Rates of Reactions PDF

**[Rates of Reactions]** Chemists use the term of rate of reaction to describe how quickly chemical changes occur. \*\*Definition: The rate of reaction is defined as the change in concentration per unit time of any one reactant or product.\*\* **[Mandatory Experiment: To monitor the rate of produ...

**[Rates of Reactions]** Chemists use the term of rate of reaction to describe how quickly chemical changes occur. \*\*Definition: The rate of reaction is defined as the change in concentration per unit time of any one reactant or product.\*\* **[Mandatory Experiment: To monitor the rate of production of oxygen from hydrogen peroxide using manganese dioxide as a catalyst]** 1. Measure out 5 cm^3^ of hydrogen peroxide and dilute to 50 cm^3^ with water. Place it in the conical flask. 2. Weigh about 0.5 g manganese(IV) oxide into the small test tube, and use the thread and stopper to suspend the test tube in the conical flask. Avoid contact between the manganese(IV) oxide and the hydrogen peroxide. 3. Place sufficient water in the trough to allow the graduated cylinder to be filled with water and inverted over the beehive shelf. Using a teat pipette, inject air into the graduated cylinder until the water level is at the 10 cm^3^ mark. 4. Arrange the delivery tube for the oxygen produced to be collected in the graduated cylinder by displacement of water. 5. Loosen the stopper momentarily to allow the thread to fall into the flask and shake vigorously, thus bringing the manganese(IV) oxide into contact with the hydrogen peroxide. The stop-clock should be started as this contact is made. Record the total volume of gas in the graduated cylinder every 30 seconds. Readings should be taken at eye level. 6. Present the results in the following table: Time (mins) 0 0.5 1 1.5 2 2.5 3 3.5 4 --------------------- --- ----- --- ----- --- ----- --- ----- --- Volume O~2~ (cm^3)^ 7. Draw a graph of total volume of oxygen against time, putting time on the horizontal axis. **\*\*Definition: The instantaneous rate of reaction is the rate of reaction at any one particular time during the reaction.\*\*** In an experiment, 25 cm^3^ of 0.1 M sulfuric acid were added to excess granulated zinc in a reaction vessel. The hydrogen evolved in the reaction was collected and its volume was measured at regular intervals of time. The results obtained are shown in the following table Time (mins) 0 1 2 3 4 5 6 7 8 9 10 11 -------------------- --- ---- ---- ------ ------ ------ ---- ---- ------ ---- ---- ---- Volume of Hydrogen 0 30 45 52.5 56.3 58.2 60 61 61.5 62 62 62 A diagram of a chemical experiment Description automatically generated(1) \(ii)  The shape of the graph shows is that the reaction is fairly fast initially. As time proceeds, the reaction becomes slower and slower as indicated by the fact that the slope of the curve becomes less and less steep and eventually levels off. \(iii) To measure the instantaneous rate at 2 minutes, we must measure the slope of the tangent to the curve at the two minute mark on the graph. To measure the slope of the line, draw a tangent at the two minute mark on the graph and then draw any vertical like and any horizontal line. The slope of the line is the length of the vertical line divided by the length of the horizontal line. Instantaneous rate as 2 mins = Slope of line = 30/2.9 = 10.34cm^3^/min **[Factors affecting Rates of Reactions]** \- Nature of Reactants\ - Particle size\ - Concentration\ - Temperature\ - Catalysts **Nature of Reactants\ **- The rate of a reaction will depend on the bonds involved -- some bonds may be easier to break or form than others.\ - Ionic bonds are fast as there is no bond breaking or forming.\ - Covalent bonds are slower as bonds need to be broken before a reaction could take place. **Particle size\ **- Smaller particle size leads to faster reactions because the particles have a higher surface area and will collide more often\ - Powdered chemicals react more quickly than large pieces of the same substance.\ - Finely divided particles can lead to a dust explosion. A dust explosion involve the build up of very fine dust in the air leading to an explosive reaction taking place. A number of conditions are needed for a dust explosion to take place:\ (i) Particles must be combustible\ (ii) There must be a source of ignition\ (iii) The particles must be dry\ (iv) Oxygen must be present **Concentration\ **- Higher concentration of reactants leads to a faster rate of reaction\ - More particles mean a greater chance of collision Since the rate of reaction is defined as the change in concentration per unit time of any one reactant or product:\ Rate = [\$\\frac{Change\\ in\\ concentration\\ (moles/L}{Time\\ (secs)}\$]{.math.inline}*\ Rate =* [\$\\frac{1}{\\text{time}}\$]{.math.inline} **Temperature\ **- An increase in temperature means an increase in the rate of reaction\ - Increasing the temperature increases the number of collisions and the amount of energy in each collision\ - Exponential relationship **[Mandatory Experiment: To study the concentration on the reaction rate using sodium thiosulfate solution and hydrochloric acid]** 1. Place 100 cm^3^ of the sodium thiosulfate solution into a conical flask. 2. Add 10 cm^3^ of 3 M hydrochloric acid to the flask, while starting the stop clock at the same time. 3. Swirl the flask and place it on a piece of white paper marked with a cross 4. Record the time taken for the cross to disappear. 5. Repeat the experiment using 80, 60, 40 and 20 cm^3^ of. sodium thiosulfate solution respectively. In each case, add water to make the volume up to 100 cm^3^ and mix before adding HCl. 6. If the initial sodium thiosulfate concentration is 0.1 M, subsequent concentrations will be 0.08 M, 0.06 M, 0.04 M and 0.02 M respectively. 7. Record the results in a table similar to the following: Conc. Of Na~2~S~2~O~3~ (moles/L) --------------------------------------------------------------- -- -- -- -- -- -- Time (secs) Rate ([\$\\frac{1}{\\text{Time}}\$]{.math.inline}) (sec^-1^) 8. Draw a graph of 1/time against concentration. Using the data shown in the table, plot a graph of reaction rate (1/t) versus concentration of sodium thiosulfate. What conclusion can be drawn from the graph about the relationship between the rate of the reaction and the concentration of the sodium thiosulfate? Conc. Of Na~2~S~2~O~3~ (moles/L) 0.2 0.16 0.12 0.08 0.04 0.02 --------------------------------------------------------------- ------ ------ ------ ------ ------ ------- Time (secs) 1.14 1.43 1.89 2.94 5.88 11.11 Rate ([\$\\frac{1}{\\text{Time}}\$]{.math.inline}) (sec^-1^) 0.88 0.70 0.53 0.34 0.17 0.09 [\[CHART\]]{.chart} The straight line through the origin tells us that the rate of reaction is directly proportional to the concentration. **[Mandatory Experiment: To study the effect of temperature on reaction rate using sodium thiosulfate and hydrochloric acid]** 1. Place 100 cm^3^ of 0.05 M sodium thiosulfate solution into a conical flask. 2\. Warm the flask gently until the temperature is about 20 ^0^C. 3\. Add 5 cm^3^ of 3 M HCl, starting a stop clock at the same time, before proceeding. 4\. Without delay, swirl the flask, place it on a piece of white paper marked with a cross, and record the exact temperature of the contents of the flask. 5\. Record the time taken for the cross to disappear 6\. Repeat the experiment, heating the thiosulfate to temperatures of approximately 30 ^0^C, 40 ^0^C, 50 ^0^C and 60 ^0^C respectively (before adding the HCl). 7\. Record the results in a table similar to the following: Temperature (^o^C) --------------------------------------------------------------- -- -- -- -- -- -- Time (secs) Rate ([\$\\frac{1}{\\text{Time}}\$]{.math.inline}) (sec^-1^) 8. Draw a graph of 1/time against temperature. **Catalysts** \*\*Definition: A catalyst is a substance that alters the rate of a chemical reaction but is not consumed in the reaction.\*\* [General Properties of Catalysts]\ - Take part in chemical reactions and are recovered chemically unchanged at the end of the reaction\ - Catalysts are specific, each catalyst acts only on one type of reaction\ - Catalysts only need to be present in very small amounts to function. Increasing the amount of catalyst does not significantly alter the rate of the reaction\ - A catalyst helps equilibrium to be achieved more quickly\ - The action of a catalyst may be destroyed by catalyst poisons e.g., lead added to petrol destroyed the catalysts in the catalytic converter of cars [Types of Catalysts]\ \*\*Definition: Homogenous catalysis is catalysis in which both the reactants and the catalyst are in the same phase, i.e., there is no boundary between the reactants and the catalyst.\*\* For example, the iodine snake experiment. Concentrated potassium iodide solution is added to washing up liquid. Hydrogen peroxide is then added. Potassium iodide catalyses the reaction of hydrogen peroxide into water and oxygen. Both liquids are in the same phase. Oxygen is quickly released and carried foam from the washing up liquid out of the vessel in the form of a 'snake'. \*\*Definition: Heterogenous catalysis is catalysis in which the reactants and the catalyst are in different phases, i.e., there is a boundary between the reactants and the catalyst.\*\* For example, the oxidation of Methanol to Methanal. Methanol is heated and some of it vaporises. A red-hot platinum wire catalyst is inserted. Methanol vapours are oxidised to form methanal, hydrogen gas and carbon monoxide. Carbon monoxide temporarily poisons the catalyst, letting it cool. When the catalyst is no longer poisoned, the reaction begins again, heating the catalyst to become red hot. The cycle continues until the methanol is used up. The hydrogen produced gets ignited by the hot wire, producing "pops". \*\*Definition: Autocatalysis is catalysis in which one of the products of the reaction acts as a catalyst for the reaction.\*\* **[Mechanisms of Catalysts]** [The Intermediate Formation Theory] A catalyst works by forming an intermediate compound. For example, the decomposition of hydrogen peroxide catalysed by the presence of I^-^ ions shows the formation of an intermediate. \- One of the hydrogen peroxide molecules reacts with an I^-^ ion to form the IO^-^ intermediate.\ - The intermediate reacts with another hydrogen peroxide molecules to form the products and regenerate the catalyst [The Surface Adsorption Theory] Most heterogenous catalysis reactions can be explained using a theory called Surface Adsorption Theory.\ Adsorption is the accumulation of substances only at the surface of another substance. \- Adsorption: Molecules diffuse to the surface and settle there. The molecules are held at the surface by the formation of temporary weak bonds.\ - Reaction on surface: The higher concentration of molecules on the surface means that molecules will collide with each other. This results in bonds being broken and new bonds being formed. The reaction causes products to be formed.\ - Desorption stage: The products leave the surface of the catalyst. More reactants are then adsorbed onto the surface of the catalyst and the process starts again. **[Catalytic Converters]** \*\*Definition: A Catalytic converter is a device in the exhaust of a motor vehicle which contains catalysts to convert pollutants in the exhaust gases to less harmful substances.\*\* A catalytic converter consists of a thin coating of Platinum, Palladium and, Rhodium which act as the catalysts. Diagram of a car exhaust system Description automatically generated(2) Catalytic conversions: **Reactant** **Product** **Environmental benefit/removal** ------------------- -------------------------- ----------------------------------------------------------------- Carbon monoxide Carbon dioxide Carbon monoxide is a poisonous gas Nitrogen monoxide Nitrogen Nitrogen monoxide is poisonous and contributes to acid rain Hydrocarbons Carbon dioxide and water Hydrocarbons cause smog and contribute to the greenhouse effect Catalytic converters need to be changed after approximately 80,000 Km due to the poisoning of the catalysts. \*\*Definition: A catalyst poison is a substance that makes a catalyst inactive.\*\*\ Lead and sulphur are common examples of catalyst poisons. They adsorb onto the catalyst surface forming permanent bonds which prevent other reactions occurring. **[Collision Theory and Activation Theory]** \*\*Definition: An effective collision is one that results in the formation of products.\*\* \- For a reaction to occur, the reacting particle must collide with each other\ - A collision only results in the formation of products if a certain minimum energy is exceeded in the collision. This is called an effective collision.\ - When an effective collision occurs between colliding particles, bonds are broken and new bonds are formed. This results in the formation of products.\ - Collisions between particles can only form products when they have enough energy. \*\*Definition: The Activation Energy is the minimum energy that colliding particles must have for a reaction to occur, i.e., the minimum energy required for effective collision between particles to occur.\*\* The size of the Activation Energy depends on the nature of the reactants. **[Reaction profile diagrams]** \*\*Definition: A Reaction Profile Diagram is a graph which shows the change in energy of a chemical reaction with time as the reaction progresses.\*\* (3)\ A diagram of a curve Description automatically generated(4)\ A catalyst reduces the activation energy for a reaction **[Exam Questions]** [2014 -- HL -- Section A -- Question 3] An experiment to investigate the effect of temperature on the rate of the reaction between 0.05 M sodium thiosulfate solution and an excess of 3 M hydrochloric acid solution was carried out as follows. A timer was started as 5 cm^3^ of the acid were added to 100 cm^3^ of the sodium thiosulfate solution in a conical flask and a value was obtained for the time taken for the reaction to progress to a certain observable stage. The reciprocal of this time (1/time) was taken as an approximate measure of the initial rate of the reaction. This procedure was repeated at a number of different temperatures. The temperatures and their corresponding reaction times and rates are shown in the table below.\ (a) Explain the term rate of reaction. Change in concentration per unit time of one reactant or product.\ (b) (i) Describe and explain the change observed in the conical flask during the reaction. Precipitate observed due to the formation of sulfur.\ (ii) Describe how this observed change was used to obtain the reaction times.\ - Stand flask on cross\ - Note time on stop clock when cross becomes invisible  \(c) Plot a graph of reaction rate (1/time) versus temperature. [\[CHART\]]{.chart}\ (d) Describe and explain the relationship shown in your graph between rate of reaction and temperature. Rate increases exponentially with temperature.\ (e) Use your graph to find the value for the reaction time at 35 ºC. Give your answer correct to the nearest second.\ Rate from graph = 0.0145 sec^-1^\ 1/0.0145 = 69s\ (f) What would be the effect on the reaction times if the experiment were repeated using 0.025 M sodium thiosulfate solution? Justify your answer\ Reaction times double original values because the rate is directly proportional to thiosulfate concentration. [2012 -- HL -- Section B -- Question 7] \(a) Define the rate of a chemical reaction. Change in concentration per unit time of any one reactant or product\ (b) Explain clearly why there is an almost instantaneous reaction between aqueous solutions of sodium chloride and silver nitrate. They are ionic compounds\ (c) When hydrogen gas and nitrogen gas are mixed in a ratio of 3 : 1 by volume at room temperature in a sealed container, the formation of ammonia (NH~3~) is very slow. Suggest two ways to increase the rate of this reaction. Explain how each of the ways you suggest speeds up the reaction.\ Increased pressure:\ Closer molecules means more collisions per unit time.\ Increases energy of molecules.\ Increased temperature:\ More collisions reach activation.\ More effective collisions\ (d) Describe how you would measure the reaction time when 10 cm^3^ of 1.0 M hydrochloric acid solution and 50 cm^3^ of 0.20 M sodium thiosulfate solution react according to the following balanced equation:\ Na~2~S~2~O~3~ + 2HCl → 2NaCl + SO~2~ + S + H~2~O\ - Place the thiosulfate solution in a vessel over a cross (mark) on a white surface. - Add the HCI and start a stopclock.\ - Note the time when the cross becomes invisible\ If you were given additional sodium thiosulfate solutions of the following concentrations: 0.04 M, 0.08 M, 0.12 M and 0.16 M, describe how you would show that the rate of this reaction is directly proportional to the concentration of the sodium thiosulfate solution.\ - Repeat the above procedure for each of the given solutions.\ - Find the reciprocals of the times plot 1/t against concentration.\ - Plot of rate against concentration gives a straight line through the origin (0, 0)\ (e) Draw a reaction profile diagram for an exothermic reaction indicating clearly on your diagram (i) the activation energy (EA) for the reaction, (ii) the heat of reaction (∆H). **[References]** 1. Slideplayer.com 2. Wix.com 3. Docbrown.info 4. Docbrown.info