Pearson Education Limited Edinburgh Gate Harlow Essex CM20 2JE England and Associated Companies throughout the world

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2 Pearson Education Limited Edinburgh Gate Harlow Essex CM20 2JE England and Associated Companies throughout the world Visit us on the World Wide Web at: Pearson Education Limited 2014 All rights reserved. No part of this publication may be reproduced, stored in a retrieval system, or transmitted in any form or by any means, electronic, mechanical, photocopying, recording or otherwise, without either the prior written permission of the publisher or a licence permitting restricted copying in the United Kingdom issued by the Copyright Licensing Agency Ltd, Saffron House, 6 10 Kirby Street, London EC1N 8TS. All trademarks used herein are the property of their respective owners. The use of any trademark in this text does not vest in the author or publisher any trademark ownership rights in such trademarks, nor does the use of such trademarks imply any affiliation with or endorsement of this book by such owners. ISBN 10: ISBN 10: ISBN 13: ISBN 13: British Library Cataloguing-in-Publication Data A catalogue record for this book is available from the British Library Printed in the United States of America

3 Evaluation of the Gas Law Constant Report Name Section QUESTIONS 1. What do you feel is the largest source of experimental error in your determination of R? 2. What would be the effect (increase, decrease, or no change) on the calculated value of R of each of the following experimental errors? Explain each answer. a. The liquid level in the eudiometer tube is lower than that in the beaker, but this is not taken into consideration in the calculations. b. The balance used to weigh the magnesium gives a mass that is higher than actual. c. Some H 2 escapes into the beaker during the reaction. 3. Most top-loading balances used in this experiment measure a mass only to the nearest mg (±0.001 g). This significantly affects the calculation of R in this experiment. Explain why this is so. How might the procedure be modified to compensate for this systematic error? 92

4 INTRODUCTION Chemical thermodynamics deals with the energy changes that accompany chemical reactions. Such energy changes are among the factors that determine the following: 1. How fast a chemical reaction takes place that is, the problem of chemical kinetics 2. How complete the reaction will be that is, the position of chemical equilibrium Thermochemistry concerns the energy changes that are manifested as the enthalpy change of reaction, H. H is the heat given off or absorbed by the reaction at constant pressure. A reaction in which heat is lost by the reactants to the surroundings has a negative H and is said to be exothermic; one in which heat is absorbed has a positive H and is said to be endothermic. The general term, enthalpy of reaction, may be classified into more specific categories: 1. The enthalpy of formation is the quantity of heat involved in the formation of 1 mole of the substance in its standard state directly from its constituent elements in their standard states. 2. The enthalpy of combustion is the quantity of heat evolved per mole of a combustible substance, such as carbon or methane, undergoing a reaction with excess oxygen. 3. The enthalpies of solution, vaporization, fusion, and sublimation are concerned with changes in state or solvation of molecules or ions. 4. The enthalpy of neutralization is the heat evolved when 1 mole of water is produced by the reaction of an acid and base. In this experiment, we shall measure the enthalpies of neutralization of HCl and of HC 2 H 3 O 2 solutions with NaOH solution, the enthalpy of solution of NaOH(s), and the enthalpy of neutralization of NaOH(s) with HCl solution. Comparison of calculated results for different parts of the experiment should permit you to verify the generalization known as Hess s law of constant heat summation, which states that a process proceeds through one or several steps, the enthalpy change for the overall process can be calculated by summing the enthalpy changes for the individual steps. Heat measurements are performed by carrying out the reaction in a calorimeter (Figure 1), in which the temperature change and mass of solution are measured. The purpose of the calorimeter is to prevent the gain of heat from the surroundings or the loss of heat to the surroundings. As a chemical reaction proceeds, heat is given off by the reaction and this same heat is gained by the solution. Thus the temperature of the solution increases. If we assume that the heat gained or lost by the calorimeter is insignificant, the heat of the reaction, q rxn, is the negative of the heat that is gained by the solution, q soln. The heat gained by the solution is calculated from the equation: q soln = mc t, where m and t are the mass and temperature change of the solution. The specific heat, C, is the quantity of heat required to raise the temperature of one gram of material (in this case, the material is the solution) by one degree Celsius (J g 1 C 1 ). The specific heat is given for the different parts of this experiment. The heat given off by the reaction and gained by the solution is measured in units of joules. The change in enthalpy, H, of a reaction is the heat evolved per mole of reactant. The H is determined by dividing the q rxn by the number of moles of reactant or product. In this experiment, the number of moles of water produced in Parts B and C equals the number of moles of each reactant. From Laboratory Manual for Experiments in General Chemistry, Ninth Edition, Thomas G. Greco, Lyman H. Rickard, and Gerald S. Weiss. Copyright 2007 by Pearson Education, Inc. Published by Prentice Hall, Inc. All rights reserved. 93

5 The assumption that no heat is gained by the calorimeter when the reaction occurs is not strictly true. However, not including the heat gained by the calorimeter introduces a typical error of 2% or less. This amount of error is well within the experimental error associated with the measurement of the temperature change in the experiment. Figure 1. Simple Calorimeter for Measuring Heat of Reaction DETERMINATION OF TEMPERATURE Three factors make it difficult to determine temperatures quickly in this experiment. 1. The calorimeter is not a perfect insulator, and heat leaks out of it (causing the temperature to drop by about 0.1 C/min). 2. It takes at least 1 min for the calorimeter to reach the same temperature as the mixture within it. 3. In Parts D and E, it takes about 3 min for the NaOH(s) to dissolve. All three effects are evident in this set of student data from Part E of the experiment. Time, min Temp, C mix These data are graphed in Figure 2. Notice at times 1, 2, and 3 that the temperature is slowly rising because the temperature of the HCl is cooler than that of the laboratory. To determine the temperature of the HCl at the time of mixing, we extrapolate these points forward in time (draw the best straight line through them) and obtain 16.8 C at the time of mixing. From time 5 to time 8 the NaOH is dissolving and reacting with the HCl and the temperature rises rapidly to 28.5 C. It then drops back quickly to 28.0 C, probably because it took this last minute to heat the calorimeter. Now the temperature begins to fall slowly and regularly (about 0.1 C/min), and we can extrapolate through these data backward in time to find out what the temperature at the time of mixing would have been if the NaOH had dissolved instantaneously and the calorimeter had warmed instantaneously. Observe carefully in Figure 2 that data are collected for 8 min after mixing in order to obtain good results. Do not stop collecting data too soon. In fact, it is a good idea to plot your data as you collect them. You should make one graph like the one in Figure 2 for each set of data you collect. 94

6 MATERIALS AND CHEMICALS Calorimeters consisting of two 250 ml (7 oz) styrofoam coffee cups nestled together covered with corrugated cardboard squares with center hole for thermometers (graduated in C), spatulas. NaOH pellets; 1.0 M solutions of HCl, NaOH, and HC 2 H 3 O 2 (5 L) Estimated Time: 3-4 hours. This may be too long for a 2-3 hour lab period. Alternatives are to schedule it over 2 periods, have students do only one part and pool class data, or not do Part C. Also, three runs are not essential. SAFETY PRECAUTIONS Review the safety rules. Be particularly careful when handling the sensitive thermometers used for this experiment. The breaking of a thermometer can occur easily and is a serious problem as mercury can be released into the environment. Mercury spills are one of the safety hazards that chemists fear the most. In addition, a broken thermometer presents the hazard of broken glass. If you spill any of the solutions (acids or bases) used in this experiment on yourself or your work area, wash thoroughly with running water and tell your instructor immediately. Figure 2. Temperature-Time Plot PROCEDURE A. Thermometer Calibration Prepare two calorimeters, each similar to the one illustrated in Figure 1, as directed. Compare your two thermometers by immersing them together in water at room temperature for 1 min, and reading the temperature of each as nearly as possible to the nearest 0.1 C. Be careful to avoid parallax in your readings. Always use the 95

7 same thermometer in the calorimeter in which the temperature change occurs, and in all subsequent readings apply any necessary correction to the other, so that the readings of both thermometers will correspond. B. The Enthalpy of Neutralization of HCl(aq) and NaOH(aq) Place 50.0 ml of 1.0 M HCl in one calorimeter and 50.0 ml of 1.0 M NaOH in the other calorimeter. With the lids and thermometers in place, read the temperatures (±0.1 C) for 3 min at 1-min intervals; quickly mix the NaOH thoroughly into the HCl solution, and continue the readings for 4 min at 1-min intervals. Extrapolate the temperatures to the time of mixing for each solution as in Part A, and calculate the enthalpy of neutralization per mole of water produced. (The density of the 0.5 M NaCl produced is 1.02 g/ml, and its specific heat is 4.00 J g 1 C 1.) DISPOSAL Solutions: Add phenolphthalein and neutralize with 1 M HCl (the first disappearance of pink color) or NaOH (first appearance of permanent pink color). Flush the resulting salt solution down the sink with running water. C. The Enthalpy of Neutralization of HC 2 H 3 O 2 (aq) and NaOH(aq) Repeat the procedure of Part B using 50.0 ml of 1.0 M HC 2 H 3 O 2 and 50.0 ml of 1.0 M NaOH. Calculate the enthalpy of neutralization as before. (Assume the same density and specific heat as for NaCl in Part B.) DISPOSAL Solutions: Follow instructions for disposal in Part B. D. The Enthalpy of Solution of NaOH(s) Carefully weigh (to ±0.01 g) about 2.00 g (0.05 mol) of NaOH(s). [Because of its hygroscopic nature, weigh this by difference, in a stoppered 50 or 125 ml Erlenmeyer flask used as a weighing bottle. Your instructor will tell you the approximate number of NaOH(s) pellets required to assist in estimating the mass needed. Be sure to clean up any spilled NaOH(s). This solid absorbs water from the air and forms a slippery solution that is also corrosive.] Place 50.0 ml of distilled water in your calorimeter. With the lid and thermometer in place, read the temperature (0.1 C) for 3 min at 1-min intervals; then add the NaOH(s), replace the lid and thermometer, and gently swirl the mixture and stir it carefully with the thermometer to dissolve the NaOH as quickly as possible. At the same time continue the temperature readings at 1-minute intervals for at least 9 min total (at least three readings after the maximum temperature reading is attained). Because of the time required for solution and complete mixing, the proper estimate of the temperature for complete solution at the time of mixing is more difficult. Make your best estimate based on extrapolations of temperatures before and after mixing as explained in the section on the determination of temperature on page 96. A plot of your data is essential. Calculate the heat of solution per mole NaOH(s) to form a 1.0 M NaOH solution. (Note that you have about 52 g of solution. The specific heat of 1.0 M NaOH is 3.90 J g 1 deg 1.) DISPOSAL NaOH solution: Add phenolphthalein and neutralize with 6 M H 2 SO 4 or HCl (the first disappearance of pink color). Flush the resulting salt solution down the sink with running water. E. The Enthalpy of Reaction of HCl(aq) and NaOH(s) Carefully weigh about 2.00 g (to 0.01 g) or slightly less of NaOH(s). Use a stoppered Erlenmeyer flask as in Part D. Measure about 55 ml of 1.0 M HCl (which provides a small excess of HCl to react with all the NaOH used) in a 100 ml graduated cylinder, and dilute this to ml, as precisely as you can. Transfer this completely to your calorimeter and, with the lid and thermometer in place, take temperature readings once each minute for 3 min. Then add the NaOH(s), replace the lid and thermometer, gently swirl the mixture and stir it carefully with the thermometer to dissolve the NaOH as quickly as possible, and at the same time continue temperature readings for 9 min total at 1-min intervals (at least three readings after the maximum temperature 96