Chapter 7: Rate of Chemical Reactions PDF

# Chapter 7: Rate of Chemical Reactions ## Physical Changes - Undergoing a physical change means no new substances are made. - These changes are often easy to reverse and mixtures produced are usually relatively easy to separate. - Examples: - Fractional distillation of crude oil - Changing state (melting, freezing,...) - Dissolving ## Chemical Changes - In chemical reactions, new chemical products are formed that have very different properties to the reactants. - Most chemical reactions are impossible to reverse. - Energy changes also accompany chemical changes and energy can be given out (exothermic) or taken in (endothermic). - Examples: - Neutralization reactions to produce salts - Rusting - Combustion - Photosynthesis ## Chemical Change vs Physical Change | | Example | | | ------------- |:-------------:|:-------------:| | Combustion | Rotting | Melting | | Shredding | Rusting | Digestion | | Boiling | Chopping | | ## Collision Theory - It's a theory that explains how Reactions happen. - For a chemical reaction to occur, reactant particles need to collide with one another. - Not every collision results in the formation of products. For products to be formed, the collision has to have a certain minimum amount of energy associated with it. - Activation energy: This minimum amount of energy needed for a reaction to take place and produce successful collisions. - Collisions which result in the formation of products are known as successful collisions. - An image depicts a collision between an acid particle and a magnesium atom. The particles in the liquid move non-stop, and for the reaction to happen, bonds must break. - The first image shows a successful collision where enough energy was used to break bonds and form new ones to create magnesium chloride and hydrogen. - The second image shows an unsuccessful collision where the acid particle just bounced away without breaking any bonds. ## Collision theory states that for a chemical reaction to occur: 1. Particles must collide 2. Collisions must have enough energy which is called activation energy ## Rate of reaction - Rate is a measure of the change that happens in a single unit of time. - Rate of reaction: the rate at which reactants are used up (decrease in mass) or the rate at which products are formed (increase in mass). - We usually don't measure concentration directly, rather we measure something that changes by changing concentration. - Examples: - Mass - Volume of gas evolved - Colour intensity - Electrical conductivity per unit time. - A graph shows the concentration of reactants and products over time. The concentration of reactants decreases over time while the concentration of products increases over time. ## For reactions that produce a gas we can measure either: 1. **A decrease in mass using a digital balance.** - As gas escapes, the mass decreases. - An image shows a conical flask containing a liquid with a digital balance and cotton wool placed above the liquid. The cotton wool allows the gas to escape and prevents splashing of the liquid. 2. **The increase in volume of gas using a gas syringe** - An image shows a conical flask with a bung, filled with a liquid (hydrochloric acid) and a magnesium ribbon placed inside. A gas syringe is connected to the flask with a stopcock. Hydrogen gas produced will rise up the flask and into the gas syringe, pushing the plunger out. - The image further shows a picture of the gas syringe where the plunger is at the initial position and a picture where it has been pushed out by 20cm3 of gas. ### Steps for carrying out the experiment using a gas syringe: 1. Clean the magnesium with sandpaper. Put dilute hydrochloric acid in the flask. 2. Drop the magnesium into the flask, and insert the stopper and syringe immediately. 3. Start the clock at the same time. 4. Hydrogen begins to bubble off. It rises up the flask and into the gas syringe, pushing the plunger out. 5. The volume of gas in the syringe is noted at intervals – for example every half a minute — until the reaction is finished (constant reading at the gas syringe means no more gas is bubbled). ### Inverted measuring cylinder - Using a flask with a delivery tube and collecting the gas over water in an inverted measuring cylinder or burette. - An image shows the setup with the flask filled with an acidic solution and a magnesium ribbon inside. The flask is connected to a delivery tube that goes into an inverted measuring cylinder filled with water. There is cotton thread placed in the flask above the magnesium ribbon and the hydrogen gas will collect at the top of the measuring cylinder. - A table shows the time in minutes and the volume of hydrogen gas collected in cm^3. - You can see that the reaction lasted about five minutes because after 5 minutes, the volume of gas is constant at 40 cm^3. - A graph shows the volume of hydrogen gas collected in cm^3 against time in minutes. The graph shows the rate of reaction increases quickly in the first few minutes, then starts to slow down until it reaches a constant point at 40 cm^3. ### Explanation of the results: 1. In the first minute, 14 cm^3 of hydrogen are produced. So the rate for the first minute is 14 cm^3 of hydrogen per minute. In the second minute, only 11 cm^3 are produced. - So the rate for the second minute is 11 cm^3 of hydrogen per minute. - The rate for the third minute is 8 cm^3 of hydrogen per minute. - So the rate decreases as time goes on. - The rate changes all throughout the reaction. - It is greatest at the start but decreases as the reaction proceeds. 2. The reaction is fastest in the first minute, and the curve is steepest then. - It gets less steep as the reaction gets slower. - The faster the reaction, the steeper the curve. 3. After 5 minutes, no more hydrogen is produced, so the volume no longer changes. - The reaction is over, and the curve goes flat. - When the reaction is over, the curve goes flat. 4. Altogether, 40 cm^3 of hydrogen are produced in 5 minutes. Therefore, the average rate for the reaction is 40/5 = 8 cm^3 of hydrogen per minute. - A graph shows the rate of reaction changes over time with the curve steepest at the beginning where the reaction is fastest, less steep as it slows down, and flat when the reaction is over. ## Factors Affecting Rate of Reaction: - The four major influences on reaction rate: 1. Concentration 2. Temperature 3. Surface area 4. Catalyst 5. Pressure of gas ### 1. Concentration: - An increase in the concentration of a solution, the rate of reaction will increase. - This is because there will be more reactant particles per unit volume, allowing more frequent and successful collisions per second, increasing the rate of reaction. - The image shows two beakers containing hydrochloric acid and magnesium, where the beaker B has an acid concentration twice as high as the beaker A. - Beaker A shows a dilute acid with less number of acid particles, hence a less chance of a successful collision. - Beaker B shows a concentrated acid with more number of acid particles, hence a higher chance of successful collisions. - If we repeated the experiment between hydrochloric acid, HCl, and magnesium, Mg — keeping everything the same each time except the concentration of the acid — in B, it is twice as concentrated as in A. - Here are both sets of results shown on the same graph: - A graph shows two curves plotted on the same graph, where the volume of hydrogen gas is plotted against time. The curve B is steeper than the curve A, showing the rate of reaction is faster in B. - The reaction in B lasts for 60 seconds, while in A it lasts for 120 seconds. - Both reactions produce 60 cm^3 of hydrogen. - So in B the average rate was 1 cm^3 of hydrogen per second, and in A, it was 0.5 cm^3 of hydrogen per second. - The average rate in B was twice the average rate in A. - So in this example, doubling the concentration doubled the rate. ### Conclusion: Increasing the concentration of reactant increases the rate of the reaction ### 2. Temperature - An increase in temperature, the rate of reaction will increase. - This is because the particles will have more kinetic energy and move faster than the required activation energy, therefore there will be more frequent and successful collisions per second, increasing the rate of reaction. - An image shows a representation of a reaction taking place between an acid particle and a magnesium atom. The acid particle is moving faster than the other particles, leading to more frequent and successful collisions. - Another image shows a graph where the volume of hydrogen gas is plotted against time. The graph shows that the volume of hydrogen gas produced increases at a higher rate at a higher temperature. - Eg: Dilute hydrochloric acid and sodium thiosulfate solution react to give a fine yellow precipitate of sulfur. ### The method: 1. Mark a cross on a piece of paper. 2. Place a beaker containing sodium thiosulfate solution on top of the paper, so that you can see the cross through it, from above. 3. Quickly add hydrochloric acid, start a clock at the same time, and measure the temperature of the mixture. 4. The cross grows fainter as the precipitate forms. Stop the clock the moment you can no longer see the cross. record the time. 5. Now repeat steps 1 - 4 several times, changing only the temperature. You do this by heating the sodium thiosulfate solution to different temperatures, before the reaction takes place. - The cross appears fainter with time. ### Results: - The cross disappears when enough sulfur has formed to hide it. (the shorter the time taken for the cross to disappear, the faster the reaction has taken place). - This took 200 seconds at 20°C, but only 50 seconds at 40°C. So the reaction is four times faster at 40°C than at 20°C. - A reaction goes faster when the temperature is raised. - A table is shown where the time for the cross to disappear in seconds against the temperature in °C is tabulated showing the rate of reaction doubles when the temperature increases by 10 °C. ### 3. Surface Area by decrease in size - An increase in the surface area of the solid, the rate of reaction will increase. - This is because more surface area particles will be exposed to the other reactant, so there will be more frequent and successful collisions per second, increasing the rate of reaction. - An image shows a cube divided into smaller cubes, each having a different surface area and volume. - When surface area is 6 cm^2, volume is 1 cm^3, and the ratio of surface area to volume is 6:1. - When surface area is 24 cm^2, volume is 8 cm^3, and the ratio of surface area to volume is 3:1. - When surface area is 54 cm^2, volume is 27 cm^3, and the ratio of surface area to volume is 2:1. - This shows that when the surface area of the solid increases, the ratio of surface area to volume decreases. - Eg: The reaction between hydrochloric acid and calcium carbonate (marble chips) Carbon dioxide gas is produced: - The image shows an experimental setup where marble chips are placed in a conical flask with hydrochloric acid inside and covered with cotton wool. The reaction produces carbon dioxide gas which will escape from the flask through the cotton wool. ### The method: 1. Place the marble in the flask and add the acid. 2. Quickly plug the flask with cotton wool to stop any liquid splashing out. 3. Then weigh it, starting the clock at the same time. 4. Note the mass at regular intervals until the reaction is complete. 5. Carbon dioxide is a heavy gas. It escapes through the cotton wool, which means that the flask gets lighter as the reaction proceeds. 6. By weighing the flask at regular intervals, you can follow the rate of reaction. 7. The experiment is repeated twice. Everything is kept exactly the same each time, except the surface area of the marble chips. - For experiment 1, large chips are used. Their surface area is the total area of exposed surface. - For experiment 2, the same mass of marble is used — but the chips are small, so the surface area is greater. - A graph shows two curves plotted on the same graph, where the loss of mass is plotted against time. The curve 2 is steeper than the curve 1 showing the rate of reaction is faster for smaller chips. The final loss of mass is 2 grams in both experiments. - For the small chips, the reaction is complete in 4 minutes while for the large chips, it takes 6 minutes. - The rate of a reaction increases when the surface area of a solid reactant is increased. - The mass of final product is always the same because we are measuring the rate (how fast a reaction is). - But we control using the same amount of reactant to produce the same amount of products. ### 4. Catalyst: - A substance that increases the rate of reaction without being chemically changed by lowering activation energy. - Catalysts reduce the activation energy by creating an alternative pathway, requiring lower activation energy, allowing more successful collision without being used up. - This shows that when a catalyst is used, the rate of reaction will increase. - An image shows the energy changes for an uncatalysed reaction and a catalysed reaction. The catalysed reaction has a lower activation energy due to the catalyst creating an alternative pathway for the reaction to proceed. - Eg: The decomposition of hydrogen peroxide. - An image shows three conical flasks, each containing hydrogen peroxide solution with either nothing, manganese (IV) oxide, or raw liver added. A glowing splint is used to test for the presence of oxygen, which indicates the rate of decomposition of hydrogen peroxide. - The first flask is the control where the splint does not relight since the decomposition is slow. - The second flask containing manganese (IV) oxide shows the splint relights indicating a faster decomposition. - The third flask containing raw liver also shows a re-lit splint indicating a faster decomposition. - So manganese (IV) oxide acts as a catalyst for the reaction. - If you add more manganese (IV) oxide, the reaction will go even faster. - Much catalyst work by providing a surface on which other molecules or atoms can react. - Others work in a more complex way. - Thus, it is wrong to say that catalysts do not take part in the reaction, some do. - But the end of the reaction, there is the same amount of catalyst as at the beginning, and it is chemically unchanged. ### 5. Pressure: - In gases only. - By increasing the pressure, same number particles are available in smaller volume, collide harder, number of successful collisions increase so the rate of reaction increases. - An image shows two containers with gas particles inside. The container on the left is under low pressure with the gas particles moving slowly, while the container on the right is under high pressure with the gas particles moving faster. - Describe the application of the above factors to the danger of explosive combustion with fine powders (e.g. flour mills) and gases (e.g. methane in mines): - **Flour mills:** - Particle size is very small. - Therefore, the surface area is very large. - Could easily combust causing an explosion due to these flammable substances that have a large surface area. - **Methane in mines:** - Increase in pressure. - Same as increasing the concentration of the reactants — because now the volume has decreased, therefore there are more particles per unit volume. - Increases the chance of successful collisions. ## Reversible reactions - Some reactions go to completion, where the reactants are used up to form the product molecules and the reaction stops when all of the reactants are used up. - Eg: calcium carbonate reacts with hydrochloric acid; reaction stops when all the calcium carbonate is used up. - In reversible reactions, the product molecules themselves react with each other or decompose and form the reactant molecules again. - We use the symbol “≒” instead of a single arrow, to show that a reaction is reversible. - When blue hydrated salt of copper sulfate is heated, it decomposes into white anhydrous copper sulfate and water. - Eg: CuSO4.5H2O ≒ CuSO4 + 5H2O ### Hydrated salts: - Are salts that contain water of crystallization which affects their molecular shape and colour. ### Anhydrous salts: - Are those that have lost their water of crystallization, usually by heating. in which the salt becomes dehydrated. - It is said that the reaction can occur in both directions: the forward reaction (which forms the products) and the reverse direction (which forms the reactants). - A reversible reaction is endothermic in one direction, and exothermic in the other, and the same amount of energy is transferred each time. - An image shows a reversible reaction of the decomposition of copper sulfate into copper sulfate and water. The forward reaction of decomposition is endothermic, while the reverse reaction of adding water is an exothermic reaction. ### Here are some important reversible reactions: | Reaction | Comments | | :----------------- | :-------------------------------------------------------------------------------------------------------------------------- | | N2(g) + 3H2 ≒ 2NH3(g) | This is a very important reaction, because ammonia is used to make nitric acid and fertilizers. | | 2SO2(g) + O2 ≒ 2SO3(g) | This is a key step in the manufacture of sulfuric acid. | | Ca SO4.5H2O | A blue hydrated crystal of CaSO4.5H2O decomposes into a white anhydrous powder (CuSO4) and water (5H2O). | | CoCl2.6H2O | A pink hydrated crystal of CoCl2.6H2O decomposes into a blue anhydrous powder (CoCl2) and water (6H2O). | ## Reactions that are reversible can reach to an equilibrium state. ## Equilibrium State: - Rate of Forward= Backward Reaction - Concentration of Reactant and Product stays constant - In a closed container - A graph shows the rate of reaction against time. The rate of the forward reaction decreases while the rate of the backward reaction increases until equilibrium is reached where both rates become equal. ### Note - The catalyst doesn't affect the position of equilibrium; however, it speeds up both the forward and backward reaction so the reaction reaches the equilibrium faster. ### Equilibrium position - Refers to the relationship between the concentration of reactants and products at the equilibrium state. - When the position of equilibrium shifts to the left, it means the concentration of the reactant increases. - When the position of equilibrium shifts to the right, this means the concentration of products increases. ### Le Chatelier's Principle: - States that when a change is made to the conditions of a system at equilibrium, the system automatically moves to oppose the change. ### There are three factors that affect equilibrium position: 1. Temperature 2. Pressure 3. Concentration ### 1. Temperature: | Change | How the equilibrium shifts | | :-------------------------- | :------------------------------------------------------------------------------------------------------------ | | Increase in temperature | Equilibrium moves in the ENDOTHERMIC direction to reverse the change. | | Decrease in temperature | Equilibrium moves in the EXOTHERMIC direction to reverse the change. | ### Example: - Iodine monochloride reacts reversibly with Chlorine to form Iodine trichloride. - ICI + Cl2 ≒ ICI3 - When the equilibrium mixture is heated, it becomes dark brown in color. - Explain whether the backward reaction is exothermic or endothermic. - Equilibrium has shifted to the left as the color dark brown means that more of ICI3 is produced. - Increasing temperature moves the equilibrium in the endothermic direction. - So the backward reaction is endothermic. ### 2. Pressure (In gases only) | Change | How the equilibrium shifts | | :---------------- | :------------------------------------------------------------------------------------------------------------------------------------------------ | | Increase in pressure | Equilibrium shifts in the direction that produces the smaller number of gas moles to decrease the pressure again. | | Decrease in pressure | Equilibrium shifts in the direction that produces the larger number of gas moles to increase pressure again. | ### Example - Nitrogen Dioxide can form Dinitrogen Tetroxide. - 2NO2 ≒ N2O4 - Explain the effect of an increase in pressure on the position of equilibrium: - The number of molecules of gas on the left is 2, and the number of molecules of the gas on the right is 1. - An increase in pressure will cause equilibrium to shift in the direction that produces the smaller number of molecules of gas. - So equilibrium shifts to the right. ### 3. Concentration | Change | How the equilibrium shifts | | :-------------------- | :---------------------------------------------------------------------------------------------------------------------------------------------------------------------- | | Increase in concentration | Equilibrium shifts to the Right to reduce the effect of an increase in the concentration of a reactant. | | Decrease in concentration | Equilibrium shifts to the Left to reduce the effect of a decrease in reactant (or an increase in the concentration of the product). | ### Example: - Iodine Monochloride reacts reversibly with Chlorine to form Iodine Trichloride. - ICI + Cl2 ≒ ICI3 - Predict the effect of an increase in concentration on the position of equilibrium: - An increase in the concentration of ICI or Cl2 causes the equilibrium to shift to the right so that more of the yellow product is formed. - A decrease in the concentration of ICI or Cl2 causes the equilibrium to shift to the left so that more of the dark brown reactant is formed. ## Haber process - It is the process of ammonia production. - 3H2(g) + N2(g) ≒ 2NH3(g) ### Conditions 1. **Temperature = 450°C** - 450°C is the (optimum temperature/compromise temperature). - Below this temperature, the rate will be very slow. - Above this temperature, the reaction will go backwards and less ammonia will be produced. 2. **Pressure = 200 atm or 20000 kPa** - The pressure is applied to the equilibrium, and the system favors the side with less number of moles of gases. - There is less number of moles on the right side (where ammonia is produced), so increasing the pressure will shift the equilibrium in the forward direction. - This also increases the rate. 3. **Catalyst = powder iron** - Disadvantages of using high pressure: - Very expensive - Dangerous - A graph shows the yield of ammonia against pressure at different temperatures. The yield of ammonia increases with an increase in pressure until it reaches a maximum point and then starts to decrease. The yield at lower temperatures is higher than at higher temperatures. ### Rates of chemical reactions - **Collision theory:** - Particles must collide, with enough energy to produce a successful collision. - A successful collision is a collision that has enough energy to produce a product. - **Activation energy:** - The minimum amount of energy needed to start a reaction and produce a successful collision. - **Exothermic:** - A reaction that releases heat to the surroundings. - **Endothermic:** - A reaction that absorbs heat from the surroundings. - An image shows a diagram to show energy changes involved in exothermic and endothermic reactions. - The exothermic reaction has a negative enthalpy change, and the activation energy is less than the enthalpy change. - The endothermic reaction has a positive enthalpy change and the activation energy is less than the enthalpy change. ### Rates of chemical reactions - **Rate:** - The change in concentration between reactants and products over time. - The slope of the graph is equal to the rate. - **Factors that can be used to measure the rate of reaction:** - pH - Color - Mass - Volume - An image shows a diagram representing the change in the amount of reactants and products over time. The graph shows the amount of reactants decreasing over time while the amount of products increasing over time. - **Product is gas** - Measure the change in volume collected in a gas syringe. - Measure the decrease in mass. - An image shows a diagram representing the setup for measuring the rate of reaction. It shows a flask with a solution with a bung on top, and a gas syringe is connected to the flask. There is also a balance underneath the flask. ### Factors Affecting Rate of Reaction 1. **Temperature:** - K.E. of particles increases. - Frequency of successful collisions increases. - Rate increases. - More particles have higher energy than activation energy, leading to an exponential increase in the rate of reaction. - A graph shows the rate of reaction against temperature. The rate of reaction increases exponentially with an increase in Temperature. 2. **Concentration:** - No. of particles per unit volume increases. - Frequency of successful collisions increases. - So no. of collisions increases and the rate increases. - A graph shows the rate of reaction against concentration. The rate of reaction increases linearly with an increase in concentration. 3. **Surface area**: - As the size of the particle increases, the surface area increases. - Frequency of successful collisions increases. - Rate increases - A graph shows the Rate of reaction against Surface area. The rate of reaction increases linearly with an increase in surface area. 4. **Catalyst**: - A substance that speeds up the rate of reaction without being chemically changed. - They do this by lowering the activation energy. - Mostly from transition elements and compounds. - A graph shows the rate of reaction against time. The rate of reaction is higher with catalyst compared to the reaction without the catalyst. It shows activation energy is lower with a catalyst, meaning the reaction can go to completion at a faster rate. 5. **Pressure (more collisions)**: - Increases the collisions, resulting in an increase in the rate of reaction. - A graph shows the rate of reaction against pressure. The rate of reaction increases exponentially as pressure increases. ### Reactions - **Irreversible:** - C + O2 ➝ CO2 - Reactions go to completion. - Exo: increasing Temperature increases the rate of reaction. - Endo: Decreasing Temperature decreases the rate of reaction. - **Reversible:** - Chemical test for water using CuSO4 and CoCl2. - 5H2O + CuSO4 ≒ CuSO4 + 5H2O - 6H2O + CoCl2 ≒ CoCl2 + 6H2O - Haber process (making ammonia) - N2 + 3H2 ≒ 2 NH3 ### Define or state features of dynamic equilibrium: - Rate of forward reaction = rate of backward reaction. - Conc. of reactant and product is constant, not the same. - In a closed container. ### Factors affecting dynamic equilibrium: 1. **Conc:** - If the concertation of A and B increases, forward reaction increases. - If the concertation of C increases, backward reaction increases. 2. **Temp:** - If temperature increases, equilibrium shifts to the endothermic direction (Backward reaction). - If temperature decreases, equilibrium shifts to the exothermic direction (Forward reaction). 3. **Pressure (gases only):** - It is the collision of particles with each other and walls of the container. - If pressure increases, equilibrium will shift to the side with less number of moles of gases. - If pressure decreases, equilibrium will shift to the side with more number of moles of gases. ### Haber process - N2(g) + 3H2(g) ≒ 2NH3(g) - Main goal = forward > Backward - 4 moles - 2 moles - **Catalyst:** - Iron powder - **Temp** - Decreasing Temperature to increase Exothermic direction (Forward reaction). - Compromise temperature = 450°C - **Pressure:** - Increase the pressure in A + B to shift to the forward reaction. - 20,000 kPa/300 atm.

Chapter 7: Rate of Chemical Reactions PDF

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