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CHEMISTRY LABORATORY MANUAL SEMESTER I & II SK015 & SK025

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LAB MANUAL SK015 & SK025

MATRICULATION DIVISION CHEMISTRY LABORATORY MANUAL SEMESTER I & II SK015 & SK025 TWELFTH EDITION MATRICULATION DIVISION MINISTRY OF EDUCATION MALAYSIA CHEMISTRY LABORATORY MANUAL SEMESTER I & II SK015 & SK025 TWELFTH EDITION First Printing, 2003 Second Printing, 2004 Third Printing, 2005 (Sixth Edition) Fourth Printing, 2006 (Seventh Edition) Fifth Printing, 2007 (Eighth Edition) Sixth Printing, 2011 (Ninth Edition) Seventh Printing, 2013 (Tenth Edition) Eighth Printing, 2018 (Eleventh Edition) Ninth Printing, 2020 (Twelfth Edition) Copyright © 2020 Matriculation Division Ministry of Education Malaysia ALL RIGHTS RESERVED. No part of this publication may be reproduced or transmitted in any form or by any means, electronic or mechanical, including photocopying, recording or any information storage and retrieval system, without the prior written permission from the Director of Matriculation Division, Ministry of Education Malaysia. Published in Malaysia by Matriculation Division Ministry of Education Malaysia, Level 6 – 7, Block E15, Government Complex Parcel E, Federal Government Administrative Centre, 62604 Putrajaya, MALAYSIA. Tel : 603-88844083 Fax : 603-88844028 Website : http://www.moe.gov.my/bmkpm Printed in Malaysia by Cataloguing-in-Publication Data Malaysia National Library Chemistry Laboratory Manual Semester I & II SK015 & SK025 Twelfth Edition eISBN: 978-983-2604-50-1 NATIONAL EDUCATION PHILOSOPHY Education in Malaysia is an on-going effort towards further developing the potential of individuals in a holistic and integrated manner, so as to produce individuals who are intellectually, spiritually, emotionally and physically balanced and harmonious based on a firm belief in and devotion to God. Such an effort is designed to produce Malaysian citizens who are knowledgeable and competent, who possess high moral standards and who are responsible and capable of achieving a high level of personal well- being as well as being able to contribute to the betterment of the family, society and the nation at large. NATIONAL SCIENCE EDUCATION PHILOSOPHY In consonance with the National Education Philosophy, science education in Malaysia nurtures a science and technology culture by focusing on the development of individuals who are competitive, dynamic, robust and resilient and able to master scientific knowledge and technological competency. CONTENTS Page i Learning Outcomes Introduction iii - iv v  Laboratory Safety vi  Preparation For Experiment  Report Writing Semester I Experiment Title 1 Determination of the formula unit of a compound 1 3 2 Acid Base Titration – Determination of the concentration of 7 hydrochloric acid solution 11 16 3 Determination of the molar mass of a metal 20 4 Charles’ Law and the ideal gas Law 5 Chemical Equilibrium 6 pH measurement and its applications Semester II 24 Experiment Title 28 31 1 Rate of reaction 34 2 Determining the heat of reaction 38 3 Electrochemical cells 41 4 Reactions of aliphatic and aromatic hydrocarbons 5 Reactions of hydroxy compounds 44 6 Aldehydes and ketones 45 References Acknowledgements SK015 & SK025 Lab Manual 1.0 Learning Outcomes 1.1 Matriculation Science Programme Educational Objectives Upon a year of graduation from the programme, graduates are: 1. Knowledgeable and technically competent in science disciplines in-line with higher educational institution requirement. 2. Able to communicate competently and collaborate effectively in group work to compete in higher education environment. 3. Able to solve scientific and mathematical problems innovatively and creatively. 4. Able to engage in life-long learning with strong commitment to continue the acquisition of new knowledge and skills. 1.2 Matriculation Science Programme Learning Outcomes At the end of the programme, students should be able to: 1. Acquire knowledge of science and mathematics fundamental in higher level education. (PEO 1, MQF LOD 1) 2. Demonstrate manipulative skills in laboratory work. (PEO 1, MQF LOD 2) 3. Communicate competently and collaborate effectively in group work with skills needed for admission in higher education institutions. (PEO 2, MQF LOD 5) 4. Apply logical, analytical and critical thinking in scientific studies and problem solving. (PEO 3, MQF LOD 6) 5. Independently seek and share information related to science and mathematics. (PEO 4, MQF LOD 7) Updated: 12/03/2020 i SK015 & SK025 Lab Manual 1.3 Course Learning Outcome 13.1 Chemistry 1 At the end of the course, student should be able to: 1. Explain basic concepts and principles of physical chemistry in novel and real life situations. (C2, PLO1, MQF LOD1) 2. Demonstrate the correct techniques in handling laboratory apparatus and chemicals when carrying out experiments. (P3, PLO2, MQF LOD2) 3. Solve chemistry related problems by applying the basic concepts and Principles in physical chemistry. (C4, PLO4, CTPS3, MQF LOD 6) 13.2 Chemistry 2 At the end of the course, student should be able to: 1. Explain basic concepts and principles of organic chemistry in novel and real life situations. (C2, PLO1, MQF LOD1) 2. Demonstrate the correct techniques in handling laboratory apparatus and chemicals when carrying out experiments. (P3, PLO2, MQF LOD2) 3. Solve chemistry related problems by applying the basic concepts and principles in organic chemistry. (C4, PLO4, CTPS3, MQF LOD 6) 1.4 Objectives of Practical Sessions The main purpose of the experiment is to give the student a better insight of the concepts of Chemistry discussed in the lectures by carrying out experiments. The aims of the experiments are to enable students to: 1. Learn and practise the necessary safety precautions in the laboratory. 2. Plan, understand and carry out the experiment. 3. Use the correct techniques in handling the apparatus. 4. Acquire scientific skills in measuring, recording and analysing data. 5. Observe, measure and record data by giving consideration to the consistency, accuracy and units of the physical quantities. 6. Determine the errors in various physical quantities obtained in the experiments. 7. Deduce logically and critically the conclusion based on observation and data analysis. Updated: 12/03/2020 ii SK015 & SK025 Lab Manual 2.0 Laboratory Safety The Science Matriculation Programme requires the students to attend practical classes two hours a week to complete six experiments each semester. In order for the laboratory to be a safe place to work in, students should learn laboratory rules and regulations, including the correct way of using laboratory apparatus and handling of chemicals before starting any experiments. Laboratory rules and regulations. 1. Attendance is COMPULSORY. If you are unable to attend any practical class, you should produce a medical certificate or a letter of exemption. 2. Read, understand and plan your experiment before pre-lab sessions and practical classes. 3. Wear shoes, lab coats and safety goggles at all times in the laboratory. 4. Tie long hair or tuck head scarf under your lab coat 5. Do not wear contact lenses during experiments. 6. Foods and drinks are not allowed in the laboratory. 7. Do not perform any unauthorised experiments! Understand and follow the specified procedures for each experiment. 8. Do not waste chemicals. Take only sufficient amount of chemicals needed for your experiments. 9. Replace the lids or stoppers on the reagent bottles or containers immediately after use. 10. Do not remove chemicals from the laboratory. 11. Handle volatile and hazardous compounds in the fume cupboard. Avoid skin contact with all chemicals, wash off any spillages. 12. Clean up spillages immediately. In case of a mercury spillage, do not touch the mercury. Notify your instructor immediately. 13. Ensure there are no flames in the vicinity before working with flammable chemicals 14. NEVER leave an ongoing experiment unattended. 15. Be aware or familiar with the location and proper way of handling safety equipment, including eyewash, safety shower, fire blanket, fire alarm and fire extinguisher. 16. Turn off bunsen flames when not in use. Notify your instructor immediately of any injury, fire or explosion 17. Do not throw any solid wastes into the sink. Dispose any organic substances in the waste bottles provided. Updated: 12/03/2020 iii SK015 & SK025 Lab Manual 18. Wash all glasswares after use and return the apparatus to its appropriate places. 19. Keep your work area clean and tidy. 20. Notify your instructor immediately of any injury, fire or explosion I have read and understood the laboratory rules and regulations as stated above. I agree to abide by all these rules, follow the instructions and act responsibly at all times. Signature : Date : Name : Practicum : Matric number : Signature Instructor : Date : Updated: 12/03/2020 iv SK015 & SK025 Lab Manual 3.0 Ethics in the laboratory 1. Follow the laboratory rules. 2. Students must be punctual for the practical session. Students are not allowed to leave the laboratory before the practical session ends without permission. 3. Co-operation between members of the group must be encouraged so that each member can gain experience in handling the apparatus and take part in the discussions about the results of the experiments. 4. Record the data based on the observations and not based on any assumptions. If the results obtained are different from the theoretical value, state the possible reasons. 5. Get help from the instructor or the laboratory assistant should any problems arise during the practical session. 4.0 Preparation for experiment 4.1 Pre-lab Sessions. i. Read and understand the objectives and the theory of the experiment. ii. Think and plan the working procedures properly for the whole experiment. Make sure you have appropriate table for the data. iii. Complete and submit the pre-lab questions provided. 4.2 Practical Sessions i. Check the apparatus provided. ii. Conduct the experiment carefully. iii. Record all measurements and observations made during the experiment. iv. Keep the work area clean and tidy. Updated: 12/03/2020 v SK015 & SK025 Lab Manual 4.3 Post-lab Sessions i. Explain what has been carried out and discuss the findings of the experiment. ii. Introduce the format of report writing as below: Objective  state clearly Theory Procedure  write concisely in your own words Results/  draw and label diagram if necessary Observation  write in passive sentences about all the Discussion steps taken during the experiment Conclusion  data tabulation with units and uncertainties  data processing (plotting graph, calculation to obtain the results of the experiments and its uncertainties)  give comments about the experimental results by comparing it with the standard value.  state the source of mistake(s) or error(s) if any as well as any precaution(s) taken to overcome them.  answer all the questions given  state briefly the results with reference to the objectives of the experiment Reminder: NO PLAGIARISM IS ALLOWED. Updated: 12/03/2020 vi CHEMISTRY 1 SK015 SK015 & SK025 Lab Manual EXPERIMENT 1 DETERMINATION OF THE FORMULA UNIT OF A COMPOUND Course Learning Objective Demonstrate the correct techniques in handling laboratory apparatus and chemicals when carrying out experiments. (P3, PLO 2, MQF LOD 2) Learning Outcomes At the end of this lesson, students should be able to: i. synthesise a zinc chloride compound. ii. ddetermine the formula unit of zinc chloride. Student Learning Time (SLT) Face-to-face Non face-to-face 2 hour 0 Introduction One of the main properties of a compound is its chemical composition which can be identified by determining the elements present. A quantitative analysis can be used to determine the composition of an unknown compound. Once the composition of the compound is known, it’s formula unit can be determined. For example, a compound containing 0.1 mole of silver and 0.1 mole of bromine will have a formula unit, AgBr. In this experiment, a simple compound composed of zinc and chlorine will be prepared. 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 compound can be calculated and the formula unit can be deduced. Apparatus Chemical reagents Hot plate 6 M HCl Glass rod Zinc powder White tile Crucible tongs 50 mL Crucible Analytical balance Measuring cylinder (10 mL) Procedure 1. Weigh the crucible and record the exact mass. 2. Place approximately 0.25 g of zinc powder into the crucible and determine the exact mass of zinc powder. Updated : 12/03/2020 1 SK015 & SK025 Lab Manual 3. Carefully add in 10 mL of 6 M HCl solution into the crucible containing the zinc powder and stir gently with 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 explosions. 4. If the zinc powder does not dissolved completely, continue adding the acid, 5 mL at a time until all zinc is dissolved. 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 content slowly so that the compound does not splatter during the heating process. 6. Heat the compound gently until it is completely dry. Remove the crucible from the hot plate immediately to avoid the compound from melting. 7. Cover the crucible and allow it to cool to room temperature. Then weigh the crucible and the compound. Record the mass. 8. Reheat the crucible to dry the compound. Let it cool to room temperature and then weigh it again. Repeat the procedure until the difference in mass does not exceed 0.02 g. 9. Determine the mass of zinc chloride from the final weight of the sample (the smallest value). Calculate the mass of chlorine in the zinc chloride. 10. Determine the formula unit of zinc chloride. EXERCISE 1. Explain why the content is not weighed while it is still hot. 2. Explain why the crucible needs to be covered during cooling. 3. Write a balanced equation for the reaction between zinc and hydrochloric acid. Updated : 12/03/2020 2 SK015 & SK025 Lab Manual EXPERIMENT 2 ACID-BASE TITRATION − DETERMINATION OF THE CONCENTRATION OF HYDROCHLORIC ACID SOLUTION Course Learning Objective Demonstrate the correct techniques in handling laboratory apparatus and chemicals when carrying out experiments. (P3, PLO 2, MQF LOD 2) Learning Outcomes At the end of this lesson, students should be able to: i. prepare a standard solution of oxalic acid. ii. standardise 0.2 M NaOH solution. iii. determine the concentration of HCl solution. iv. acquire the correct techniques of titration Student Learning Time (SLT) Face-to-face Non face-to-face 2 hour 0 Introduction Titration is a laboratory technique used to determine the concentration of a solution using another solution with a known concentration. Standards in acid-base titrations One of the solutions involved in a titration is used as a standard solution. The standard solution can be classified as either primary or secondary. A primary standard solution is prepared by dissolving an accurately weighed pure solid of a known molar mass in a known volume of distilled water. A primary standard is used to determine the molarity of the other standard solution, known as 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). The NaOH solution used in titrations need to be standardized because they contain impurities. Solid NaOH is hygroscopic (it absorbs moisture). Thus, it is difficult to obtain its accurate mass. The standardized NaOH becomes the secondary standard and can then be used to determine the concentration of other acids such as HCl acid. Equivalence point and end point An equivalence point is the point in a titration at which the added titrant reacts completely with the electrolyte according to stoichiometry.To detect this equivalence point, an indicator which produces a change in colour is often used. The point at which the indicator changes colour is called the end point. The end point and equivalence point should ideally be the same. Updated : 12/03/2020 3 SK015 & SK025 Lab Manual Chemical equations In this acid-base titration, the neutralisation reactions involved are: H2C2O4(aq) + 2NaOH(aq)  Na2C2O4(aq) + 2H2O(l) . . .(1) HCl(aq) + NaOH(aq)  NaCl(aq) + H2O(l) . . .(2) Apparatus Chemical reagents Burette x M HCl Glass rod 0.2 M NaOH White tile Distilled water Retort stand Phenolphthalein Filter funnel Hydrated oxalic acid, H2C2O4.2H2O 50 mL beaker 25 mL pipette Analytical balance 250 mL conical flask 250 mL volumetric flask 50 mL measuring cylinder Procedure (A) Preparation of standard solution 1. Weigh to the nearest 0.0001 g about 3.00 g of hydrated oxalic acid, H2C2O4.2H2O in a 50 mL beaker. 2. Add approximately 30 mL of distilled water to dissolve the oxalic acid. 3. Transfer the solution into a 250 mL volumetric flask. Rinse the beaker and pour the content into the flask. Add distilled water up to the calibrated mark of the volumetric flask. 4. Stopper and shake the flask to obtain a homogeneous solution. 5. Calculate the concentration of the standard oxalic acid solution. NOTE: Use this solution to standardize the NaOH solution in Part (B). (B) Standardisation of 0.2 M NaOH solution 1. Rinse a burette with a given NaOH solution to be standardized. 2. Fill the burette with the NaOH solution. Ensure there are no air bubbles trapped at the tip. 3. Record the initial burette reading to two decimal places. 4. Pipette 25 mL of oxalic acid solution from Part (A) into a 250 mL conical flask. Add 2 drops of phenolphthalein to the oxalic acid solution. Updated : 12/03/2020 4 SK015 & SK025 Lab Manual 5. Place a white tile underneath the flask so that any colour change can be clearly observed. 6. Titrate the acid with the NaOH solution from the burette. During the titration, swirl the flask continuously. 7. Rinse the unreacted solutions at the inner wall of the conical flask with distilled water. 8. Upon reaching the end point, a temporary pink solution appears but fades when the solution is swirled. Continue titrating until a pale pink colour persists for more than 30 seconds. This is the end point. 9. Record the final burette reading to two decimal places. 10. Repeat the titration three times. 11. Calculate the molarity of the NaOH solution. (C) Determination of the molar concentration of HCl solution. 1. Pipette 25 mL of a given HCl solution into a 250 mL conical flask. 2. Add two drops of phenolphthalein. 3. Repeat steps 5-9 as in Part (B). 4. Calculate the concentration of HCl. EXERCISE Does the addition of water in step 7 (Part B) affect the result of the titration? Explain. Updated : 12/03/2020 5 SK015 & SK025 Lab Manual DATA SHEET EXPERIMENT 2 ACID-BASE TITRATION RESULTS (A) Preparation of standard oxalic acid solution i. Exact mass of hydrated oxalic acid = ii. Moles of hydrated oxalic acid = iii. Molarity of oxalic acid = (B) Standardisationof 0.2 M NaOH solution Burette reading / mL Gross I II III Final reading Initial reading Volume of NaOH used / mL Average volume of NaOH used = Calculate the molarity of the NaOH solution. (C) Determination of the molar concentration of HCl solution Burette reading / mL Gross I II III Final reading Initial reading Volume of NaOH used / mL Average volume of NaOH used = Calculate the molarity of the HCl solution. Updated : 12/03/2020 6 SK015 & SK025 Lab Manual EXPERIMENT 3 DETERMINATION OF THE MOLAR MASS OF A METAL Course Learning Objective Demonstrate the correct techniques in handling laboratory apparatus and chemicals when carrying out experiments. (P3, PLO 2, MQF LOD 2) Learning Outcomes At the end of this lesson, students should be able to: i. standardize the hydrochloric acid solution. ii. determine the molar mass of an alkaline earth metal by back- titration method. Student Learning Time (SLT) Face-to-face Non face-to-face 2 hour 0 Introduction A reactive metal, for example an alkaline earth metal, would readily react with a strong acid such as hydrochloric acid. The general reaction between a metal, M and an aqueous hydrochloric acid, HCl is as follows: M(s) + 2HCI(aq)  MCl2(aq) + H2(g) The molar mass of M can be determined by a back-titration. A back titration is a two-stage analytical technique. The first stage involves the reaction of a metal with an excess amount of acid of a known concentration. In the second stage, the unreacted acid is titrated with a standardized base solution to determine the amount of the remaining excess reactant. In this experiment, the concentration of the acid is initially determined by the normal titration before the reaction with metal M is carried out. M reacts completely according to stoichiometric equation and if the amount of acid used exceeds the amount of metal in terms of equivalence, then the resulting solution would be acidic. The excess acid can be determined by performing back-titration with sodium hydroxide solution. The amount in moles of the reacted metal is determined by comparing the moles of acid before and after the reaction. Apparatus Chemical Reagents Scissors Distilled water White tile Phenolphthalein Pipette filler Filter funnel Dilute hydrochloric acid, HCl Retort stand 0.1 M Sodium hydroxide, NaOH, 50 mL beaker An unknown alkaline earth metal, M 50 mL burette 25 mL pipette 25 mL Analytical balance 250 mL conical flask Abrasive cloth no.3 (36) Aluminium oxide Updated : 12/03/2020 7 SK015 & SK025 Lab Manual Procedure (A) Standardization of HCl solution 1. Rinse a clean burette with 0.1 M NaOH. 2. Fill the burette with 0.1 M NaOH solution. 3. Record the initial burette reading to two decimal places. 4. Pipette 25 mL HCl solution into a 250 mL conical flask. Add 2 drops of phenolphthalein to the acid. 5. Place a piece of white tile underneath the flask. 6. Titrate the acid with the NaOH solution. Swirl the flask continuously. 7. Upon reaching the end point, a temporary pink solution will appear but the colour will fade when it is swirled. Continue titrating until the pale pink colour persists for more than 30 seconds. This is the end point. 8. Record the final reading of the burette. 9. Repeat the titration three times. 10. Calculate the concentration of the HCI solution. (B) Determination of the molar mass of a metal 1. Pipette 25 mL of HCl solution into 2 separate conical flasks. 2. Clean two pieces of metal M, each of approximately 4 cm long, with a piece of abrasive cloth. 3. Weigh accurately the mass of each sample. 4. Cut each sample into smaller pieces. 5. Place the samples separately into the HCl solution. Swirl occasionally until the metal is completely dissolved. 6. Add 2 drops of phenolphthalein. 7. Record the initial burette reading. 8. Titrate the unreacted HCl with the NaOH solution. 9. Record the final burette reading. 10. Repeat titration with the other sample. Updated : 12/03/2020 8 SK015 & SK025 Lab Manual DATA SHEET EXPERIMENT 3 DETERMINATION OF THE MOLAR MASS OF A METAL RESULTS 1. Titration of standard HCl solution Concentration of NaOH = ___________ M Volume of HCl = ___________ mL Burette reading / mL Gross I II III Final reading Initial reading Volume of NaOH / mL Average volume of NaOH = 2. Reaction of metal and HCl Mass of metal (sample I) (g) Mass of metal (sample II) (g) 3. Titration of unreacted HCl Sample I Sample II Burette reading / mL Final reading Initial reading Volume of NaOH (mL) CALCULATION 1. Calculate the molarity of the standard HCl solution. 2. Determine the number of moles of HCl in 25 mL of the standard solution. 3. Calculate the number of moles of the unreacted HCl solution. Sample I: Sample II: Updated : 12/03/2020 9 SK015 & SK025 Lab Manual 4. Calculate the number of moles of the reacted metal. Sample I: Sample II: 5. Determine the molar mass of metal in each sample. Sample I: Sample II: Average molar mass of metal = _______ 6. By comparing the results with elements in the periodic table, determine the metal M. Updated : 12/03/2020 10 SK015 & SK025 Lab Manual EXPERIMENT 4 CHARLES’ LAW AND THE IDEAL GAS LAW Course Learning Objective Demonstrate the correct techniques in handling laboratory apparatus and chemicals when carrying out experiments. (P3, PLO 2, MQF LOD 2) Learning Outcomes At the end of this lesson, students should be able to: i. verify Charles’ Law. ii. determine the molar mass of a volatile liquid. Student Learning Time (SLT) Face-to-face Non face-to-face 2 hour 0 Introduction 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) 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 and a small pressure exerted by the plug. Thus, the pressure of the trapped air is constant. 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= A x h 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 constant pressure. Ideal Gas Equation: By combining the relationships govern by the gas laws, a general equation known as the ideal gas equation can be obtained. Boyle’s Law Volume of a fixed mass of a given gas is inversely proportional to its pressure at constant temperature. V  1 (n, T constant) p Updated : 12/03/2020 11 SK015 & SK025 Lab Manual Avogadro’s Principle All gases of equal volume will contain the same number of molecules at the constant temperature and pressure. V  n (T, P constant) Charles’ Law Volume of a fixed mass of a given gas is directly proportional to its absolute temperature at constant pressure. V  T (n, P constant) Thus, combining the three laws, we get nT Vp The above expression can be written as RnT PV = nRT ...........(1) V = or P This is the ideal gas equation and R is called the gas constant. The number of moles, n, mass n= Molar mass,Mr Therefore, the ideal gas equation can also be written as RT ..........(2) PV = m ( ) Mr Apparatus Chemical reagents Needle Ice Wire gauze Methanol Tripod stand Unknown liquid Rubber band Thermometer Bunsen burner Aluminum foil Beaker (600 mL) Analytical balance Open tube manometer Retort stand and clamp Charles’ law apparatus Conical flask (100 mL) Measuring cylinder (100 mL) Updated : 12/03/2020 12 SK015 & SK025 Lab Manual Procedure (A) Charles’ Law 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 as shown in Figure 4.1. Figure 4.1 Charles’ law apparatus 2. Fill a 100 mL measuring cylinder with tap water. Place the tube and the thermometer into the water until the air column in the tube is immersed. 3. Leave for 5 minutes to ensure that the temperature of the trapped air is equivalent to the temperature of the tap water. 4. Record the temperature and measure the height of the air column. 5. Repeat Steps 2 – 4 using : i. warm water (40 – 50°C) ii. a mixture of ice and water iii. a mixture of ice and 5 mL methanol NOTE: Ensure that the mercury plug does not split into small droplets. Updated : 12/03/2020 13 SK015 & SK025 Lab Manual (B) Determination of the molar mass of a gas 1. Cover a 100 mL conical flask with a piece of aluminium foil and tie it loosely around the neck with a rubber band as shown in Figure 4.2. Figure 4.2 Figure 4.3 2. Prick a tiny hole in the middle of the foil with a needle. 3. Weigh the apparatus accurately. 4. Remove the foil and place 5.0 mL of the unknown liquid into the flask. 5. Replace the foil and tie it with a rubber band. 6. Clamp the neck of the flask and immerse it into a 600 mL beaker containing water as shown in Figure 4.3. 7. Heat the water until all of the unknown liquid in the flask has vaporised. 8. Record the temperature of the water bath when all the unknown liquid has evaporated. 9. Take out the flask immediately by using the clamp. 10. Wipe the outer wall of the flask and the aluminium foil when the flask is cooled. 11. Weigh the flask with the aluminium foil, rubber band and the condensed unknown liquid. 12. 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. 13. Calculate the molar mass of the unknown liquid using the ideal gas equation. Updated : 12/03/2020 14 SK015 & SK025 Lab Manual DATA SHEET EXPERIMENT 4 CHARLES’ LAW AND THE IDEAL GAS LAW (A) Charles’ law TABLE 1 Volume Temperature (Height of gas column) Condition Warm water Tap water Ice-water Ice-methanol 1. Complete TABLE 1. 2. Plot the height of the column, h, against temperature, T, in celsius on a graph paper. Based on the graph, state the relationship between volume and temperature. 3. Extrapolate the line until h = 0, to obtain the absolute zero temperature. (B) Determination of the molar mass of the gas Reading TABLE 2 No Item 1. Mass of flask + rubber band + cover (g) 2. Mass of flask + rubber band + cover + condensed liquid (g) 3. Mass of condensed liquid (g) 4. Temperature of water bath (oC) 5. Barometric pressure (mm Hg) 6. Volume of flask (mL) 1. Complete TABLE 2. 2. Calculate the molar mass of the unknown liquid. Updated : 12/03/2020 15 SK015 & SK025 Lab Manual EXPERIMENT 5 CHEMICAL EQUILIBRIUM Course Learning Objective Demonstrate the correct techniques in handling laboratory apparatus and chemicals when carrying out experiments. (P3, PLO 2, MQF LOD 2) Learning Outcomes At the end of this lesson, students should be able to: i. study the effect of concentration and temperature on chemical equilibrium. ii. determine the equilibrium constant, Kc, of a reaction. Student Learning Time (SLT) Face-to-face Non face-to-face 2 hour 0 Introduction There are two types of chemical reactions, namely irreversible and reversible. A reversible reaction will reach a dynamic equilibrium when the rate of the forward reaction is equal to the rate of the reverse reaction. At this stage, one cannot observe any changes in the system as the concentration of reactants are constant. This does not mean that the reactions have stopped, instead, the reactions are still occurring but at the same rate. The factors that influence chemical equilibrium are: i. concentration ii. temperature iii. pressure (for reactions that involve gases) A change in one of the factors on a system that is already at equilibrium, will cause the reaction to move to the direction that minimizes the effect of change. The direction of the change can be determined by applying Le Chatelier’s Principle. Le Chatelier’s Principle states that if a system at equilibrium is disturbed by a change in temperature, pressure or concentration of one or more components, the system will shift its equilibrium position in such a way so as to counteract the effect of the disturbance. The effect of concentration According to the Le Chatelier’s principle, the change in concentration of any substance in a mixture at equilibrium will cause the equilibrium position to shift in the forward direction or reverse direction to re-attain the equilibrium. Consider a general reaction as follows: A+B C+D If substance 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 re-established. Updated : 12/03/2020 16 SK015 & SK025 Lab Manual On the other hand, if substance 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 re-established. The effect of temperature The effect of temperature on an equilibrium system depends on whether the reaction is exothermic or endothermic. Consider the following system: E+F G + Heat If the forward reaction is exothermic then the heat released is considered as one of the products. Heating the system will cause the equilibrium to shift in the reverse direction so as to reduce the excess heat. Thus, the concentrations of E and F increase while 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 shift in equilibrium through changes in colour o r phases such as precipitation or dissolution. Apparatus Chemical reagents Burette 6 M HCl Ice bath 0.2 M CoCl2 Test tube 2.5 M NaOH Water bath 0.1 M KSCN Pipette (10 mL) 0.1 M Fe(NO3)3 Beaker (100 mL) 0.5 M SbCl3 in 6 M HCl Conical flask (100 mL) Measuring cylinder (10 mL and 100 mL) Procedure (A) The effect of concentration in the formation of thiocyanoiron(III) complex ion The thiocyanoiron(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) + 2SCN- (aq) [Fe(SCN)2]+(aq) (Yellowish brown) (blood-red) 1. Place 2 mL of 0.1 M Fe(NO3)3 solution and 3 mL of 0.1 M KSCN solution in a 100 mL beaker. 2. Add 50 mL of distilled water to reduce the intensity of the blood red solution. Updated : 12/03/2020 17 SK015 & SK025 Lab Manual 3. Place approximately 5 mL each of this solution into four test tubes. (a) To the first test tube, add 1 mL of 0.1 M Fe(NO3)3. (b) To the second test tube, add 1 mL of 0.1 M KSCN. (c) To the third test tube, add 6-8 drops of 2.5 M NaOH. (d) The fourth test tube serves as a control. 4. Tabulate the observations. (B) The Effect of Temperature The reaction between hexaaquocobalt(II) complex ion with chloride ion produces tetrachlorocobalt(II) ion. The equation for the reaction is given below: [Co(H2O)6]2+(aq) + 4Cl-(aq) [CoCl4]2-(aq) + 6H2O(l) (pink) (blue) 1. Place 2 mL of 0.2 M CoCl2 solution into a conical flask. 2. Add 20 mL of 6 M HCl and swirl the flask. 3. A purple solution should form, indicating a mixture of pink and blue. If the solution appears pink, add more HCl; if it is blue, add more distilled water. 4. Divide the purple solution into 3 separate test tubes. (a) Leave one test tube at room temperature. (b) Place the second test tube in an ice bath. (c) Place the third test tube in a water bath at 80 – 90oC. 5. Record the colour of the solution in each test tube. Remove the second and the third test tubes and leave them at room temperature. Observe the change in colour. EXERCISE Determine whether the forward reaction is exothermic or endothermic. Discuss. Updated : 12/03/2020 18 SK015 & SK025 Lab Manual (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 [SbCl 3 ] Procedure 1. Pipette 5.0 mL of 0.5 M SbCl3 in 6 M HCl into a conical flask. 2. Carefully add distilled water from a burette into the conical flask while swirling until a faint white precipitate is obtained. 3. Record the volume of water added. 4. Calculate the value of the equilibrium constant, Kc. EXERCISE Explain why the concentration of pure liquid and solid are excluded from the equilibrium constant expression for a heterogeneous system. Updated : 12/03/2020 19 SK015 & SK025 Lab Manual EXPERIMENT 6 pH MEASUREMENT AND ITS APPLICATIONS Course Learning Objective Demonstrate the correct techniques in handling laboratory apparatus and chemicals when carrying out experiments. (P3, PLO 2, MQF LOD 2) Learning Outcomes At the end of this lesson, students should be able to: i. use various methods to measure the pH of acids, bases and salts. ii. determine the dissociation constant, Ka, of acetic acid. Student Learning Time (SLT) Face-to-face Non face-to-face 2 hour 0 Introduction pH is a measure of acidity or basicity of a solution. pH is defined as the negative logarithm of hydrogen ion concentration, [H+]. pH = -log [H+] ....................(1) The pH scale ranges from 0 to 14. At 25°C, a neutral solution has a pH of 7. An acidic solution has a pH of less than 7 while a basic solution has a pH greater than 7. There are two methods to determine pH in the laboratory. The first method involves the use of indicators such as pH paper and the universal indicator. The second method is using the pH meter. Acids or bases which ionise 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 ionise completely. An example of a weak acid is acetic acid, CH3COOH, and that of a weak base is ammonia, NH3. Consider the ionisation of a weak acid, HA. HA(aq) H+(aq) + A-(aq) ....................(2) The equilibrium constant expression for the above reaction is written as: Ka  [H ][A ] …..................(3) [HA] where [H+], [A-] and [HA] represent the molar concentrations of species that exist at equilibrium. Kais the dissociation constant for acid HA. A similar expression of Kb can be written for weak bases. One of the methods to determine Ka is by adding a weak acid solution to its conjugated base solution. The product of this process is an acidic buffer solution. The conjugated base is obtained from the salt produced using the titration method. Updated : 12/03/2020 20 SK015 & SK025 Lab Manual In this method, a known weak acid, HA is divided into two equal portions, X and Y. The first portion, X is titrated with NaOH solution using phenolphthalein as an indicator to detect the formation of a salt solution. A change in colour, from colourless to light pink, indicates the end point. The equation for the reaction is:- OH-(aq) + HA(aq) A-(aq) + H2O(l) ………………(4) In this reaction, HA reacts with NaOH to form NaA and H O. NaA ionises completely to form 2 + A- and The number of moles of A- formed is the same as the number of moles of HA in Na . the second portion, Y, which has not been titrated. The second portion of the weak acid HA is added to the conical flask containing the salt NaA. In this mixture, the concentration of HA is equal to the concentration of A- from the salt. Since [A-] = [HA], and from Equation 3, + Ka = [H ] + The value of [H ] is obtained by measuring the pH; hence the value of Ka can be calculated. Apparatus Chemical reagents Burette pH paper pH Meter Methyl red Test tube Methyl orange 25 mL pipette Alizarin yellow 250 mL conical flask Phenolphthalein Universal indicator 0.1 M NaCl 0.1 M NH4NO3 0.1 M CH3COONa 0.1 M and 1.0 M NH3 0.01 M and 1.0 M HCl 0.1 M and 1.0 M CH3COOH 0.1 M, 0.2 M and 1.0 M NaOH Procedure (A) Determination of pH of acidic and basic solutions 1. (a) Place 2 mL of the following solutions into separate test tubes. i. 0.01 M HCl ii. 1.0 M HCl iii. 0.1 M CH3COOH iv. 1.0 M CH3COOH v. 0.1 M NaOH vi. 0.1 M NH3 Use pH paper to determine the pH of the solutions. Updated : 12/03/2020 21 SK015 & SK025 Lab Manual (b) Use a pH meter to determine the pH of the following solutions: i. 0.01 M HCl ii. 1.0 M HCl iii. 0.1 M CH3COOH iv. 1.0 M CH3COOH 2. Fill the test tubes with 2 mL of each of the following solution: i. 0.01 M HCl ii. 0.1 M CH3COOH iii. 0.1 M NH3 Add two drops of methyl red to each test tube. Record the observation. Determine the pH range by comparing the colour of the solutions with the chart provided. Repeat step 2 with methyl orange. 3. Fill the test tubes with 2 mL of each of the following solution: i. 0.1 M NaOH ii. 1.0 M NaOH iii. 0.1 M NH3 iv. 1.0 M NH3 Add two drops of alizarin yellow to each test tube. Record the observation. Determine the pH range by comparing the colour of the solutions with the chart provided. (B) Determination of pH of salt solutions 1. Fill the test tube with 2 mL of each of the following solution: i. 0.1 M NaCl ii. 0.1 M CH3COONa iii. 0.1 M NH4NO3 Using pH paper and universal indicator, determine the pH and state whether the salt solutions are acidic, basic or neutral. Updated : 12/03/2020 22 SK015 & SK025 Lab Manual (C) Determination of the dissociation constant of a weak acid, Ka 1. Pipette 25 mL of 0.1 M CH3COOH into two conical flasks, X and Y. 2. Add 2 - 3 drops of phenolphthalein into the conical flask X, and titrate it with 0.2 M NaOH. When the volume of base reaches 10 mL, add the titrant drop by drop. The end point is reached when the solution becomes pink. Record the initial and the final readings of the burette. 3. Mix the solution in step 2 with 25 mL of 0.1 M CH3COOH in the conical flask Y. Determine the pH of this mixture using a pH meter. 4. Calculate Ka from the value of pH obtained in step 3. EXERCISE 1. Calculate the percentage of ionisation of 0.1 M and 1.0 M acetic acid. How does the percentage of ionisation change with its concentration? 2. Refer to the pH value of acetic acid in Part (A). Calculate its Ka and compare this valueto that obtained from Part (C). Updated : 12/03/2020 23 CHEMISTRY 2 SK025 SK015 & SK025 Lab Manual EXPERIMENT 1 RATE OF REACTION Course Learning Objective Demonstrate the correct techniques in handling laboratory apparatus and chemicals when carrying out experiments. (P3, PLO 2, MQF LOD 2) Learning Outcomes At the end of this lesson, students should be able to study the effect of concentration, temperature and catalyst on the reaction rate Student Learning Time (SLT) Face-to-face Non face-to-face 2 hour 0 Introduction The reaction rate is the change in concentration of the reactants or products per unit time. The factors that influence the rate of reaction are temperature, pressure, catalyst, size of particles and concentration of reactants. The rate of a reaction can be studied by observing the change in the chemical or physical properties of species involved in the reaction. The reaction rate is inversely proportional to the time of the reaction, i.e. the faster the reaction occurs, the shorter is the time for the reaction to complete. Apparatus Chemical reagents Glass rod 0.1 M HCl Water bath 10% MnSO Stopwatch Boiling tube 4 Thermometer Pipette (10 mL) 2.0 M H SO Burette (50 mL) 24 Conical flask (100 mL) Laminated white paper with ‘X’ mark 0.02 M KMnO Measuring cylinder (10mL) 4 0.2 M Na S O 22 3 0.25 M H C O 224 Updated: 12/03/2020 24 SK015 & SK025 Lab Manual Procedure (A) The effect of concentration on the reaction rate 1. Place 50 mL of 0.2 M sodium thiosulphate, Na S O using a burette into a 100 mL 22 3 conical flask. Put the conical flask on the white paper with ‘X’ mark. 2. Pipette 10 mL of 0.1 M HCl into the conical flask and immediately start the stopwatch. Stir continuously with a glass rod until the mark is no longer visible and record the time. Note: The ‘X’ mark should be observed from the top of the conical flask. 3. Repeat steps 1-2 with the addition of distilled water to the sodium thiosulphate as instructed in Table 1.1. Table 1.1 Concentration of reactant Volume of Volume of Concentration Volume of Time 1 0.2 M Na S O distilled of 0.1 M HCI (s) water t 22 3 (mL) Na S O (M) solution 22 3 (mL) (s-1) solution 0.00 (mL) 10.00 50.00 10.00 10.00 40.00 20.00 10.00 30.00 30.00 10.00 20.00 40.00 10.00 10.00 4. Calculate the concentration of the sodium thiosulphate solution after the dilution and the value of 1 . t 1 5. Plot a graph of against the concentration of sodium thiosulphate solution. t 6. Based on the graph, state the relationship between the concentration of the sodium thiosulphate solution with time and the rate of reaction. Updated: 12/03/2020 25 SK015 & SK025 Lab Manual (B) The effect of temperature and catalyst on the reaction rate 1. Label 4 boiling tubes as A1, A2, B1 and B2. 2. Place 10 mL of 0.25 M oxalic acid, H C O solution into boiling tubes A1 and A2. 224 3. Fill boiling tubes B1 and B2 with 5 mL of 0.02 M KMnO4 solution. Then add 10 mL of 2.0 M H2SO4 solution to both tubes. 4. Add 5 drops of 10% MnSO solution to B . Stir the mixture. 42 5. Place tubes A1 and B1in a water bath at temperature of 30C for about 3 minutes. 6. While tube A1 is still in the water bath, pour the solutions from tube B1 into tube A1. Start the stopwatch immediately. 7. Record the time taken for the mixture to turn colourless. 8. Repeat Steps 5 - 7 for tubes A2 and B2. 9. Follow Steps 2-7 for the temperatures of 35C, 40C and 50C. Record your results in Table 1.2. Table 1.2 Effect of temperature and catalyst on reaction rate Temperature Without catalyst MnSO4 With catalyst MnSO4 (C ) (A1 + B1) (A2 + B2) 30 t (s) 1 (s-1) t (s) 1 (s-1) 35 t t 40 50 1 10. Plot against the temperature for the mixtures of A1+ B1 and A2 + B2solutions on t the same graph. 11. Based on the graph, deduce the relationship between i. temperature and rate of reaction. ii. catalyst and rate of reaction. Updated: 12/03/2020 26 SK015 & SK025 Lab Manual EXERCISE 1. What is the function of the catalyst in the above reactions? 1 2. What does represent? t Updated: 12/03/2020 27 SK015 & SK025 Lab Manual EXPERIMENT 2 DETERMINING THE HEAT OF REACTION Course Learning Objective Demonstrate the correct techniques in handling laboratory apparatus and chemicals when carrying out experiments. (P3, PLO 2, MQF LOD 2) Learning Outcomes At the end of this lesson, students should be able to: i. determine the heat capacity of a calorimeter. ii. determine the heat of neutralisation of HCl and NaOH Student Learning Time (SLT) Face-to-face Non face-to-face 2 hour 0 Introduction Heat released or absorbed during chemical reactions 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, -q is absorbed by the solution and the calorimeter. rx -q = q + q (1) rxn s c where q = heat absorbed by solution s q = heat absorbed by calorimeter c The heat absorbedby a calorimeter is proportional to the change in temperature. The proportionality constant, C, is known as the heat capacity of a calorimeter. Heat capacity is defined as the amount of heat required to increase the temperature of the calorimeter by 1C. qc = C∆T ……….(2) For a solution, the heat absorbed is proportional to the mass of the solution and the change in temperature. The constant, c, is known as the specific heat capacity of solution per unit mass. The specific heat capacity of a very dilute solution is equivalent to the specif ic heat -1 -1 capacity of pure water ,4.18J g C The mass of the solution can be calculated by assuming the density of the solution is the same as the densityof water. qs = mscs ∆ T …………(3) Heat released can be determined by measuring the temperature before and after the reaction. -q = C ∆T + m c ∆T ……………(4) rxn c ss where final temperature of system – initial temperature of system ∆T = mass of solution heat capacity of calorimeter m= s specific heat capacity of solution C= c c= s Updated: 12/03/2020 28 SK015 & SK025 Lab Manual Apparatus Chemical reagents Pipette (25 mL) 1.0 M HCl Beaker (100 mL) 1.0 M NaOH Thermometer Calorimeter or styrofoam cup Procedure (A) Determination of the heat capacity of a calorimeter 1. Set up a simple calorimeter as shown in Figure 2.1. Figure 2.1 A simple calorimeter (Chang, 2005) 2. Measure the temperature, T1, of an empty calorimeter. 3. Pipette 50 mL of distilled water into a 100 mL beaker. 4. Heat the beaker to a temperature between 50 - 60°C. 5. Pour the hot water into the calorimeter. Close the lid immediately and measure the initial temperature of the hot water, T2. 6. Observe the decrease in temperature every 10 seconds for 2 minutes. Record the temperature that remains constant, T3. 7. Determine the heat capacity of the calorimeter. Updated: 12/03/2020 29 SK015 & SK025 Lab Manual (B) Determination of the heat of neutralisation of 1.0 M HCl and 1.0 M NaOH 1. Pipette 25 mL of 1.0 M NaOH solution into the calorimeter and 25 mL of 1.0 M HCl solution into a beaker. Record the initial temperature of each solution. 2. Without removing the thermometer, lift the lid slightly and quickly pour the HCl solution into the calorimeter. 3. Quickly replace the lid of the calorimeter. 4. Stir the solution and record the maximum temperature reached. 5. Calculate the heat of neutralisation. Updated: 12/03/2020 30 SK015 & SK025 Lab Manual EXPERIMENT 3 ELECTROCHEMICAL CELLS Course Learning Objective Demonstrate the correct techniques in handling laboratory apparatus and chemicals when carrying out experiments. (P3, PLO 2, MQF LOD 2) Learning Outcomes At the end of this lesson, students should be able to: i. arrange Al, Zn, Mg, Fe and Cu in an electrochemical series. ii. determine the Faraday’s constant by electrolysis of CuSO4 solution. Student Learning Time (SLT) Face-to-face Non face-to-face 2 hour 0 Introduction Electrochemistry is a study of the relationship between electricity and chemistry. Generally there are two types of electrochemical cells, namely galvanic and electrolytic cells. A galvanic cell is an electrochemical cell in which redox reaction occurs spontaneously to generate electricity.For a galvanic cell, oxidation occurs at the anode and electrons flow to the cathode where reduction occurs. A standard reduction potential is defined as a reduction potential obtained at a standard condition where the concentration of solution is 1.0 M, the gas partial pressure is 1 atm and temperature is 25 °C. The standard reduction potential values are arranged in a certain order and the list is known as the Standard Reduction Potential Table or the emf Series. The potential difference between the two half cells in an electrochemical cell is called cell potential. The cell potential or the cell voltage at the standard condition can be written as: Eocell = Eocathode - Eoanode The cell potential at non-standard condition can be calculated by using the Nernst equation. Ecell  Eo  0.0592 logQ cell n In this experiment, the cell potential is obtained from the voltmeter reading. By inserting the value and the concentration of the electrolyte in the Nernst equation, the standard c ell potential, Eocell can be determined. An electrolytic cell uses electricity to produce chemical changes in an electrolyte. The cell is made up of two electrodes connected to a battery which functions as a source of direct current. During electrolysis, cations are reduced at the cathode while anions are oxidised at the anode. The amount of substance formed at each electrode can be predicted based on Faraday’s first law. Updated: 12/03/2020 31 SK015 & SK025 Lab Manual Apparatus Chemical reagents Tongs 0.1 M CuSO4 Ammeter 0.1 M ZnSO4 Hair dryer 0.1 M FeSO4 Voltmeter 0.1 M MgSO4 Stopwatch 0.1 M Al(NO3)3 Transformer Zinc strip Sand paper / abrasive cloth Crocodile clips Copper strip Beaker (50 mL) Magnesium strip Analytical balance Salt bridge Iron strip (nail) Measuring cylinder (50 mL) Aluminium strip Carbon rod Saturated KNO3/KCl Note: 1. Clean the electrodes with sand paper / abrasive cloth before use. 2. Ensure that the filter paper to be used as salt bridge is completely soaked in saturated KNO3/KCl solution. Avoid handling the salt bridge with bare hands. Procedure (A) Galvanic cell 1. Clean the metal strips with sand paper / abrasive cloth. 2. Fill a 50 mL beaker with 35 mL of 0.1 M CuSO4 and the other beaker with 35 mL of 0.1 M ZnSO4. 3. Set up the apparatus as shown in Figure 3.1. Zn Cu Salt bridge CuSO4 ZnSO4 Figure 3.1 Galvanic cell 4. Record the cell potential. Updated: 12/03/2020 32 SK015 & SK025 Lab Manual 5. Repeat Steps 1 – 4 by replacing Zn2+/Zn half cell with a (a) magnesium strip in 0.1 M MgSO4 (b) aluminium strip in 0.1 M Al(NO3)3 (c) iron strip in 0.1 M FeSO4 6. Arrange the metals in ascending order of strength as reducing agents. 7. Verify the above order by calculating the standard reduction potential, Eored, of each electrode. (B) Determination of Faraday’s constant 1. Clean a copper electrode with a piece of sand paper / abrasive cloth. 2. Weigh the copper electrode accurately. 3. Set up apparatus as show in Figure 3.2. Fill a 50 mL beaker with 35 mL 0.1 M CuSO4. + DC − Carbon Copper (anode) (cathode) 0.1 M CuSO4 Figure 3.2 An electrolytic cell 4. Complete the circuit by connecting the wires from each electrode to the ammeter and transformer. Set the transformer to supply the direct current with a voltage of 3 V. 5. Run the electrolysis for 15 minutes. 6. Record the ammeter reading and your observation of each electrode. 7. Disconnect the circuit and record the exact time of electrolysis. 8. Dry the copper strip using a hair dryer. 9. Weigh again the copper strip. 10. Calculate the mass of copper deposited. Determine the Faraday’s constant. Updated: 12/03/2020 33 SK015 & SK025 Lab Manual EXPERIMENT 4 REACTIONS OF ALIPHATIC AND AROMATIC HYDROCARBONS Course Learning Objective Demonstrate the correct techniques in handling laboratory apparatus and chemicals when carrying out experiments. (P3, PLO 2, MQF LOD 2) Learning Outcomes At the end of this lesson, students should be able to: i. study the chemical properties of an alkane, alkene and arene. ii. differentiate an alkane from an alkene and arene. Student Learning Time (SLT) Face-to-face Non face-to-face 2 hour 0 Introduction Hydrocarbons are organic compounds that contain only carbon and hydrogen. Alkanes which are also known as paraffins are saturated hydrocarbons. They do not contain double or triple bonds. Hence, alkanes are relatively inert to chemical reactions. Example of alkanes: H H HH HC HCCH H HH Methane Ethane Cyclohexane Alkanes undergo free radical substitution reaction. CH4 + CH2Cl2 CH3Br + HBr Br2 uv Alkenes are unsaturated hydrocarbons with at least one double bond between two carbon atoms. Example of alkenes: H2C CH2 Ethene Cyclohexene Updated: 12/03/2020 34 SK015 & SK025 Lab Manual Alkenes can easily undergo addition reactions at the C=C bond. For example, alkenes undergo hydrogenation and halogenation to form alkanes and dihalides, respectively. +CH3CH=CH2 CH2Cl2 Br2 CH3CHBrCH2Br Alkenes also react with potassium permanganate solution in two different conditions: a. In basic medium to form a diol. H2C CH2 KMnO4/OH- HH + MnO2 room temp HC CH (brown precipitate) OH OH b. In hot acidic medium to form a carboxylic acid. H3C CH3 KMnO4/H+ 2 H3C O C C C H H OH Arenes are aromatic hydrocarbons with stable molecular structures. Example of aromatic hydrocarbons: CH3 Toluene Naphthalene Anthracene Updated: 12/03/2020 35 SK015 & SK025 Lab Manual Although arenes have a very high degree of unsaturation, they are relatively inert towards all addition reactions except at a very high pressure and temperature. Ni + H2 high pressure, 200 °C Arenes undergo electrophilic aromatic substitution reactions in the presence of a Lewis acid catalyst. + Br2 FeBr3 Br + HBr Apparatus Chemical reagents Dropper Toluene Test tube Cyclohexane Rubber band Cyclohexene Labeling paper Dichloromethane Test tube rack 0.01 M KMnO4 Black sugar paper (6 x 12 cm) 4% bromine in dichloromethane Procedure (A) Reaction with bromine in dichloromethane 1. Label 6 dry, clean test tubes, A to F. 2. Place 1 mL of cyclohexane in test tubes A and B, 1 mL of cyclohexene in test tubes C and D, and 1 mL of toluene in test tubes E and F. 3. Wrap test tubes A, C and E with black sugar papers. 4. Add 4 to 5 drops of 4% bromine in dichloromethane into each test tube. 5. Keep test tubes A, C and E in a dark place, and test tubes B, D and F in the sunlight. Leave them for 15 minutes. 6. Record the observations. Updated: 12/03/2020 36

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SK015 Experiment 1: Determination of The Formula Unit of A Compound.

Determination of The Formula Unit of A Compound. In this experiment, we are going to synthesis zinc chloride and then, deduce its formula unit. Students will use analytical balance to weight zinc and mix zinc with concentrated hydrochloric acid in fume cupboard.

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