Hess Law;Peter Jeschofnig, Ph.D.;Version 42-0158-00-01;Review the safety materials and wear goggles when;working with chemicals. Read the entire exercise;before you begin. Take time to organize the materials;you will need and set aside a safe work space in;which to complete the exercise.;Experiment Summary;Students will have the opportunity to measure;temperature changes taking place in a calorimeter;during neutralization reactions and use the;measurements to calculate enthalpy of reaction.;They will illustrate the validity of Hazy Law by;comparing the values of enthalpy of two chemical;reactions.;Objectives;To measure temperature changes taking place in a calorimeter during neutralization;reactions;and use the measurements to calculate enthalpy of reaction.;To compare the enthalpy of two chemical reactions and use these measured values;to illustrate;the validity of Hess Law.;Materials;Materials From: Label or;Box/Bag: Qty Item Description;Student Provides Distilled water;Watch;Coffee cups;Paper towels;From LabPaq 1 Thermometer - Digital;1 Goggles-Safety;4 Cup, Styrofoam, 8 oz;1 Cylinder-25-mL;From Experiment Bag;Hess' Law 2 Ammonia, NH3 (comes as aqueous;ammonia, NH4OH), - 2 M - 10 mL;2 Ammonium chloride, NH4Cl - 2M - 10mL;2 Hydrochloric acid, HCl - 2 M - 20 mL;2 Pipet, Long Thin Stem;2 Sodium hydroxide, NaOH - 2M - 20 mL;Note: The packaging and/or materials in this LabPaq may differ slightly from that which;is listed;above. For an exact listing of materials, refer to the Contents List form included in the;LabPaq.;Discussion and Review;Thermochemistry is the study of the heat energy involved in chemical reactions and;changes of physical state. Nearly all chemical reactions involve the release or;absorption of heat, a form of energy. The burning of any fuel such as gasoline, coal, or;wood is an example of a heat-releasing reaction. Heat energy is called thermal energy;and it is always spontaneously transferred from hotter to colder matter.;The First Law of Thermodynamics is the Law of Energy Conservation. It states that;the total energy of the universe must remain constant. Therefore, all energy transferred;between a system and its surroundings must be accounted for as heat or work.;The standard S.I. unit for heat energy is the joule, J. It takes 4.184 joules, the;equivalent of 1;calorie, to raise the temperature of one gram of water by 1 C. The kilojoule, kJ, is;commonly used in many applications: 1000 joule = 1 kilojoule.;When a chemical reaction takes place in a stable environment where the temperature;and;pressure remain constant, the system defined by the reactants and products either;produces or;releases heat energy.;If the reacting system releases heat energy to its surroundings, a concurrent;increase in;surroundings temperature is observed, and the reaction is exothermic;If the system absorbs heat energy from its surroundings, a decrease in the;surroundings;temperature is observed, and the reaction is endothermic.;A measure of the amount of heat given off or absorbed in any chemical reaction is;called the;enthalpy change or heat of reaction, and is given the symbol H.;When thermodynamic measurements are carried out at standard-state conditions where;the;pressure is constant at 1 atm and the temperature is constant at 25oC, the reaction;enthalpy is;designated as the standard enthalpy change or H. It is important to have;standardized values because the enthalpy of a reaction can vary with different reaction;conditions.;The following reaction for the formation of water from its constituents is exothermic;H2(g) + O2(g) H2O(l), H f = -286 kJ;For every mole of H2O (l) formed at standard-state conditions, 286 kilojoules of heat;energy are;released. When the standard enthalpy change of reaction describes the formation of 1;mol of;compound directly from its elements in their standard states as in this example, the;value of H of is called the standard heat of formation.;To determine the enthalpy change for a given reaction (Hrxn), the summation of the;heats of;formation (H f) for the reactants are subtracted from the summation of the heats of;formation (H f) for the products.;H rxn = [n Hf (products)] - [n Hf (reactants)];Tables containing the standard heats of formation for a number of compounds are;available in the appendices of any general chemistry textbook.;Hess's Law states that if a reaction is the sum of two or more other reactions, the H;for the;overall process must be the sum of the H values of the constituent reactions.;Enthalpy change (H) is independent of the path that a reaction follows to move from;reactants;to products. It only depends on the relative energy difference between the reactant and;product;molecules at constant pressure. Enthalpy change is referred to as a state function due;to its;independent of pathway. Since the enthalpy of a substance is not commonly;determined, the;change in enthalpy when reactants are converted to products is often used to describe;a chemical;or physical process.;The thermal energy absorbed or produced by a chemical process reflects a difference;between;the enthalpy between the reactants and products (H). For example, in the;decomposition of;liquid water into its component elements, H 2 (g) and O2 (g), there are two successive;changes.;First, the liquid water is vaporized. Second, the water vapor decomposes into its;constituent;elements shown below. The H value for this overall process can be determined by;adding the;H values from the equations for each step as shown below.;(1) H2O (l) H2O (g), H 1 = +44 kJ;(2) H2O (g) H2 (g) + O2 (g), H 2 = +242 kJ;(1) + (2) H2O (l) H2 (g) + O2 (g), Hnet = +286 kJ;In order to determine H for the reaction NH3 + HCl NH4Cl in this experiment, H rxn;for the;following two reactions will be measured;1. NaOH (aq) + HCl (aq) H2O (l) + NaCl (aq);2. NaOH (aq) + NH4Cl (aq) NH3 + NaCl + H2O (l);Comparison of the calculated results for different parts of the experiment will verify the;generalization known as Hess's Law of Constant Heat Summation. In this case the;target reaction NH3 + HCl NH4Cl can also be performed directly and the results;compared to reactions 1 and 2.;A Styrofoam coffee cup calorimeter will be used to measure the amount of heat energy;evolved;or absorbed during the chemical reactions of this experiment. A digital thermometer is;used to;measure the change in temperature between the final and initial temperatures of the;solutions.;Unfortunately, it is impossible to have perfect insulation and some of the heat energy will;be lost to the surroundings, including to the material from which the calorimeter is;constructed.;Calibrating the calorimeter before using it to make measurements on an unknown;system usually solves the problem of heat losses. A known amount of heat energy from;a known process is released into the calorimeter system, and the temperature change is;measured. A simple calculation is done to determine the amount of heat energy loss;called the heat capacity of the calorimeter or calorimeter constant. For this experiment it;assumed that the heat capacity of the calorimeter is insignificant and it is ignored.;Another practical problem is that heat energy exchanges do not occur instantaneously;i.e., it takes time for energy to move from a hot object to a cold one. An acceptable;solution to this problem is to obtain a cooling curve for the heat energy exchange in;question and then extrapolate the data back to the exact time that the exchange began.;Below is a sample graph from hypothetical data. Notice that at the time of combining the;two solutions, their starting temperature is 20 oC. Since the starting temperatures are at;room;temperature no initial temperature adjustment is needed. From 0 to 40 seconds the;temperature;rises rapidly to 34.2oC. The temperature then drops gradually 31.1 oC and will continue;to drop.;Usually recording the temperature in 20-20 second intervals for 5 minutes is enough to;provide a;good cooling curve. Extrapolation of these data backward in time determines what the;temperature;at the time of mixing would have been if the temperature of the reaction had been;instantaneous;and the calorimeter had warmed instantaneously. In this example, the temperature at;the time;of mixing determined by extrapolation is 34.3 oC.;Calculations: The equation used to calculate heat gained or lost is;qsolution = (mass of solution) x (specific heat) x T;Density = 1.02 g/mL for all solutions in this experiment;Specific Heat = 4.184 J;T = Final temperature Initial temperature;A small amount of heat is lost to the surroundings which in this case is the calorimeter.;This;heat loss can be accounted for by using a calorimeter constant, c, which can be;determined;experimentally. However, the amount of heat lost to the calorimeter is so insignificant;that it is;often left off, or simply assumed to be 1 J* T. (q cal = c x T).;If a correction was to be made for the heat absorbed by the calorimeter, the heat of the;reaction;qrxn, could be determined by taking the negative of the heat gained by the solution, qsoln;plus that;gained by the calorimeter, qcal;qrxn = -(qsoln + qcal);Once the total thermal energy transfer is known, the enthalpy of reaction can be;determined;using the following equation;H = qrxn /moles NaOH or HCl;Moles of NaOH or HCl can be determined from the equation: M = moles/L;10 mL = 0.01L, 2M = moles/0.01L = 0.02 moles;Exercise 1: Hess Law;Procedure;Part 1: Reaction: HCl & NaOH NaCl + H2O;1. Before beginning, set up data tables similar to the Data Tables 1 & 2 in the Lab;Report Assistant;section.;2. Construct a calorimeter from 2 Styrofoam cups: Trim the lip of one cup and use that;cup as;the top of the calorimeter. Make a small hole in the top so a thermometer can be;inserted, as;shown below. Be careful when inserting the thermometer into the calorimeter since it;has a;pointed tip that could puncture the lower cup if inserted too forcefully. Place the;calorimeter;assembly into an empty coffee cup to help prevent it from tipping over.;Figure 2;3. Use a graduated cylinder to accurately measure 10 mL of 2M HCl. Use an empty;thin-stem;pipet to remove or add drops of HCl so that the meniscus level is on the 10 mL mark.;Pour the;10 mL HCl into the Styrofoam calorimeter. Rinse the thin-step pipet according to this;manuals;instructions on Use, Disposal, and Cleaning of Common Materials.;4. Rinse and dry the graduated cylinder and accurately measure 10 mL of 2M NaOH;using the;same technique in step 2 above. Pour the 10 mL NaOH into another Styrofoam cup and;place;the cup into a second empty coffee cup to prevent it from tipping over.;5. Turn on the digital thermometer and place it into the HCl solution. Wait 5 minutes and;record;the temperature of the solution in Data Table 1.;6. Remove the thermometer, rinse the tip with distilled water, dry it with a paper towel;and;place it into the NaOH solution. Wait 5 minutes and record the temperature of the;solution;in Data Table 1. Remove the thermometer, rinse the tip with distilled water, and dry it;with a;paper towel for future use.;7. Pour the contents of one Styrofoam cup into the second one, combining the two;solutions.;Quickly place the Styrofoam lid on top of the cup containing the combined solutions and;insert;the thermometer through the hole in the lid. Be careful when inserting the thermometer;to;ensure its pointed tip does not puncture the lower Styrofoam cup.;8. Record the temperature every 20 seconds for 5 - 6 minutes and record in Data Table.;9. Graph the data points using an Excel spreadsheet, time in seconds on the x-axis and;temperature on the y-axis. The graph should look similar to the sample cooling curve;below.;10. Place a ruler on the declining temperature portion of the curve and extrapolate to;the 0-line.;Read the extrapolated temperature where the straight line intersects the 0-time line.;This;temperature represents the final temperature of the mixture. Enter this temperature in;Data;Table 1.;11. Dispose of the solution in the calorimeter by flushing it down the drain with water.;Recall that;the solution results from a neutralization reaction and is simply salt water.;12. Rinse all equipment used in preparation for reaction 2. This includes the;calorimeters;graduated cylinders, pipets, etc.;Part 2: Reaction 2: NH4Cl + NaOH NH3 + NaCl + H2O;1. Repeat the Procedures from Part 1, but using 10 mL of 2M NH 4Cl and 10 mL of 2 mL;of NaOH.;2. Dispose of the solution in the calorimeter by flushing it down the drain with water.;3. Rinse all equipment used in preparation for reaction 3. This includes the calorimeters;graduated cylinders, pipets, etc;Part 3: Reaction: NH3 + HCl NH4Cl;1. Repeat the Procedures from reaction 1, but using 10 mL of 2M NH3 and 10 mL of 2;mL of HCl.;2. Dispose of the solution in the calorimeter by flushing it down the drain with water.;3. Rinse all equipment used in preparation for future experiments. This includes the;calorimeters;graduated cylinders, pipets, etc.;Hess Law;Peter Jeschofnig, Ph.D.;Version 42-0158-00-01;LabReportAssistant;This document is not meant to be a substitute for a formal laboratory report. The Lab;Report;Assistant is simply a summary of the experiments questions, diagrams if needed, and;data tables;that should be addressed in a formal lab report. The intent is to facilitate students;writing of lab;reports by providing this information in an editable file which can be sent to an;instructor.;Part 1: Reaction: HCl & NaOH NaCl + H2O;Part 1: Reaction: HCI & NaOH NaCI +H20;Data Table 1: Sample Data;Initial Temperature of HCl - oC;Initial Temperature NaOH - oC;Average Initial Temperature - oC;(extrapolated);Final Temperature of mixture;Change in Temperature of mixture, T;Data Table 2: Sample Data;Time after mixing- seconds;Temperature - C;20;40;60;80;100;120;140;160;180;200;220;240;260;280;300;Part 2: Reaction 2: NH4Cl + NaOH NH3 + NaCl + H2O;Data Table 3;Initial Temperature of NaOH - oC;o;Initial Temperature NH Cl - C;Average Initial Temperature - oC;Final Temperature of mixture;(extrapolated);Change in Temperature of mixture, T;Data Table 4;Time after mixing- seconds;20;40;60;80;100;120;140;160;180;200;220;240;260;280;300;Part 3: Reaction: NH3 + HCl NH4Cl;Data Table 5;Temperature - C;Initial Temperature of HCl - oC;o;Initial Temperature NH - C;Average Initial Temperature - oC;Final Temperature of mixture;(extrapolated);Change in Temperature of mixture, T;Data Table 6;Time after mixing- seconds;20;40;60;80;100;120;140;160;180;200;220;240;260;280;300;Temperature - C;Questions;For A. through E. See the calculations for the Data Tables above.;A. Using the data from your data tables calculate T for all three reactions;B. Calculate the heat loss or gain of the three solution mixtures;C. Use Hess Law and H for the first two reactions;NaOH (aq) + HCl (aq) H2O (l) + NaCl (aq);NaOH (aq) + NH4Cl (aq) NH3 + NaCl + H2O (l);to determine H for this reaction: NH3 + HCl NH4Cl;D. Compare the results of step 3 above with the experimental results of the;NH3 + HCl NH4Cl;E. Use the thermodynamic quantities given below to calculate the theoretical H for this;reaction: NH3 + HCl NH4Cl;Hf for NH3 (aq) = - 80.29 kJ/mol;Hf for HCl (aq) = - 167.2 kJ/mol;Hf for NH4 (aq) = - 132.5 kJ/mol;Hf for Cl- (aq) = - 167.2 kJ/mol;F. What was the H value obtained for NH3 + HCl NH4Cl from Hess Law method?;G. What was the H value obtained for NH3 + HCl NH4Cl experimentally?;H. What was the calculated H value obtained for NH3 + HCl NH4Cl using published;thermodynamic data?;What was the % error of the various methods used? (i.e. comparing the results of the;results of Hess Law method and the experimental results to the calculated value?;J. Name three examples of the practical application for the use of H values.
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