Chemistry Lab (DKK 1711) 1 EXPERIMENT 1 PART 1: INTRODUCTION TO THE LABORATORY AND USAGE OF LABOTORY EQUIPMENT. Objectives: 1. To understand chemistry laboratory safety rules and regulation and abide by them. 2. To know the correct techniques of handling laboratory apparatus. Introduction: The chemistry laboratory should be a safe place to work in. Because of this, you should know all laboratory rules and regulations, including the correct way of using lab apparatus and handling of chemicals. Before any experiment work can begin in the laboratory, the rules and regulations must be understood and agreement be made to abide by these rules. Laboratory Rules and Regulation: 1. Attendance Attendance is COMPULSORY. All students are required to attend practical classes of 3 hours each. They are required to complete 10 experiments each semester. If a student is unable to attend any practical classes, he/she should produce a medical certificate or letter of exemption. Every student should: i) Have a jotter/log book. ii) Wear a lab coat, proper shoes (no sandals or slipper) and safety goggles (when needed) Chemistry Lab (DKK 1711) 2 2. Apparatus: SAFETY IN THE LAB IS PRIORITY. Every student must check the condition of all the apparatus to be used before starting the experiments, if there is a shortage of apparatus or breakage, please report it to the lecturer immediately. Every student is required to use and handle apparatus with care. Apparatus and the work area must be cleaned after completing experiments. Check all glassware for cracks before using it. Report to your lecture in case of any breakage or malfunction of apparatus during the experiment. 3. Chemicals: All volatile or dangerous chemicals such as concentrated acids are usually placed in the fume cupboards. Chemical that are less dangerous are usually placed on shelves or on the table. However, it is wise to treat every chemical as if it were hazardous. Be careful not to contaminate the chemicals. To avoid contamination, NEVER put your pipette into the reagent bottle and NEVER return unused chemicals to their respective bottles. When pouring out reagents, hold the stopper in your hand. Do not put it on the table. When replacing the stopper, place it first at the opening to ensure that any drips do not spill outside the reagent bottle. Take only sufficient amounts of chemicals for your experiments and use them with care. Share any excess. Do not waste chemicals. Dispose of excess chemicals in appropriate waste container. 1. Handling Apparatus: Cleaning lab apparatus In order for experiments to run smoothly and yield accurate results, all apparatus used must be clean. Apparatus should be washed with a soap solution and brushed with a Chemistry Lab (DKK 1711) 3 suitable brush. Rinse with tap water and then with distilled water. Clean apparatus does not have traces of fat or oil on the surface. Using some common lab apparatus, lecturer would observe and demo to their student how to handle apparatus such as; 1. Graduated cylinder 2. Pipette 3. Volumetric flask 4. Burette 5. Bunsen Burner 6. Analytical balance. Exercise 1- Using the analytical balance 1) Take any object e.g. beaker, porcelain dish, watch etc and measure the mass of the object 5 times. (You can weigh to the closest milligram on the balance provided) 2) Record all readings in your jotter book ( State the object that you selected) 3) Calculate the average mass of each object. Determine the deviation in the average mass of the object. 4) Compare the maximum deviation with the average deviation, Explain. PART 2: DETERMINATION OF THE FORMULA UNIT OF A COMPOUND Objective 1. To synthesis a zinc chloride compound 2. To determine the formula unit of zinc chloride Introduction One of the main properties of a compound is its chemical composition, which can be identified simply by determining the elements presents in the compound. Although Chemistry Lab (DKK 1711) 4 qualitative analysis is very useful, it would be an added advantage to chemists if they know the quantitative aspect of the compound. The outcome from a quantitative analysis can be used to determine the composition of unknown compound. Once the composition of the compound is known, its formula unit can be determined. For example, a compound containing 0.1 mole Ag and 0.1 mole Br will have formula unit of AgBr. In this experiment, you will prepare a simple compound which is composed of two elements, zinc and chlorine. Once the mass of zinc and the mass of the compound are known, the mass of chlorine can be determined. Using these masses, the percentage composition of the product can be determined and the formula unit can be calculated. Material Glass rod, crucible, wire gauze, tripod stand, porcelain tiles, analytical balance, measuring cylinder 10 mL, bunsen burner, HCl 6.0M and zinc powder Method 1. Weigh the crucible and record the exact mass 2. Place approximately 0.25 g Zinc powder into the crucible. Weigh the crucible with its contents and determine the exact mass of zinc. 3. Carefully add in 10ml of 6.0M HCl solution into the crucible containing the zinc powder and stir the content gently using a glass rod. A vigorous chemical reaction will occur and hydrogen gas will be released. (Caution: Carry out this step in a fume cupboard. Do not work near a fire source. Wet hydrogen gas can cause an explosion). 4. If the zinc powder does not dissolve completely, add another 5 ml of the 6.0M HCl solution. Continue adding the acid 5 mL at a time until all of the zinc dissolves. (Not all zinc will dissolve since the zinc used is not pure. The amount of acid to be used must not exceed 20 mL) 5. Place the crucible on a hot plate in the fume cupboard and heat the contents slowly so that the compound will not splatter during the heating process. 6. Heat the compound until it is completely dry, but make sure that the compound does not melt. Chemistry Lab (DKK 1711) 5 7. Allow the crucible to cool to room temperature, and then weigh it. 8. Reheat the crucible, let it cool to room temperature and weigh it again. Reheat this procedure until the difference in mass is not more than 0.02 g. 9. Determine the mass of zinc chloride from the final weight of the sample. Calculate the mass of chlorine in the zinc chloride. 10. Once the mass of zinc and chlorine is obtained, calculate the formula unit of zinc chloride. Questions 1. Why the content is not weighed while it is still hot? Explain. 2. Compare the calculated formula unit with the theoretical formula unit. Explain why there is a difference, and suggest ways to overcome the error. 3. Based on your result, write the equation for the reaction between zinc and HCl. Chemistry Lab (DKK 1711) 6 EXPERIMENT 2 PART A: ACID–BASE TITRATION: PREPARATION OF PRIMARY STANDARD AND STANDARISATION OF A SOLUTION. Objective: 1. To determine concentration of sodium hydroxide solution using a primary standard. 2. To acquire the correct technique in carrying out a titration. Introduction Titration Titration is a procedure in which one solution is used to analyses another. The purpose of a titration is to determine the quantity or concentration of one of the reagents that of the other being know before hand. Standards in Acid Base Titrations One of the substances involved in a titration must be used as a standard for which the amount of substance present is accurately known. The standard can either be in the form of a pure substance or a standard solution (a solution of a accurately known concentration). The concentration of standard solution is often determined by using primary standard such as a substance that can be accurately weighed and pure. In its pure form, the number of moles present can be accurately determined from the measured weight and the known molar mass. A primary standard is used to find molarity of another solution, called a secondary standard. For example, oxalic acid, H2C2O4 and potassium hydrogen phthalate, KHC8H4O4 are two common primary standards used to determine the concentration of bases (secondary standard). Solution of sodium hydroxide, NaOH and hydrochloric acid, HCL that are used in titration often need to be standardized because they contain impurities that cannot be Chemistry Lab (DKK 1711) 7 removed economically during their manufacture. In addition, solid NaOH is hygroscopic (it absorbs moisture), which makes it difficult to obtain its accurate weight. The standardized base can then be used to determine the concentration of other acids. Equivalence point and end point In a titration, the goal is to determine the equivalence point of stoichiometric point which is the point in a titration where stoichiometric amounts of reactants have been reached. Indicators are often used to show when this equivalence point is obtained. A change of colour will be observed when this takes place and this point is known as the end point. The end point and equivalence point should ideally be the same. Chemical equations In an acid-base titration, neutralization takes place at the equivalent point. For example HCl(aq) + NaOH(aq) NaCl (aq) + H2O(l) H2C2O4 (aq) + 2NaOH (aq) Na2C2O4(aq) + 2H2O As in first reaction above, at the end point, the mole ratio of HCl: NaOH is 1:1 while the mole ratio of is H2C2O4(aq) : NaOH is 1:2. If the concentrations of one of the reactants is known, the other can be determined based on this ratio. Material Beaker, burette, conical flask, weighing bottles, analytical balance, NaOH (0.2M), measuring cylinder 50ml, Phenolphthalein, distilled water (in a wash bottle) and hydrated oxalic acid (H2C2O4.2H2O). Chemistry Lab (DKK 1711) 8 Procedure A. Preparation of primary standard 1. Calculate the mass of hydrated oxalic acid, H2C2O4.2H2O needed to neutralize 15 ml of 0.2 M NaOH solution. 2. Weigh out 3 samples of the calculated mass in weighing bottles to the nearest 0.001 g using analytical balance. Record the exact weigh of your samples. 3 Transfer each of these samples into labeled 250 ml conical flasks and add 50 ml of distill water into each using a measuring cylinder. 4. Swirl the flasks to dissolve the acid. B. Standardisation of NaOH solution. 1. Clean and then rinse a burette with distilled water. Next, rinse it through several times with 5-10ml portions of 0.2 M NaOH solution. Discard the used solution. 2. Fill the burette with the 0.2 M NaOH solution. Drain off some of the solution through the tip of the burette to ensure there are no air bubbles. 3. Record the initial burette reading to two decimal points. 4. Add 2 drops of phenolphthalein to the oxalic acid prepared in A. 5. Place a white tile or a piece of white paper under the flask so that the colour of solution is easily observe. 6. Titrate the sample by adding the NaOH solution from burette. During the titration, swirls the flask continuously. Occasionally rinse the sides of the flask with distill water from the wash bottle to bring down any untreated acid solution that may cling to the wall of the flask. 7. Upon reaching the end point, there will be a temporary appearance of a pink color that fades when the solution is swirled. Titrate until the faint pink color persists for more than 30 seconds. This is the actual end point. 8. Record the final burette reading to two decimal points. 9. Repeat the titration with the second sample. 10. Calculate the molarities of NaOH solution from each titration. If the molarities do not agree within 0.006 units, carry out a third titration. Chemistry Lab (DKK 1711) 9 Questions: 1. Calculate the actual molarity of the 0.2 M solution. 2. Why is it unnecessary to know the exact concentration of oxalic acid solution in order to determine the concentration of the base? 3. In step 6 of the section B, you have added a substantial amount of distill water into the conical flask. Does this affect the result? Explain your answer. PART B. ACID-BASE TITRATION: DETERMINATION OF THE CONCENTRATION OF A SOLUTION Objective: 1. To determine the concentration of HCl solution and vinegar using a standard NaOH solution. Apparatus Beaker, burette, conical flasks, graduated cylinder 10 ml, Vinegar, HCl 6.0M, distilled water and standard NaOH solution 0.2 M. Method: A. Determination of molar concentration of diluted HCl 1. Fill a 400ml beaker with 150 ml distills water. 2. Measure 5 ml of 6.0 M HCl using a graduated cylinder. 3. Slowly and carefully, add the acid to water while stirring it with a glass rod. 4. Rinse a clean 25 ml pipette with the acid solution from step 3. Pipette three samples of the acid solution into three separate conical flasks. 5. Fill the burette with a standard 0.2 M NaOH solution. (The exact concentration will be given by your instructor). 6. Just before titration, add 2 drops of phenolphthalein. Chemistry Lab (DKK 1711) 10 7. Titrate with the standard NaOH solution. This first titration will give you a gross reading (an approximate volume of titrant needed for subsequent titrations). 8. Repeat the titration 2 more times. The volume of titrant obtained from the second and third titrations should agree within 0.10 ml. 9 Do another titration if the above precision is not obtained B. Determination of molar concentration of vinegar 1. Pipette 5.0 ml vinegar into a conical flask. 2. Add approximately 50 ml distilled water and 2 drops phenolphthalein. 3. Titrate with the standard NaOH solution using the same procedures as in part A Questions 1. Write the equations for the reactions in A and B 2. Calculate the concentration of the dilute HCL solution. 3. Calculate the molarity of acetic acid in vinegar. 4. Suggest other indicators that can be used for each of the above titration. Chemistry Lab (DKK 1711) 11 EXPERIMENT 3 REDOX TITRATION: TITRATION USING SODIUM THIOSULPHATE Objective 1. To prepare a standard solution of potassium iodate for determining the concentration of sodium thiosulphate solution accurately. 2. To acquire the proper technique of carrying out a titration. Introduction: Redox titrations using sodium thiosulphate as a reducing agent is known as iodometric titration since it is used specifically to titrate iodine. The reaction involved is: I2 + 2 Na2S2O3 2NaI + NaS4O6 I2 + 2 S2O32- 2I- + S4O62- In this equation I2 has been reduced to I- 2S2O32- S4O62- + 2e I2 + 2e 2I- The iodine/thiosulpfate titration is a general method for determining the concentration of oxidizing agent solution. A known volume of an oxidizing agent is added into an excess solution of acidified potassium iodide. The reaction will release iodine: Example: a) With KMnO4 2MnO4- + 16H + 10I- 2Mn2+ + 5I2 + 8H2O b) With KiO3 IO3- + 5I- + 6H+ 3I2 + 3H2O The iodine that released is titrated against a standard thiosulphate solution. From the stoichiometry of the reaction, the amount of iodine can be determined and the concentration of the oxidizing agent which released the iodine can be calculated. Chemistry Lab (DKK 1711) 12 In an iodometric titration, a starch solution is used as an indicator as it can absorb the iodine that released. This absorption will cause the solution to change to a dark blue colour. When this dark blue solution is titrated with the standardized thiosulphate solution, iodine will react with the thiosulpahte solution. When all the iodine has reacted with the thiosulphate solution, the dark blue colour disappears. Iodine dissolved in water by adding an excess of KI so that KI3 which has similar properties to iodine is formed. I2 + KI KI3 I3- + 2e 3I- Oxidizing agents used other than thiosulphate are iron (II) salts, arsenic (III) oxide, sulphur dioxide and stibium(III) oxide. The following are reactions of sulphur dioxide and stibium(III) oxide with iodine : I2 + SO2 + 2H2O 2HI + H2SO4 2I2 + Sb2O3 + 2H2O 4HI + Sb2O5 Experiment aspects to be consider: 1. The indicator (starch) in the iodometric titration is not added in the early stage of the experiment as in acid-base titrations. Starch is only added after titration has begun, i.e. when the colour of the reaction mixture has changed from brown to a light yellow colour. Starch is a colloid that can absorb iodine and form a complex. When this happens, it would be difficult to release the iodine when titrating with the thiosulphate. This will influence the determination of the end point. Hence, the addition of the starch should only be done when the solution is light yellow in colour. At this point, there is only a small amount of released iodine left and the complex formed would also be small in quantity and can easily react with the thiosulphate. 2. As soon as the solution is mixed with KI, titrate immediately with the thiosulphate. This will prevent the iodide from evaporating. Chemistry Lab (DKK 1711) 13 Material Burette, glass rod, filter funnels, pipette filler 25 ml, pipette 100 ml, beaker 250 ml, conical flask, analytical balance, weighing bottle, 50 ml beaker 5 ml and 25 ml measuring cylinder. Solution H2SO4 1.0M, distilled water, starch solution, potassium iodide solids, potassium iodate crystal and sodium thiosulphate solution 0.1M. Method A) Preparation of Potassium Iodate Solution (To be share among two students) 1. Weigh approximately 0.75g of potassium iodate crystal in a 50 mL beaker. 2. Add 25ml of distilled water into the beaker and stir with a glass rod to dissolve all the potassium iodate. 3. Pour the potassium iodate solution through a filter funnel into a 250ml volumetric flask. Rinse the beaker with distilled water and pour it into the volumetric flask. 4. Add more distilled water to the volumetric flask up to the mark on the neck of the flask. Put the stopper in place and shake the flask until you get a homogenous solution. The standard solution is now ready to be used in part B. B) Standardization of 0.1 M Sodium Thiosulphate solution (To be done individually) 1. Rinse and fill a clean burette with 0.1 M sodium thiosulphate solution that is to be standardized. Make sure that there are no air bubbles in your burette. 2. Record the initial reading of your burette. ( Note : You should read your burette at eye level. Accuracy of reading the burette reading should be within the range of 0.05 mL.) 3. Pipette 25.0mL of the standard potassium iodate solution that has been prepared in part A into a 250ml conical flask. 4. Weigh approximately 1 g potassium iodide crystals and add it into the solution in the conical flask. 5. Then add in 10.0 ml of 1.0M sulphuric acid solution and swirl the conical flask Chemistry Lab (DKK 1711) 14 until all the KI has dissolved. 6. i) Titrate immediately the released iodine with the sodium thiosulphate solution while swirling the conical flask until a light yellow solution is obtained. ii) Dilute this solution with distilled water until the total volume is about 100 mL. iii) Add in 1 ml of the starch solution and continue titrating until the blue colour disappears and the solution becomes colourless. This is the end point of the titration process. 7. Record the final reading of your burette. 8. Repeat the titration three times. Questions: 1. Write the equation for reactions between i) iodate ion and iodide ii) iodine and thiosulphate ion 2. Calculate the molarity of KIO3 solution i) From the reaction equations in 1 (i) and (ii) determine the mole ratio between the iodate ion and thiosulphate ion. 3. Calculate the molarity of the sodium thiosulphate solution using the formula : M1V1 =M2V2 Chemistry Lab (DKK 1711) 15 EXPERIMENT 4: BOYLE’S LAW, CHARLES’ LAW & THE IDEAL GAS LAW Objectives 1. To verify Boyle’s Law. 2. To verify Charles’ Law. 3. To determine the molar mass of a gas. Introduction Boyle’s Law states that the volume of a fixed mass of a gas is inversely proportional to its pressure at constant temperature. Mathematically, the law is written as V 1 / P (n, T constant) Charles’ Law states that the volume of a fixed mass of a given gas is directly proportional to its absolute temperature at constant pressure. The law is written as V T (n, P constant) Avogadro’s Principle states that at the same temperature and pressure, all gases of equal volume will contain the same number of molecules. V n (T, P constant) It has been proved that the volume of one mole of an idela gas is 22.4 L at STP (273 K and 1 atm). Thus, combining the three laws, we get V nT / P The above expression is rewritten as V = R( n T/ P) or PV = nRT ………………… (1) This is the ideal gas equation and R is called the gas constant. The number of moles, n, is the mass in gram, m, divided by its molar mass, M. n = m/M Therefore, the ideal gas equation can also be written as PV = (m/M) RT ……………………… (2) Chemistry Lab (DKK 1711) 16 Materials Needle, wire gauze, barometer, tripod stand, rubber band, thermometer, 600 mL beaker, aluminium foil, analytical balance, 125 mL conical flask, Boyle’s Law apparatus, retort stand and clamp, bunsen burner/ hot plate, 100 mL measuring cylinder, Charles’ Law experiment apparatus, ice, methanol, tap water, unknown liquid. A. Boyle’s Law In this experiment, an open-tube manometer, as shown in the Figure 4.1 is used to verify Boyle’s Law. Figure 4.1 The manometer consists of two glass tubes with a uniform cross-sectional area which are connected together by rubber tubing. One end of the tube is left open whilst the other end is closed by attaching a piece of rubber tubing and pinching it with a screw clip. The system contains mercury. The volume of the trapped gas is the volume of the tube between the level of the mercury and the screw clip at the closed end of the tube. It is measured by the height of the gas column. The level of mercury at the open end of the tube is determined by the pressure of the trapped gas. The pressure exerted by the mercury column, PHg, is measured by the difference in height between the two levels. The temperature of the trapped gas is assumed to be the same as room temperature and is constant throughout the experiment. Another assumption is that the cross-sectional area of the tube is constant throughout the tube. Chemistry Lab (DKK 1711) 17 Methods 1. Loosen the clamp holding the open end of the tube from the retort stand. Move the tube up and down gently until the level of the mercury at both ends are equal. (Caution: Be careful not to spill the mercury). 2. Record the heights (in millimeters) of the level of mercury at both the open and closed ends of the tube. 3. Determine the height of the column of gas. 4. Carefully raise the open end of the tube to the highest point possible. Again, record the level of mercury at both ends of the tube. 5. Carefully lower the open end of the tube to the lowest point possible and record the readings for both ends. Refer to Diagram 4.1 and 4.2. 6. Take four more sets of readings by adjusting the open end of the tube to various heights. 7. Record the atmospheric pressure, Pa, reading from a barometer. 8. Calculate the pressure of the gas, Pg, where Pg Pa PH g Pg < Pa ; Pg = Pa - PHg Pg > Pa ; Pg = Pa + PHg Diagram 4.1 Diagram 4.2 Chemistry Lab (DKK 1711) 18 B. Charles’ Law In this experiment, a quantity of air is trapped between the sealed end of a thick- walled glass tube (with a small cross-sectional area) and a movable plug of mercury. If the glass tube is held upright, the plug of mercury will move to a position where the pressure of the air in the tube is equal to the atmospheric pressure plus a small pressure exerted by the plug. Thus, the pressure of the trapped air is constant throughout the experiment. Diagram 4.3 The volume, V, of the trapped air is obtained by multiplying the cross-sectional area of the tube, A, with the height of the air column, h. V=Axh Assuming that the cross-sectional area is constant, the volume is directly proportional to the height, i.e., V h . Therefore, the height of the air column can be used as a measure of the volume in this experiment. By measuring this height at different temperatures we can determine the relationship between the volume of the trapped air and its temperature at a constant pressure. Chemistry Lab (DKK 1711) 19 Note: Mercury is toxic and harmful because it can be absorbed into the skin. Mercury vaporises at high temperature and the vapour is poisonous. Do not touch or break any mercury lump with your fingers. Report any spillage immediately. Method 1. Tie a thermometer to a glass tube containing a plug of mercury with a rubber band. The bulb of the thermometer is placed approximately half-way up the column of the trapped air. Refer to Diagram 4.3. 2. Fill a 100 mL measuring cylinder with tap water and immerse the tube and the thermometer into the water until the air column in the tube is totally submerged. 3. Leave for 5 minutes to ensure that the temperature of the air is equivalent to the temperature of the tap water. 4. Record the temperature and measure the height of the air column, then remove the tube. 5. Repeat steps 3-5 using: (i) warm water (40 °C – 50 °C) (ii) a mixture of ice and water, and (iii) a mixture of ice and methanol instead of tap water in the measuring cylinder. Chemistry Lab (DKK 1711) 20 C. Determination of the molar of a gas Method 1. Cover a 125 mL conical flask with a piece of aluminums foil and tie it loosely around the neck with a rubber band (Refer to Figure 4.2). 2. Pick a tiny hole in the middle of the foil with a needle. Weight the apparatus accurately. 3. Remove the foil and place 5.0 mL of the unknown liquid into the flask. 4. Put back the foil and tie it with a rubber band. Figure 4.2 Figure 4.3 5. Clamp the neck of the flask and immerse it as far as possible into a 600 mL beaker containing water. Refer to Figure 4.3. 6. Place 2 – 3 anti-bumping granules into the water and heat the water to a boil. 7. Record the boiling temperature of the water bath and the barometer reading (atmospheric pressure). 8. Observe the liquid in the flask when the water is boiling. Take the flask out by using the clamp as soon as all the liquid has evaporated, including the liquid which has condensed at the neck of the flask. 9. When cooled, dry the outer wall of the flask and the aluminums foil. 10. Weight the vapour-filled flask with the aluminums foil, rubber band and the condensed unknown liquid. Chemistry Lab (DKK 1711) 21 11. Discard both the foil and the condensed liquid. Fill the flask up to the brim with water and pour it into a measuring cylinder. Record the volume of water. 12. Calculate the molar mass of the unknown liquid using Equation (2). Questions 1. If the amount of the unknown liquid is less, what would be the effect on the value of the molar mass obtained from the experiment? 2. What is the main source of error in the determination of the molar mass? 3. If the flask is not wiped dry, what would be the effect on the molar mass? EXPERIMENT REPORT A. Boyle’s Law 1. Complete the table below: Level of Level of Difference Gas Height of mercury at mercury at = PHg pressure, air Volume, PV the open the closed (mm) Pg (mm) column, h V (mm3) end (mm) end (mm) (mm) Assuming cross-sectional area of tube is constant. 2. Compare the PV values obtained for each set of readings. Discuss. 3. Draw a graph of volume against pressure on a graph paper. By using the graph, explain relationship between pressure and volume. Chemistry Lab (DKK 1711) 22 B. Charles’ Law 1. Complete the following table: Condition Temperature Height of gas column Warm water Tap water Ice-water Ice-methanol 2. Plot the height of the column, h, against temperature, T, in Kelvin on a graph paper and extrapolate the line until h = 0. this temperature is the absolute zero temperature. 3. Based on the graph, (i) state the relationship between volume and temperature, (ii) measure the gradient and state the unit. C. Determination of the molecular mass of a gas Readings 1. Mass of flask + rubber band + cover + vapour 2. Mass of flask + rubber band + cover 3. Mass of vapour (condensed liquid) 4. Temperature of boiling water (°C) 5. Barometric pressure (mm Hg) 6. Volume of flask Calculation for the molar mass of the unknown liquid (gas): Molar mass of gas: _________________________ Chemistry Lab (DKK 1711) 23 EXPERIMENT 5: MOLECULAR GEOMETRY Objectives 4. To view the structure of molecules in 3-dimension. 5. To determine the shape of molecules using the VSEPR theory. 6. To determine the type of hybridisation of the central atom. Introduction The first step toward visualizing what a molecule looks like is to convert its molecular formula to its Lewis structure. A Lewis structure is a 2-dimensional structural formula that shows how the atoms are attached to each other within the molecule but does not reveal the overall shape. Together with Lewis structures, we use the Valence Shell Electron Pair Repulsion (VSEPR) theory to determine the true geometry of a molecule by converting Lewis structures into 3-dimensional shapes. In the VSEPR theory, the repulsion between two pairs of electrons (i.e. bonding pairs or lone pairs) in the valence shell is the dominant factor that determines the geometry of a molecule. The electron pairs around the central atom are oriented as far apart as possible to reduce the repulsion between them. In this case, the lone pairs repel more strongly than the bonding pairs. The true shape of a molecule is determined by the position of the terminal atoms which are bonded to the central atom. The lone pairs will only influence the bond angles but not the shape. From the basic geometry of the molecule, you can predict the hybridisation of the central atom. The number of hybrid orbitals is the same as the number of electron pairs around the central atom. Materials 3-D molecular model set. Chemistry Lab (DKK 1711) 24 Methods A. Basic Geometry 1. Using the 3-D model set, construct structures with the following geometry: Linear Tetrahedral Octahedral Trigonal planar Trigonal bipyramidal 2. Draw the structures in the table provided. 3. Determine the bond angles and write the general formula to represent the molecule. Use the symbol A as the central atom and X as the terminal atom. B. Molecular Geometry of HCl, CO2 and BH3 1. Construct the molecular models of HCl, CO2 and BH3. 2. Determine the geometry and bond angles of the molecules that you have constructed. 3. Write the general formula in the table. C. Molecular Geometry of CH4, NH3 and H2O 1. Construct the molecular models of CH4, NH3 and H2O. 2. Determine the geometry and bond angles of the molecules that you have constructed. 3. Write the general formula in the table. D. Molecular Geometry of PF5, SF4, AsCl5, ICI3, XeF2 1. Construct the molecular models of PF5 and AsCl5. 2. Determine the geometry and bond angles of the molecules that you have constructed. Write the general formula. 3. Construct all the possible molecular models of ICI3, SF4 and XeF2 that you can think of. Determine the most stable geometry for each molecule. Write the bond angles and write the general formula in the table. Chemistry Lab (DKK 1711) 25 E. Molecular Geometry of SF6, XeF4 and BrF5. 1. Construct all the possible geometry of SF6, XeF4 and BrF5. 2. Determine the most stable geometry based on VSEPR theory for each molecule and the bond angles. Questions 1. What is the most important factor in determining the geometry of a molecule or an ion? 2. List down the steps to determine the molecular shape of a compound. 3. Explain why the geometry of H2O is not linear whereas CO2 is linear. Chemistry Lab (DKK 1711) 26 EXPERIMENT REPORT Basic Geometrical Drawings and Bond Angles Molecular Geometry Bond Angles General Formula Chemistry Lab (DKK 1711) 27 Molecular Geometry and Bond Angles for Particular Compounds Compounds Molecular Geometry Bond Angles General Formula HCl CO2 BH3 CH4 NH3 H2 O PF5 AsCl5 SF4 ICI3 XeF2 SF6 BrF5 XeF4 Chemistry Lab (DKK 1711) 28 EXPERIMENT 6: CHEMICAL EQUIBLIRIUM Objectives 1. To study the effect of concentration and temperature on chemical equilibrium. 2. To determine the equilibrium constant, Kc, of a reaction. Introduction There are two kinds of chemical reactions, i.e. irreversible and reversible reactions. A reversible reaction will reach a dynamic equilibrium when the rate of the forward reaction equals to the rate of the reverse reaction. At this stage, one cannot see any change in the system. However, this does not mean that the reactions have stopped. The reactions are still occurring, but at the same rate. The factors that influence chemical equilibrium are: Concentration Temperature Pressure (for reactions that involve gases) Le Chatelier’s principle is used to determine the position of the equilibrium when one of the above factors is changed. Le Chatelier’s principle states that if a system at equilibrium is disturbed by a change in temperature, a change in pressure or a change in the concentration of one or more components, the system will shift its equilibrium position in such a way as to counteract the effects of the disturbance. The Effect of Concentration Consider a general reaction as follows: A + B C + D According to the Le Chatelier, when the concentration of any substance in mixture at the equilibrium is changed, the equilibrium position will shift either forward or backward to restore the equilibrium. For example, if reactant A (or B) is added to a mixture at equilibrium, the reaction will shift forward to reduce the concentration of A (or B) until equilibrium is reached again. Chemistry Lab (DKK 1711) 29 On the other hand, if reactant C (or D) is added, the equilibrium will shift in the direction that will reduce the concentration of C (or D), i.e. from right to left until equilibrium is reached again. The effect of temperature The effect of temperature on an equilibrium system depends on wheter the reaction is exothermic or endothermic. Consider the following system: E+F G + heat Assume the forward reaction is exothermic (i.e. heat is considred as one of the products). Heating the system will cause the equilibrium to shift in the reverse direction so as to reduce the added heat. Thus, the concentration of E and F increases whilst the concentration of G decreases. However, when the system is cooled, the equilibrium will move forward to increase the heat in the system. The same principle can be applied to explain an endothermic system. In this experiment, you will study the effect of changes in concentration and temperature on two equilibrium systems. You can notice the change in equilibrium through changes in colour or phases such as formation of precipitates or solvation. Materials Burette, ice bath, test tube, water bath, 10 mL pipette, 100 mL beaker, 100 mL conical flask, 10 mL measuring cylinder, 100 mL measuring cylinder, CoCl2 0.2M, KSCN 0.1M, 2.5 M NaOH, distilled water, Fe(NO3)3 0.1M, Antimony chloride (0.5 M SbCl3 in 6.0M HCl) Chemistry Lab (DKK 1711) 30 A: The effect of concentration in the formation of thiocyanoferum(III) complex ion The thiocyanoferum(III) complex ion is formed when iron(III) ion, Fe3+, is added to the thiocyanate ion, SCN-. The equation for the reaction is Fe3+(aq) + SCN-(aq) Fe(SCN) (aq) 2 Yellowish Blood red Methods 1. Place 2 mL of 0.1M iron(III) nitrate, Fe(NO3)3, solution and 3 mL of 0.1 M potassium thiocyanate, KSCN, solution into a 100 mL beaker. 2. Add 50 mL distilled water to reduce the intensity of the dark red colour formed. 3. Place about 5 mL of this solution into three separate test tubes. a) To the first test tube, add 1 mL of 0.1 M iron(III) nitrate solution. b) To the second test tube, add 1 mL of 0.1 M potassium thiocyanate solution. c) To the third test tube, add 6-8 drops of 2.5 M NaOH. 4. Record your observations in table form. Question Discuss the effect of adding iron(III) nitrate, potassium thiocyanate and sodium hydroxide on the position of the above equilibrium. Chemistry Lab (DKK 1711) 31 B: The effect of temperature In aqueous solution, the cobalt(II) ion, Co2+, exists as hexaaquocobalt(II) complex ion, Co(H2O)62+, which is pink. Other cobalt(II) complex ions have different colours, e.g., CoCl42- is blue. For the following equation, the position of the equilibrium can be changed to produce a solution which is either blue or pink. This is because the position of the equilibrium depends on the relative concentration of the complex ions. Co(H2O)62+ + 4Cl- CoCl42- + 6H2O Pink Blue Method 1. Measure 4 mL of 0.2 M CoCl2 solution with a measuring cylinder and place it into conical flask. 2. Measure 20 dops concentrated hydrocloric acid and add it into the same flask (Note: This should be done in the fume cupboard). 3. Shake the flask. Make sure you get a purple solution, indicating a mixture of pink and blue. If you get a pink solution, add more acid, and if it is blue, add some distilled water. 4. Divide the purple solution into three test tubes: (a) Leave one test tube at room temperature. (b) Place the second test tube in the ice bath. (c) Place the third test tube in a water bath at 80°C – 90°C. 5. Record the colour of the solution in each test tube. Remove the second and the third test tube and leave them at room temperature. Observe the change in colour. Question Determine whether the above system is exothermic or endothermic. Explain your answer based on your results. Chemistry Lab (DKK 1711) 32 C. Determination of the equilibrium constant. The following reaction is an example of a heterogenous system, SbCl3(aq) + H2O(l) SbOCl(s) + 2HCl(aq) The expression for the equilibrium constant is, Kc HCl 2 SbCl3 Note that the concentration of pure liquid (i.e. water) and the concentration of solid (i.e. SbCl3) are not included in the equilibrium constant expression. Method 1. Pipette 5.0 mL of antimony chloride, SbCl3 solution (0.5 M SbCl3 in 6.0 M HCL) into a conical flask. 2. Carefully release distilled water from a burette into the conical flask until a chalky solution is obtained. 3. Record the volume of water added. Questions 1. Calculate the concentration of antimony chloride and hydrochloric acid in the solution at equilibrium. 2. Calculate the value of the equilibrium constant, Kc. 3. Explain why the concentration of the pure liquid and solid is excluded from the equilibrium constant expression for the heterogeneous system. Chemistry Lab (DKK 1711) 33 EXPERIMENT 7: pH MEASUREMENT AND ITS APLLICATIONS Objectives 1. To use various methods to measure the pH of acids, bases and salts. 2. To determine the dissociation constant of acetic acid Introduction The concentration of hydrogen ions in a solution is not expressed in moles or molarity but is expressed in terms of its pH. pH is defined as the negative logarithm of hydrogen ion concentration, (H+). pH = - log (H+)…………………………………………… (1) Example: (H+) = 1 x 10-4 mol/L pH = -log (1 x 10-4) =4 The pH scale ranges from 1 to 14. A neutral solution has a pH of 7. An acidic solution has a pH value of less than 7 while a basic solution has a pH value greater than 7. There are two methods of measuring pH in the laboratory. The fist method involves the use of indicators, pH paper, litmus paper and universal indicator. The second method is using the pH meter. Acids or bases which ionize completely are called strong acids or strong bases. An example of a strong acid is HCl and a strong base is NaOH. Weak acids and weak bases do not ionized completely. An example of a weak acid is acetic acid, CH3COOH, and that of a weak base is ammonia, NH3. Consider the ionization of a weak acid, HA. Chemistry Lab (DKK 1711) 34 HA (aq) H+ + A- ………………………………………….. (2) The equilibrium expression for the above reaction is written as: ( H )( A ) Ka ………………………………………….. (3) HA where (H+), (A-) and (HA) are species concentration at equilibrium and Ka is the ionization constant for acid HA (The subscript ‘a’ in Ka stands for acid). The same equation can be written for weak bases. The ionization constant for weak bases is written as Kb. One method to determine Ka is by measuring the pH of an acid with a known concentration. (H+) at equilibrium can be calculated by substituting the pH value in equation (1). In equation (2), the stoichiomectric coefficient for both H+ and A- is the same, therefore (A) = (H+). For acids that undergo very little ionization, the value for (HA) in equation (3) is the same as the initial molarity of the acid (HA). Example The molarity of acid HA, (HA) = 0.001 M pH = 4 From Equation (1), (H+) = 1 x 10-4 M From Equation (2), (A-) = (H+) = 1 x 10-4 M ( H )( A ) (1x104 )(1x104) From Equation (3), Ka = ( HA) 0.01M An easier method to determine Ka is by adding a weak acid solution to its conjugated base solution. The product of this process is in an acidic buffer solution. The conjugated base is obtained from the salt that is produced using the titration method. Chemistry Lab (DKK 1711) 35 In this method, a known weak acid (HA) is divided into two equal parts, X and Y. The first part, X is titrated with NaOH (a strong base) using phenolphthalein as indicator to detect the formation of a slat solution. A change of colour from colorless to pink indicates the end point. The equation for the reaction is as follows: OH-(aq) + HA(aq) → A-(aq) + H2O(l) In this reaction, all the HA in the conical flask reacts with NaOH to form NaA and H 2O. NaA ionizes completely to form A- and Na+. The number of moles of A- formed is the same as the number of moles of HA in the second part, Y (which has not been titrated). The second part of the weak acid HA is added to the contents in the conical flask (which contains the salt NaA). In this mixture, the concentration of HA will be equal to the concentration of A- (from the salt). From equation (3), ( H )( A ) Ka ( HA ) Since (A-) = (HA), Ka = (H+) The value of (H+) can be obtained by measuring the pH and from this; the value of Ka can be calculated. Materials Burette, pH Meter, pH Paper, Test tubes, 25 mL Pipette, 250 mL Conical flask, NaCl (0.1 M), NaOH (0.2M) Methyl violet, NH4NO3 (0.1M), Alizarin yellow, Metyl orange, Phenolphtalein, Universal inidicator, CH3COONa (0.1 M), NH3 (0.1 and 1.0 M), HCl (0.01 M and 1.0 M), CH3COOH (0.1 and 1.0 M). Chemistry Lab (DKK 1711) 36 Methods A. Determination of pH of acidic and basic solutions 1. i) Determine the pH of 0.01 M, 1.0 M HCl using; a. pH paper. Refer to the pH colour chart of the pH paper. b. Universal indicator. Refer to pH colour chart of universal indicator. ii) Determine the pH of 0.01 M, 1.0 M HCl using a pH meter. 2. i) Fill in two test tubes with 2 mL HCl 0.01 M and HCl 1.0 M respectively. ii) Add 2 drops of Methyl violet in both test tubes and note the colour change. 3. i) Using pH paper with an acidic range and a pH meter, determine the pH of 0.1 M and 1.0 M CH3COOH. ii) Using Methyl violet and Methyl orange, determine the colour of the indicator in both solutions. 4. i) Using pH paper with a basic range, determine the pH of 0.1 M and 1.0 M NaOH. ii) Determine the colour of Alizarin yellow in 0.1 N and 1.0 M NaOH. iii) Determine the colour of Alizarin yellow in 0.0 M and 1.0 M NH3 B. Determination of pH of salt solutions Using pH paper with appropriate range and universal indicator, determine the pH of the following salt solutions; NaCl 0.1 M, NaOAc 0.1 M and NH4NO3 0.1 M. Determine whether the salt solutions are acidic, basic or neutral. Chemistry Lab (DKK 1711) 37 C. Determination of the dissociation constant of a weak acid, Ka 1. Pipette 25 mL of 0.1 CH3COOH into two conical flasks, labeling them X and Y. 2. Add two to three drops of phenolphthalein into the conical flask labeled X and titrate it with 0.2 M NaOH. (Titrate slowly base when the volume of base reaches 10 mL). The end point is reached when the solution becomes pink. Record the initial and the final reading of the burette. 3. Mix the resulting solution in step (2) with 25 mL of 0.1 M CH3COOH in conical flask Y. Determine the pH of this mixture using the pH meter. 4. Calculate Ka from the value of pH obtained (3). Question: 1. From the pH values obtained from both concentrations of acetic acid, calculate Ka for acetic acid. Compare these values with those obtained from part (C) Chemistry Lab (DKK 1711) 38 EXPERIMENT 8 THERMOCHEMISTRY: DETERMINING THE HEAT OF REACTION Objectives a. To determine the value of heat capacity of a calorimeter. b. To determine the hat of neutralization of HCl and NaOH. Introduction Chemical reactions can release or absorb heat. These can be measured by using a calorimeter. A calorimeter is a container that is thermally isolated from the environment. Heat released by the chemical reaction (-qr) is absorbed by the solution and the calorimeter. -qr = qr + qc …………………………………………… (1) Where qs = heat absorbed by solution qc = heat absorbed by calorimeter The heat absorbed is proportional to the change in temperature. The proportionality constant, C is known as the heat capacity of a sample. Heat capacity is defined as the amount of heat required to increase the temperature by 1oC. q = Cc∆T ………………………………………………... (2) For a homogeneous sample such as solutions, the heat absorbed is proportional to the mass of the sample and the increase in temperature. Proportionality constant, c is known as specific heat capacity of solution per unit mass. The specific heat capacity for water is 4.18 JK-1g-1. The specific heat capacity of a very dilute solution is equivalent to the specific heat capacity of pure water. Therefore, the value of the specific heat capacity of all the solutions used in this experiment will be 4.18 JK-1g-1. Chemistry Lab (DKK 1711) 39 Heat that is being released or absorbed (qr) can be determined by measuring the temperature before and after the reaction. - qr = Cc∆T + mcs∆T …………………………………(3) where ∆T = final temperature of system – initial temperature of system ms = mass solution Cc = heat capacity of calorimeter Cs = specific heat capacity of solution For this experiment the density of the solution is assumed to be the same as the density of water (1 g cm-3). Therefore the mass of the solution can be calculated. Materials Stop watch, Calorimeter, Conical flask, Thermometer, 100 mL beaker, Measuring cylinder, Bunsen burner/water bath, HCl (1.0 M), NaOH (1.0 M, Distilled water.) Methods A. Determination of the heat capacity of a calorimeter 1. Measure the temperature of an empty calorimeter by putting a thermometer inside the calorimeter, T1. 2. Pour 50 mL distilled water into a 100 mL beaker. 3. Heat the beaker to a temperature between 50 – 60oC. 4. Pour the hot water into the calorimeter. Immediately measure the initial temperature of the hot water, T2. 5. Observe the decrease in temperature until it says constant for a few minutes. Record this temperature, T3. Temperature/ oC T1 T2 T3 Chemistry Lab (DKK 1711) 40 B. Determination of the heat of neutralization of HCl 1.0 M and NaOH 1.0 M. 1. Using a measuring cylinder, measure 25 mL NaOH 1.0 M and 25 mL HCl 1.0 M. 2. Pour the NaOH solution into the calorimeter and the HCl solution into a conical flask. Record the initial temperature of each solution. 3. Without removing the thermometer, lift the cover slightly and quickly pour all the HCl solution the calorimeter. 4. Quickly replace the cover of the calorimeter. 5. Stir the solution and begin to record the temperature every 15 seconds for at least 2 minutes. Note the maximum temperature. 6. Repeat this experiment twice. Tinitial (NaOH) = _________oC Tinitial (HCl) = _________oC Time/ seconds 15 30 45 60 75 90 105 120 o Tsolution / C Question Calculate the value of the heat of neutralization of the above reaction and compare it with the theoretical value. Give your comments. Chemistry Lab (DKK 1711) 41 EXPERIMENT 9 NERNST EQUATION: DETERMINATION OF ELECTRODE POTENTIALS OF ZINC AT VARIUOS CONCENTRATIONS Objectives 1. To determine the effect of concentration of electrolyte solutions on the reduction potential of Zn2+/ Zn half cell. 2. To compare reduction potential values of Zn2+/Zn half cell obtained from the Nernst equation. Introduction Galvanic cell is a type of electrochemical cell where redox reactions (oxidation and reduction) occurs spontaneously and generates electric flow. For a cell that is connected by a salt bridge, oxidation will occur at the anode and this will cause electrons to flow through a wire circuit to the cathode where reduction reactions will occur. It can now be said that a full circuit is obtained. Reduction potential that is obtained from a standard condition where solution concentration is 0.1 M, gas partial pressure is 1 atm and temperature is 25oC is known as standard reduction potential. These standard reduction potential values are arranged in a certain order and this list is known as the Reduction Potential Table. In this table, species with more positive values are more easily reduced while species with more negative values will undergo oxidation. Cell potential at conditions that are not standard is given by this equation; Ecell = Ecathode - Eanode Where Ecathode and Eanode are the electrode potentials or half cell potentials. Chemistry Lab (DKK 1711) 42 At 25oC the electrode potential is given by the Nernst equation: 0.591 1 E electrode E o electrode log n ( X n ) for a half reaction of Xn+ + ne → X Where, Eelectrode = Reduction potential of electrode Eoelectrode = Standard reduction potential of electrode n = no. of moles electron involved in the half reaction n+ (X ) = ionic concentration in unit mole L-1 The above Nernst equation shows clearly that electrode potential values and consecutive cell potential values depend on the concentration of ions involved. Materials Voltmeter, Salt bridge, Sand Paper, Pipette filler, Crocodile clips, 50 mL beaker, 25 mL Pipette, 25 mL volumetric flask, 50 mL volumetric flask, 5 mL graduated pipette, 1 mL graduated pipette, CuSO4 (0.1 M), ZnSO4 (0.1 M), Zinc electrode, Copper electrode. Diagram 1: Galvanic cell Chemistry Lab (DKK 1711) 43 Methods A. Measurement of Cell Potential between 0.1 M ZnSO4 and 0.1 M CuSO4 1.Pipette 25 mL 0.1 M ZnSO4 solution and 0.1 M CuSO4 solution previously prepared into two separate 50 mL beakers. 2.Connect both solutions with a salt bridge as seen in the diagram for apparatus setup. 3. By using a 2V range from the voltmeter, record the cell potential, Ecell (up to three decimal points) into your log book. Note: Do not discard the 0.1 M CuSO4 solution. It can be reused in part B and C. D. Measurement of Cell Potential between 0.01 M ZnSO4 and 0.1 M CuSO4. 1.Preparation of 0.001 M ZnSO4 solution. a. By using a graduated pipette, pipette an ascertained volume (needs to be calculated) of 0.1 M ZnSO4 into a 25 mL volumetric flask. b. Add distilled water until the mark, put the stopper in its place and shake the flask until you get a homogeneous solution. 2.Pour all the 0.001 M ZnSO4 solution into a 50 mL beaker. 3. Set up the apparatus as in part A by replacing the 0.1 M ZnSO4 with the 0.001 M ZnSO4 solution. 4.Record the cell potential value as in part A. C. Measurement of Cell Potential between 0.001 M ZnSO4 and 0.1 M CuSO4 1. Preparation of 0.01 M ZnSO4 solution. a. By using a graduated pipette, pipette an ascertained volume (needs to be calculated) of 0.1 M ZnSO4 into a 25 mL volumetric flask. Chemistry Lab (DKK 1711) 44 b. Add distilled water until the mark, put the stopper in its place and shake the flask until you get a homogeneous solution. 2. Pour all the 0.01 M ZnSO4 solution into a 50 mL beaker. 3. Set up the apparatus as in part A by replacing the 0.1 M ZnSO4 with the 0.01 M ZnSO4 solution. 4. Record the cell potential value as in part A. Method Cell Potential (Ecell) A B C Calculation 1.Calculate the value of half cell potentials, EZn2+/Zn for every concentration by using:- Ecell = E Cu2+/Cu – Ezn/Zn2+ where EoCu2+/Cu = +0.34 V 2.Calculate the value of half cell potential EZn2+/Zn by using the Nernst equation:- 0.0591 1 E Zn2+/Zn = EoZn2+/Zn - where EoZn2+/Zn = -0.76 V 2 log Zn 2 3. Complete the report sheet in the appendix. Question: Based on your results, state the effect of concentration changes on cell potential 2+ values of electrode potential E Zn /Zn i.e. whether the values have become more positive or more negative when the concentration is decreased Chemistry Lab (DKK 1711) 45 EXPERIMENT 10: REACTION RATES Objectives To study the effect of concentration, temperature and catalyst on the rate of reaction. Introduction The rate of a reaction is the change in concentration of the reactants or products per unit time. Among the factors which influence the rate of a reaction are the concentration of the reactants, temperature and catalyst. The rate of a reaction can be studied by observing the change in the chemical properties or the change in physical properties of species involved in the reaction. The rate of a reaction affects the time factor, i.e. the faster the reaction occurs, the shorter the time for the reaction to complete. The relationship between the concentration of the chosen reactants and time can be used to measure the rate of reaction. The rate of reaction can also be influenced by changes in temperature and the presence of catalyst. Materials Glass rod, Boiling tubes, Steam bath, Stop-watch, thermometer, Bunsen Burner, 10 mL Pipette, 50 mL Burette, 100 mL Conical flask, HCL (0.10 M), MnSO4 (10 %), KMnO4 (0.2 M), H2SO4 (2.00 M), Na2S2O3 (0.10 M), H2C2O4 (0.25 M) Methods C. Determination of the effect of concentration on the rate of reaction a. Place 50 mL of 0.10 M sodium thiosulphate, Na2S2O3 into a conical flask. Put the conical flask on a piece of white paper on which a cross has been drawn. b. Add 10 mL of 0.10 M HCl into the conical flask and start the stop-watch. Stir continuously with a glass rod until the cross on the paper is not available. c. Note the time required for the cross to disappear. d. Repeat steps 1-3 with the addition of distilled water to the sodium thiosulphate as instructed in Table 3.1 Chemistry Lab (DKK 1711) 46 Volume of 0.1 M Volume of Concentration Volume of 0.10 Time for cross Na2S2O3 distilled water of Na2S2O3 M HCl solution to disappear Solution (mL) added (mL) (mL) (seconds) 50.00 0.00 10.00 40.00 10.00 10.00 30.00 20.00 10.00 20.00 30.00 10.00 10.00 40.00 10.00 Table 3.1: Concentration of reactants D. Determination of the effect of temperature and catalyst on the rate of reaction a. Place 10 mL og 0.25 M oxalic acid, H2C2O4 solution into two boiling tubes labeled A1 and A2. b. Fill boiling tube B1 with 5 mL of 0.2 M KMnO4 solution and 10 mL of 2.0 M H2SO4 solution, stir with a glass rod. c. Fill boiling tube B2 with 0.2 M KMnO4 and 2.0 M H2SO4 solution and 10 mL as boiling tube B1, and then add 5 drops of 10 % MnSO4 solution. d. Place boiling tubes A1, B1, A2 and B2 in a hot water bath with preset temperature. e. i) Mix the solutions in boiling tubes A1 and B1. Note the time taken for the mixture to become colorless. ii) Mix the solutions in boiling tubes A2 and B2. Note the time taken for the mixture to become colorless. f. Repeat steps 1-5 for the other temperatures as in Table 3.2. Temperature Time taken for solution to become colorless (seconds) (oC) Without catalyst MnSO4 With catalyst MnSO4 (A1+B1) (A2+B2) 30 40 50 Table 3.2: Effect of temperature and catalyst on reaction rate Chemistry Lab (DKK 1711) 47 Calculations and Discussions A. The effect of the concentration on the rate of reaction 1. Calculate the concentration of the sodium thiosulphate solution after the dilution 1 and the value of ( s 1 ) t 2. Plot a graph of the concentration of the sodium thiosulphate solution after the 1 dilution against the value of . t 3. Determine the relationship between the concentration of the sodium thiosulphate solution with time and the rate of reaction. B. The effect of the temperature and catalyst on the rate of reaction. 1 1. Calculate the value of . t 1 2. Plot a graph of ( s 1 ) against the temperature of the reactions for the mixtures of t A1 + B1 solutions and A2 + B2 on the same graph. 3. based on your graph, deduce the relationship between i. Temperature and the rate of reaction ii. Catalyst and the rate of reaction.