Higher Chemistry Controlling the Rate NEW LEARNING Collision Theory and Activation Energy Kinetic energy and temperature Reaction profiles Catalysts REVISION Reactions monitored and graphs interpreted Average Rate of Reaction Calculated from a Graph

Starter Questions S3 Revision

Starter Questions S3 Revision

Lesson 1: Collision Theory Today we will learn to Explain why reaction rate is increased by increasing temperature, concentration or surface area. We will do this by Using collision theory to explain some demonstrations and familiar reactions We will have succeeded if We can explain any rate change using collisions.

Activation Energy Activity 1.1: Activation Energy Your teacher may demonstrate some reactions that use heat or light to give the reactant particles the required activation energy. Watch the demonstration of bromine reacting with a cycloalkane. At National 5, we learned that alkenes rapidly decolourise bromine water. The reaction with saturated hydrocarbons is slower because light is needed to provide the activation energy. Heat is also sometimes the source of the activation energy.

Factors Affecting Reaction Rate Particle Size The smaller the particle size, the greater the surface area, and the faster the reaction Concentration The higher the concentration of the reactants, the faster the reaction. Temperature The higher the temperature, the faster the reaction These observations are explained by the collision theory.

Collision Theory In order to react particles must collide. A chemical reaction will only occur if the reacting particles collide with enough kinetic energy or speed. The energy is required to overcome the repulsive forces between the atoms and molecules and to start the breaking of bonds. The minimum kinetic energy required for a reaction to occur is called the activation energy (EA). When the reactant particles collide with the required activation energy they form an activated complex. This unstable intermediate breaks down to form the products of the reaction.

Collision Geometry and the Activated Complex E,g, The reaction of hydrogen and bromine H Br  2HBr Sometimes the collisions do not result in a reaction, despite having the minimum kinetic energy. This is thought to be because the particles have not collided with the correct geometry (angle) to allow the activated complex to be formed. In the above reaction of hydrogen and bromine the particles collided side on but if they collided end on… H-H + Br-Br  H----H-----Br----Br no reaction occurs as the activated complex cannot be formed if only 2 of the atoms come into contact with one another.

Surface Area/Particle Size Increasing Surface Area A reaction’s speed can be altered by increasing the surface area of one of the reactants. This can be achieved by breaking up lumps into chips or even better grinding it into a powder. When this occurs more of the reactant’s particles are exposed and therefore more particles can come in contact with the other reactant. This means that there will be a greater number of collisions, and therefore more successful collisions, resulting in a faster rate of reaction. Activity 1.2: After looking at the available chemicals and equipment and discussion with others in your class and your teacher, devise an experiment that would show clearly and fairly that increasing the surface area of a solid reactant speeds up the reaction.

Collision Theory and Particle Size The smaller the particle size, the faster the reaction as the total surface area is larger so more collisions will occur.

Starter Task Complete questions 1 and 2 in the tutorial booklet and we will review as a class. You have 10 minutes! S3 Revision

Starter Answers S3 Revision

Starter Answers S3 Revision

Lesson 2: Concentration Today we will learn to Investigate the effect of concentration on the rate of a reaction We will do this by Designing an experiment to monitor different concentrations but keep other variables the same. We will have succeeded if Alter concentration effectively to influence rate.

Concentration Increasing Concentration A reaction’s speed can be altered by increasing the concentration of a solution or pressure in a gas. When this occurs there are more reactant particles in a given volume of space and therefore more particles can come in contact with the other reactant. This means that there will be a greater number of collisions, and therefore more successful collisions, resulting in a faster rate of reaction.

Concentration Activity 1.3: Your teacher will show you what happens during the iodine clock reaction. After looking at the available chemicals and equipment and discussion with others in your class and your teacher, can you alter the concentration of the chemicals but not the total volume to get the clock to show exactly 1 minute.

Collision Theory and Concentration The straight line graph means rate is directly proportional to the concentrations of the reactants, i.e. double the concentration and you double the rate. This is true of many reactions. The faster rate is due to the increased number of collisions which must occur with higher concentrations of reactants.

Starter Task Complete questions 1 and 2 in Tutorial 3 booklet and we will review as a class. You have 10 minutes! S3 Revision

Starter Answers S3 Revision

Starter Answers S3 Revision

Lesson 3: Temperature Today we will learn to Investigate the effect of temperature on the rate of a reaction. We will do this by Designing an experiment to monitor different temperature but keep other variables the same. We will have succeeded if We can plot a graph of rate vs temperature to show the effect.

Temperature Increasing Temperature A reaction’s speed can be altered by increasing the temperature of the reaction mixture. When this occurs all the reactant particles have more energy and so a much greater number will have the required activation energy. This means that there will be a greater number of collisions with the activation energy, and therefore more successful collisions, resulting in a faster rate of reaction. Shaded areas represents number of particles which can collide successfully at each temperature. Activity 1.4: After looking at the available chemicals and equipment and discussion with others in your class and your teacher, devise an experiment that would show clearly and fairly that increasing the temperature speeds up the reaction.

Kinetic Energy and Temperature Temperature is a measure of the average kinetic energy or speed of the particles of a substance. At any given temperature, the particles of a substance will have a range of kinetic energies and this can be shown on an energy distribution graph. NB The maximum height of T2 is always lower than T1 E.g. 20oC E.g. 30oC

The graph above shows the kinetic energy distribution of the particles of a reactant at two different temperatures. It shows that at the higher temperature (T2), many more molecules have energies greater than the activation energy and will be able to react when they collide.

Explain why this student statement is wrong… Exit Task Mr Wrong Explain why this student statement is wrong… ‘Reaction rate is directly proportional to temperature, concentration and particle size’ Pupils could answer this, and if time, do some examples using the show me boards.

Starter Questions From your S2-S4 Coursework, describe what a catalyst is. A substance which speeds up a chemical reaction without being used up in the reaction. Give an example of a catalyst. Pt in a catalytic converter, Fe in the Haber process. Can the same catalyst be used in different chemical reactions? Sometimes. S3 Revision

Lesson 4/5: Catalysts Today we will learn to Explain how a catalyst speeds up a reaction. We will do this by Investigating different catalysts in the decomposition of H2O2 We will have succeeded if We can explain the effect of a catalyst on the energy changes in a reaction.

Catalysts A reaction’s speed can be altered by using a catalyst. A catalyst allows the reaction to take place via an alternative route that has a lower activation energy. This means that there will be a greater number of collisions with the activation energy at this temperature, and therefore more successful collisions, resulting in a faster rate of reaction. The enthalpy of the reaction stays the same.

Catalysts and Activation Energy A catalyst provides an alternative pathway for the reaction with a lower activation energy. (We will look at more of this type of energy diagram in the next week or so) ∆H

Catalysts and Reaction Rates A catalyst is a substance which changes the speed of a chemical reaction without being permanently changed itself. Catalysts speed up chemical reactions by providing an alternative reaction pathway which has a lower activation energy. There are 2 main types of catalyst: Homogeneous Catalysts Heterogeneous Catalysts

Homogeneous Catalysts Homogeneous catalysts are in the same state as the reactants. e.g. Heat Very little reaction occurs Fast reaction, solution turns green, gases evolve rapidly Solutions of potassium sodium tartrate and hydrogen peroxide (colourless) CoCl2(aq) + Reaction complete, solution turns pink again

Heterogeneous Catalysts Heterogeneous catalysts are in a different state to the reactants. e.g. Decomposition of hydrogen peroxide (solution) using manganese (IV) oxide (solid) as a catalyst. Manganese (IV) oxide (s) 2H2O2 (aq) 2H2O (l) + O2(g)

Common Catalysed Reactions Zymase Alcohol & Carbon dioxide Maltose & glucose Brewing Platinum Carbon Dioxide Carbon Monoxide Catalytic Converter Vanadium (V) oxide Sulphuric acid Sulphur Dioxide & oxygen Contact Aluminium oxide or silicate Short –chain hydrocarbons Long-chain Hydrocarbons Cracking Nitric acid Ammonia & Oxygen Ostwald Iron Ammonia Nitrogen & Hydrogen Haber Catalyst Products Reactants Process Homogeneous catalysis – hydroxide ions used in manufacture of soap from fats & oils.

Catalysts Activity 1.5: After looking at the available chemicals and equipment and discussion with others in your class and your teacher, devise an experiment that could identify the best catalyst for the decomposition of hydrogen peroxide.

How Heterogeneous Catalysts Work Active sites This type of catalyst is called a surface catalyst. It works by adsorbing the reacting molecules on to active sites and holding them with weak bonds on its surface. This not only causes the bonds within the molecule to weaken but also helps the collision geometry. The reaction occurs on the surface with less energy needed to form the activated complex (lower activation energy). The products are formed and leave the catalyst surface free for further reactions

Catalyst Poisoning A surface catalyst can be poisoned when another substance attaches itself to the ‘active sites’. This is very often irreversible so prevents reactant molecules from being adsorbed onto the surface. For this reason, catalysts have to be regenerated or renewed. E.g. Lead and its compounds are poisons of transition metal catalysts. This is why unleaded petrol must be used in cars with catalytic converters. Arsenic and its compounds are also common poisons. Catalysts can also be made ineffective by side-reactions. E.g. the iron catalyst used in the Haber Process rusts due to the presence of air and water, so needs to be replaced every so often.

Enzymes Enzymes catalyse the chemical reactions which take place in living cells. Enzymes are complex protein molecules which are very specific- they usually only speed up one particular reaction and work best at specific temperatures and pH (optimum).

Examples in nature are: Amylase – breaks down starch during digestion. Catalase – breaks down hydrogen peroxide Many enzymes are used in industry: Invertase – used in chocolate industry for the hydrolysis of sucrose to form fructose and maltose. Zymase - converts glucose into alcohol in the brewing industry. Protease (and others) – used in biological washing powders to dissolve natural stains like protein .

Objective Traffic Lights How do you feel about the lesson objectives? Red = don’t think I have grasped this Amber = feeling OK about this, have just about got there Green = Confident I have achieved this If there are no coloured cards available (e.g. in planners), Mark corners in the room for each colour and ask pupils to move to indicate confidence.