary reaction or elementary process. Chemical Kinetics
Reaction Mechanisms • The molecularity of a process tells how many molecules are involved in the process. • The rate law for an elementary step is written directly from that step. Chemical Kinetics
Multistep Mechanisms • In a multistep process, one of the steps will be slower than all others. • The overall reaction cannot occur faster than this slowest, rate-determining step. Chemical Kinetics
Slow Initial Step NO 2 (g) + CO (g) NO (g) + CO 2 (g) • The rate law for this reaction is found experimentally to be Rate = k [NO 2]2 • CO is necessary for this reaction to occur, but the rate of the reaction does not depend on its concentration. • This suggests the reaction occurs in two steps. Chemical Kinetics
Slow Initial Step • A proposed mechanism for this reaction is Step 1: NO 2 + NO 2 NO 3 + NO (slow) Step 2: NO 3 + CO NO 2 + CO 2 (fast) • The NO 3 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. Chemical Kinetics
Fast Initial Step • The rate law for this reaction is found (experimentally) to be • Because termolecular (= trimolecular) processes are rare, this rate law suggests a two-step mechanism. Chemical Kinetics
Fast Initial Step • A proposed mechanism is Step 1 is an equilibriumit includes the forward and reverse reactions. Chemical Kinetics
Fast Initial Step • The rate of the overall reaction depends upon the rate of the slow step. • The rate law for that step would be • But how can we find [NOBr 2]? Chemical Kinetics
Fast Initial Step • NOBr 2 can react two ways: ØWith NO to form NOBr ØBy decomposition to reform NO and Br 2 • The reactants and products of the first step are in equilibrium with each other. • Therefore, Ratef = Rater Chemical Kinetics
Fast Initial Step • Because Ratef = Rater , k 1 [NO] [Br 2] = k− 1 [NOBr 2] Solving for [NOBr 2] gives us k 1 [NO] [Br ] = [NOBr ] 2 2 k− 1 Chemical Kinetics
Fast Initial Step Substituting this expression for [NOBr 2] in the rate law for the rate-determining step gives Chemical Kinetics
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. Chemical Kinetics
Catalysts One way a catalyst can speed up a reaction is by holding the reactants together and helping bonds to break. Chemical Kinetics
Enzymes • Enzymes are catalysts in biological systems. • The substrate fits into the active site of the enzyme much like a key fits into a lock. Chemical Kinetics
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