Chemistry 12 Unit I Reaction Kinetics Study Guide

Transcription

1 Chemistry 12 Unit I Reaction Kinetics Study Guide I.1 - Introduction: Reaction kinetics is the study of rates (speeds) of chemical reactions and the factors affect them. Rates of reaction are usually expressed in terms of a concentration change per unit of time. 1. The speed at which a reaction occurs is called the rate of reaction. We can express the rate of a reaction in terms of the amount of a reactant consumed or amount of a product formed per unit of time Reaction Rate = = or or, basically Example: = Δ amount Δ time 1. Which of the following could not be units for reaction rate? A. sec 1 B. g * ml -1 C. M/min D. o C/hour (where Δ means change in ) 2. If 16.0g of HCl are used up after 10 min in a certain reaction, calculate the average rate of the reaction in g HCl per second. Calculate the average consumption rate in g HCl / sec. 3. Consider the following reaction: (3 marks) C 12 H 22 O 11(s) 11H 2 O (g) + 12C (s) The rate of decomposition of C12H22O11 is 0.75mol/min. How many grams of C is produced in 10.0 seconds? (Ans: 18 g) 1

2 g/s of solid iron is reacted with excess gaseous Cl2 to produce Iron (III) chloride. Find the consumption rate of Cl2 in g/min. (Ans: 3.44 x 10 2 g Cl2/min) I.2 - Methods of Measuring Reaction Rates: ***Do Hebden Questions #1-6, pgs 2-3*** In order to calculate reaction rate, one must monitor observable changes over time. The observable change depends on the reaction itself. In order to illustrate the methods used to determine the rate of a reaction, we will consider a specific reaction, the reaction of copper metal with nitric acid in an open system. Cu(s) + 4HNO3(aq) Cu(NO3)2(aq) + 2H2O(l) + 2NO2(g) + HEAT red-brown colourless blue colourless brown As the reaction takes place, the amount of reactant will decrease whereas the amount of product will increase. To determine the rate of the above reaction, we could measure at least five different physical properties: a) Color changes The Cu(NO3)2 has a characteristic blue color. The intensity of the blue color could be measured in a SPECTROPHOTOMETER. Rate = Δ (color intensity)) Δ time b) Temperature changes Since the reaction is producing heat (exothermic reaction), we could perform the reaction in a CALORIMETER or with a thermometer and measure and graph the temperature changes. Rate = Δ(temperature changes) Δ time c) Pressure changes If the reaction were completed in a CLOSED system, the gas NO2 produced would cause pressure increase as it is the only gas species in the reaction. So we could measure the rate at which NO2 is produced by measuring the change in pressure. Rate = Δ (pressure) Δ time 2

3 d) Mass changes Since gaseous NO 2 is produced and released to the surrounding (OPEN system), we could measure the rate at which the gas is produced up by measuring its change in mass. Rate = Δ (mass) Δ time e) Acidity changes Since the reaction uses up nitric acid, HNO3, we can measure the change in the [H + ] with a ph meter. (Recall that [H + ] means molar concentration of H + in moles/litre.) Example: (April 2003) Rate = Δ [H + (aq)] Δ time ***Do Hebden Questions #7-9, pg 5*** I.3 - Factors That Affect Reaction Rates: The rate of a reaction depends on both the nature of the reactants and on the conditions under which the reaction occurs. In general there are six ways to speed up/ slow down chemical reactions. a) Temperature (As T increases, reaction rate increases) An increase in temperature will always increase the rate thus time required for the reaction decreases. Increasing the temperature increases the average kinetic energy of the particles. In general for a slow reaction, each increase in temperature of 10 C will double the reaction rate. b) Concentration of the Reactants (As [conc.] increases, reaction rate increases) The rates of both homogeneous and heterogeneous reactions are increased by an increase in concentration of reactants (gas and aqueous...not SOLIDS!). Increasing concentration increases the number of particles in the reaction and thus the probability of successful collisions. Increasing the pressure of a gaseous reaction increases the number of gaseous reactant particles and therefore reaction rate. c) Pressure of the Reactants (Gas) (As P increases, reaction rate increases) Pressure is just another way to define concentration for a gaseous species at a given temperature. Pressure can be increased by either adding reactant particles into a system or decreasing the volume (size) of the reaction vessel. 3

4 Examples: 10 mol = 2sss M 5L 10 mol = 2ss M 5L 20 mol = 4sss M 10 mol = 1ss0 M 5L A(g) + B(g) à 2C(g) A(g) + B(s) à 2C(l) d) The ature of Reactants Catalysts are NOT used up in a reaction. After the reaction is complete, there will be as much of the 4 1L pressure ds not affect reaction rate if pressure of A is increased, Some reactions are naturally slow because the bonds involved are very strong and unreactive, or the electrons involved are tightly held. Other reactions are naturally fast because the reaction involves the breakage of weak bonds or the removal of loosely held electrons. We have no control over these big differences in the rate of a reaction. Example: N 2 à N + N slow reaction as the reaction involves Na + + Cl - à NaCl faster reaction as reactants are already e) Surface Area, State, or Phase (As SA increases, reaction rate increases) (SA only affects heterogeneous reactions, not homogeneous reactions) The greater the surface area of the reactants, the greater the rate of the reaction. This is because the greater the surface area, the more reactant particles are exposed to each other which leads to better chance of collision. Example: powdered Zn reacts faster than chunk of Zn State and phase considerations The states (solid, liquid, gas) of reactants also have an effect on reaction rates. Homogeneous Reactions (A reaction involving one phase) Heterogeneous Reactions (A reaction involving multiple phases) Solid phase reactions tend to be slow because the reacting species cannot move freely. Liquid phase reactions tend to be fast because of the close proximity of the particles. Gas phase reactions tend to fast because of the speed of gas particles. Aqueous ion reactions have the fastest reaction rates because of their close proximity, their ability to move throughout the solution, and their strong positive- negative attraction. Aqueous >> Gases and Liquids >> Solids f) Catalysts and Inhibitors A catalyst is a chemical that can be added to a reaction to INCREASE the rate of the reaction.

5 catalyst as was originally put into the reaction. Catalysts provide an easier pathway to the products. Example: enzymes, platinum. An inhibitor is a chemical that REDUCES the reaction rate by combining with a catalyst or one of the reactants in such a way as to prevent the reaction from occurring. Examples of inhibitors are poisons and antibiotics. Examples: ***Do Hebden Questions #10-19, pgs 7-11*** I.4 Experimental Measurement of Reaction Rates Given a reaction A à B, plot the general shape of the curves showing [A] vs. time, and [B] vs. time on the axes below. 5

6 Initially, the rate of using up A (the reactant) will be very high. This makes sense, since initially, the [A] is pretty high and you learned that if the [reactant] is high, rate will be high. Later on, when a lot of A has been used up, the rate of the reaction will be slow. This is because low [reactant] gives slow reaction rate. The rate of a reaction can always be found by finding the slope of the amount vs. time graph. There are two types of rates: 1. Instantaneous rate (rate at a specific time) Draw a tangent line at a specific point by choosing the point for the specific time you are interested in. Calculate the slope of the tangent line. 2. Average rate (rate over a specific time interval) Draw a line connecting the points of interest. Calculate the slope of the line. Example: 1. What is the rate of reaction (g/s) for the Zn in the following reaction? Calculate without graphing. (Ans: 1 x 10 1 g/s) Zn (s) + 2HCl (aq) à ZnCl (aq) + H 2(g) Mass (g) Time (s)

7 I.5 - Reaction Rates and Collision Theory: Collision theory (or K.M.T) states that molecules act like small hard spheres that bounce off each other and transfer energy among themselves during their collisions. In order to react, molecules must collide with each other. For the collision to be considered effective, the particles must have: 1. Sufficient kinetic energy (speed) to collide, AND 2. Correct geometry hit at the right (proper) angle Kinetic Molecular Theory is used to explain why the following factors affect reaction rate. Effect of Temperature: If you increase the temperature, you increase the rate. This is because increasing temperature leads to: 1. Increase in the kinetic energy (speed) which 2. Increase the frequency of collisions Effect of Concentration: If you increase the concentration (reactants), you increase the rate. This is because increasing concentration leads to: 1. Increase in the number of particles in the system which 2. Increase the frequency of collisions Effect of Pressure: If you increase the pressure (gaseous reactants), you increase the rate. This is because increasing pressure (by decreasing volume) leads to: 1. Decrease in the space between particles in the system which 2. Increase the frequency of collisions Effect of Surface Area: If you increase the surface area (solid reactants), you increase the rate. This is because increasing surface area leads to: Examples: 1. Increase in the number of particles exposed to other reactants which 2. Increase the frequency of collisions ***Do Hebden Questions #20-22, pg 12*** 7

8 I.6 Enthalpy Changes in Chemical Reactions: The amount of energy required to break a bond between two atoms is bond energy. If we break a bond, an amount of energy equal to the bond energy must be added to the bond: Cl2(g) + 243kJ 2Cl(g) Conversely, if two atoms form a bond, an amount of energy equal to the bond energy is released by the atoms: 2Cl(g) Cl2(g) + 243kJ When a chemical reaction occurs, new molecules are formed as chemical bonds are broken and formed. Enthalpy (H) = total kinetic and potential energy which exists in the system at constant pressure. (Def) Enthalpy is the Heat of the reaction which often has a unit kj. ΔH = Delta H this is the change in enthalpy (heat). Its sign indicates whether: Energy has been added to the reactants in order to form products (Endothermic) or The reactants will give off their excess energy in order to allow the products to form (Exothermic) Exothermic: products have less energy than reactants (Heat is released). Consider the reaction: 4Fe + 3O2 2Fe2O3 4Fe + 3O2 ***ΔH is negative for exothermic reactions*** Enthalpy (H) ΔH is Negative _2Fe2O3 Reaction Proceeds There are three ways to incorporate the heat term into a chemical equation: Endothermic: products have more energy than reactants (Heat is added). Consider the reaction: 2N 2 + O2 2N 2 O +164kJ ***ΔH is positive for endothermic reactions*** 8

9 Examples: I.7 - Kinetic Energy Distributions: ***Do Hebden Questions #23-28, pgs 13-16*** Kinetic Energy = Energy due to movement (kj) Temperature = Average kinetic energy of the system ( o C) At a given temperature, not all molecules have the same KE. Instead, some are moving quickly and some are moving slowly. Thus, some molecules DO react and some DO NOT, so there MUST be CONTINUOUS DISTRIBUTION of energies among the molecules. (Label graph below using p. 18). The area underneath of the curve represents the number of particles. The location of the peak represents the temperature. 9

10 Note that some molecules have low KE and some have very high KE What happens when the temperature is increased Average KE increases = The peak moves to the right Maximum KE increases = The base gets wider Total % of particles remains the same = The area underneath the curve remains the same = The height gets shorter What happens to the rate? It goes up!!! That means there should be a bigger fraction of particles that have Minimum KE (See the graph below). Number of Molecules 25 C 100 C 200 C Minimum energy required before a molecule can react These molecules can react Kinetic Energy of Molecules The increased reaction rate due to an increase in temperature is primarily due to the increased number of molecules with sufficient energy to react, and not due to the increased number of collisions.ssssdddddsfdasuofjniasodfjmioa Rule of thumb: For a SLOW reaction, a 10 o C increase in temperature DOUBLES the rate. Why? Increasing T by 10 o C doubles the area under the curve past the Minimum KE line. What happens when concentration is increased Average KE remains the same = The peak remains the same Maximum KE remains the same = The base remains the same Total % of particles increases = The area underneath the curve increases = The height gets taller 10

11 What happens when surface is increased Average KE remains the same = The peak remains the same Maximum KE remains the same = The base remains the same Total % of particles remains the same = The area underneath the curve remains the same = The height remains the same What happens when catalyst is added Average KE remains the same = The peak remains the same Maximum KE remains the same = The base remains the same Total % of particles remains the same = The area underneath the curve remains the same = The height remains the same BUT, less KE is required. Kinetic energy of particles (kj) Catalysts lower the M.E, so more particles have enough KE to react. ***Do Hebden Questions #29-32, pgs 19-20*** 11

12 I.8 - Activation Energies: The existence of a minimum energy requirement before a molecule can react means the molecule has an energy barrier to overcome. This energy is called Activation Energy. Potential Energy Reactants The top of the hill. This is the point at which the actual reaction occurs. If reactant molecules do not posses sufficient energy to here, they bounce off each other and do not react. Products Reaction Proceeds As molecules approach each other, the outer electrons of each molecule repel the other, thus slowing down the molecules and converting their KE into PE. If the electrons can gain enough PE, the reactant bonds can be broken and the product bonds can be formed. At this critical moment a very short lived, high energy chemical species is formed the ACTIVATED COMPLEX. After the reaction, the newly formed product molecules repel each other (outer electrons) and the molecules start to move away. (Label graph below) Gain PE by slowing down Reactant molecules very briefly fuse together to form the ACTIVATED COMPLEX. Potential Energy Reactants Reactants approach each other. Product molecules lose PE by speeding up and converting it to KE. Products 12

13 Reaction Proceeds The activation energy exists in all reactions because of the mutual repulsion of the outer electrons as the reactants approach each other as well as the energy required to begin breaking bonds. The barrier gets bigger as a result of more electrons at the reaction site or when there are greater bond strength Activated Complex (*): The very short lived (unstable) molecule that forms at the moment of collision when the reactant molecules are in the process of rearranging to form products. The formula of an activated complex is found by simply adding up (connecting) all the atoms involved in the two reacting molecules. Since it is unstable, it is difficult to observe activated complex. Activation Energy (Ea) : The minimum potential energy required to change the reactants into the activated complex. A collision between a pair of particles is said to be effective if the collision results in a reaction occurring. We can now explain why a reaction proceeds at a fast rate, a slow rate, or even an effectively zero rate. The graph below shows that when two molecules collide, they need sufficient kinetic energy (KE) (that is, they have to be moving fast enough) for the reaction to occur. Potential Energy (PE) is gained by converting the kinetic energy. Number of Particles Reaction proceeding The KE diagram shows how many particles have the minimum AMOUNT OF KINETIC ENERGY to react. We can see in the graph above that less than 50% of the particles possess a KE greater than or equal to the minimum energy required for a reaction to occur. As a result this reaction would be relatively slow. For a fast rate, the PE hill would be lower and the a lot more molecules would possess the necessary M.E At the start of the reaction, most of the energy is stored as KE. As the particles (molecules) approach each other, they lose KE and gain PE. After the reaction, the molecules separate and the KE increases as the PE decreases. Assuming no energy is lost from the system, PE + KE = constant 13

14 Examples: ***Do Hebden Questions #33-40, pgs 23-24*** 14

15 We now introduce an extremely important concept prepare yourselves Many reactions can go FORWARD or BACKWARD That is: REACTANTS PRODUCTS The following Potential Energy diagram is the same as those we have seen before, however it shows that for all reversible reactions there is a Forward Activation Energy (Ea(f)) and a Reverse Activation Energy(Ea(r)) where: Ea(f) = the activation energy for the FORWARD reaction. Ea(r) = the activation energy for the REVERSE reaction *NOTE: Activation energy is always positive, that is energy must be added to get to the top of the hill.* Endothermic Reactions PE Ea(f) Ea(r) ΔH _ You can see from the diagram that: Ea(f) = Ea(r) + ΔH Reaction proceeds Exothermic Reactions _ Ea(f) In an exothermic rxn as well: PE Ea(r) Ea(f) = Ea(r) + ΔH Example: ΔH _ Reaction proceeds ***Do Hebden Questions #41-45, pg 25*** 15

16 I.9 - Reaction Mechanisms: Most reactions occur in more than one step or elementary step. Elementary Step: An individual step in a reaction mechanismshukshukshu. Reaction Mechanism: The actual sequence of steps that make up an overall reaction is a reaction mechanism.sdfhukuhsdfhuksdfhuksdfhuksdfhkj Consider the following reaction: 4HBr (g) + O 2(g) 2H 2 O (g) + 2Br 2(g) It is unlikely that 4 HBr molecules and 1 O2 molecule will collide simultaneously to undergo a chemical change. Instead, experiments show that the reaction occurs in more than one step. Complex reactions CANNOT go in a single step, but must consist of more than one step in getting from reactants to products. Consider the following reaction mechanism: Step 1 (slow): HBr (g) + O 2(g) à HOOBr (g) Step 2 (fast): HBr (g) + HOOBr (g) à 2HOBr (g) Step 3 (fast): HBr (g) + HOBr (g) à H 2 O (g) + Br 2(g) Step 1, 2, and 3 are called ELEMENTARY STEPS. The slowest step in the mechanism determines the overall rate of the reaction and is called the rate determining step or RDS. Unless RDS is altered to speed up, the overall rxn sill not be affected. In this example, step 1 is the RDS. In other reactions, the rate determining step can be a first, last, or middle step and still control the overall rate. To determine the overall reaction from the reaction mechanism, we simply add up all the steps in the reaction equations and cancel any species that occur on both sides of the final equation. Step 1 (slow): HBr (g) + O 2(g) à HOOBr (g) Step 2 (fast): HBr (g) + HOOBr (g) à 2HOBr (g) Step 3 (fast): HBr (g) + HOBr (g) à H 2 O (g) + Br 2(g) In the above example HOOBr and HOBr are reaction intermediates or simply intermediates. They are formed in one elementary process but subsequently, consumed in another and are therefore, not found written in the overall reaction. Unlike an activated complex, intermediates have lower PE and they are often observed. The formula of an activated complex is found by simply adding up all the atoms involved in the two reacting molecules. For example, in Step 2 (above) the activated complex is H2O2Br2 (the order of the atoms is not important). NOTE: YOU WILL NEVER HAVE TO PREDICT A REACTION MECHANISM. Years of precise work and complex analysis through experimentation are required to determine a mechanism. 16

17 Examples: *** Do Hebden Questions #46-53, pg 28*** 17

18 I.10 Potential Energy Diagram of a Reaction Mechanism: Each step ( elementary process ) of a reaction mechanism involves an individual activated complex and activation energy. For the reaction: 4HBr(g) + O2(g) 2H 2 O(g) + 2Br2(g) There are three steps and therefore three humps (Ea) in the PE Diagram. HOOBr Potential Energy (PE) HBr + O2 HOBr H2O + Br2 Reaction proceeds Each step of the mechanism has its own activation energy but the overall activation energy of a reaction does not equal the sum of the activation energies. Instead the overall Ea is the same as the Ea of the RDS (slowest step). The activation energy for a particular step in a reaction can be defined as: Ea = PE(activated complex) PE(reactants for the step) ***Do Hebden Questions #54-55, pg 30*** 18

19 I.11 - The Effect of Catalysts on Activation Energy: I.12 - The Effect of Catalysts on the Reaction Mechanism: A catalyst is a substance that provides an alternative reaction mechanism having lower Ea. Catalysts are not completely used up in chemical reactions but are regenerated. Catalyst work by orienting colliding particles or stabilizing the activated complex thereby lowering the Ea. If we lower the activation energy by adding a catalyst, more reactant molecules will possess the minimum kinetic energy required to form the activated complex. Therefore, more molecules can react and the rate increases. **Important Point: It is not only the forward reaction rate that is increased; by lowering the energy hump we also lower the reverse activation energy and increase the rate of the reverse reaction. Label the diagram Ea(f) catalyzed Potential Energy )cat Reactants Ea(r)cat Ea(r) Products Reaction proceeds For the catalyzed reaction, the value of H is the same as for the Uncatalyzed reaction. This is because catalyst has no effect on the PE of the reactant and product. Both catalysts and intermediates will NOT appear in the overall reaction BUT do not confuse them! Catalysts are used and then regenerated Intermediates are made then used up 19

20 Example: The reaction of hydrogen gas and oxygen gas to form water occurs slowly at room temperature since it has numerous steps. If a platinum catalyst is added, the reaction happens much faster with a two step mechanism. Pt catalyst 2H 2(g) + O 2(g) 2H 2 O (g) H H H H + H H H H O2 Pt H H H H H H Oxygen molecule sticks on the platinum which breaks the oxygen bond. A hydrogen molecule approaches. Hydrogen molecules collide with oxygen atoms to form water ***Do Hebden Questions #56-61, pg 34*** 20

21 I.13 - Some Uses of Catalysts: Catalysts are presently used extensively in industry, and are actively being researched for further uses. Hebden outlines multiple examples of catalysts in real life including the following: Most biological reactions are initiated or aided by catalysts called enzymes. The particular molecule upon which an enzymes acts is called its substrate. A particular reaction may involve several enzymes working together. Examples: a. The enzyme maltase breaks down the sugar maltose into glucose Maltase C 12 H 22 O 11 + H 2 O 2C 6 H 12 O 6 Maltose Glucose The industrial process for making sulphuric acid, H2SO4 involves several steps. First, sulphur is burned to give SO2: S (s) + O 2(g) SO 3(g) Next, the SO2 is passed over a catalyst consisting of finely divided particles of platinum (Pt) or vanadium pentoxide (V2O5) to form SO3: 2SO2(g) + O2(g) + catalyst SO3(g) + catalyst The resulting SO3 can then be added to water to make sulphuric acid: SO 3(g) + H 2 O (l) H 2 SO 4(l) Automobiles are equipped with a catalytic converter to reduce the amount of harmful pollutants that are released into the environment. The production of nitric oxide (NO), NO2, carbon monoxide and unburned hydrocarbons are the chief concern. The following two equations show how NO and NO2 ( NOx ) are made. N 2 + O 2 à 2NO and 2 NO + O2 à 2 NO2 ***Do Hebden Questions #61-63, pg 36*** 21