PLEASE REORD ALL DATA DIRETLY INTO YOUR LAB NOTEBOOKS Introduction Heating a substance is one of the simplest processes carried out in the chemical laboratory, and is usually accompanied by a rise in the temperature of the substance being heated. The amount of heat energy required to raise the temperature by 1 depends both upon the amount (mass) of substance and the identity (chemical composition) of the substance. The heat capacity, symbolized by, is the ratio of the heat added (q) to the observed temperature rise () q q hemical reactions are often accompanied by the release or absorption of heat. For example, when hydrogen gas is combined with oxygen gas and heated the reaction produces an enormous amount of heat. Reactions that produce (release) heat are said to be exothermic. Reactions that absorb heat are said to be endothermic. The heat released in a chemical reaction is often determined by measuring the temperature change of the material surrounding the chemical reactants and products. This works because the amount of heat associated with a reaction is equal to the amount of heat that is either transferred to or from its surroundings. The study of heat associated with chemical reactions is called thermochemistry. The measure of temperature changes associated with chemical reactions is called calorimetry. To the right is a diagram of a simple calorimeter, a device in which the heat of a reaction is measured. The calorimeter can be as simple as a styrofoam cup and a thermometer. In a calorimeter the reactants are placed into the container and allowed to react. The reactants and the products of a chemical reaction, are called the system. The materials surrounding the system are called the surroundings. As the reaction proceeds, the temperature of the liquid in which the reaction is occurring changes because heat is transferred between the system and the surroundings. We can measure the change in enthalpy ng the process. We use the system as the reference and use the following sign convention: Heat is absorbed by the system H > 0 endothermic Heat is released by the system H < 0 exothermic If the reaction is exothermic then the temperature of the solvent (surroundings) will increase. If the reaction is endothermic then the temperature of the solvent (surroundings) will decrease. The change in temperature of the solvent can then be used to determine the amount of heat transferred: H = m * s *T where H is the change in enthalpy for the reaction (heat of reaction), m is the mass of the solution in which the reaction is occurring, s is a constant called the specific heat capacity, and is the change in
temperature of the solution as the reaction takes place. The specific heat capacity depends on the material used. Values of s are given below for several materials. Material s c J/(g K) J/(mol K) s c Material J/(g K) J/(mol K) Bi 0.1221 25.52 Zn 0.3886 25.40 Pb 0.1276 26.44 o 0.4210 24.81 Au 0.1290 25.42 Ni 0.4440 26.07 Pt 0.1326 25.86 Fe 0.4494 25.10 Hg (l) 0.1395 27.98 Ti 0.5223 25.02 Sb 0.2072 25.23 a 0.6315 25.31 I 2 (s) 0.2145 54.44 Si 0.7121 20.00 Sn 0.2274 26.99 K 0.7565 29.58 d 0.2311 25.98 Al 0.9025 24.35 Ag 0.2350 25.35 Mg 1.0238 24.89 Se 0.3212 25.36 Na 1.2284 28.24 Ge 0.3216 23.35 Li 3.5609 24.77 u 0.3846 24.44 Water 4.184 75.40 Experimental Procedure The experimental procedure has three parts. You first obtain the heat capacity of the calorimeter by adding measured portions of hot and cold water. This value is used for the second part of the experiment in which you are given a sample of an unknown metal. This sample must be weighed, heated, and placed in the calorimeter. The observed temperature rise is used to calculate the specific heat capacity of the metal. Part I: alibration of the Temperature Probe lick ALIBRATION OF THE TEMPERATURE PROBE. There is a written set of instructions with pictures that will take you through the steps of calibrating your temperature probe. THE WATER THAT YOU USE FOR ALIBRATION SHOULD BE NEAR THE RANGE OF THE DATA THAT YOU WILL OLLET, SO USE OLD/OOL WATER AROUND 20 DEGREES AND WARM WATER AROUND 40 DEGREES. WARM WATER AN BE OBTAINED FROM THE TAP BY THE BENH WHERE THE LAB MATERIALS AN BE FOUND. Part II: Specific Heat apacity of a Metal 1.) Dry out your calorimeter. 2.) Add some cool water (about 20 ) to your calorimeter. Use about 50 g of water, but record exactly (to 4 decimals) the mass of water you use. 3.) Obtain about 15 g of an unknown metal. Record the mass of the metal that you actually obtain, as well as which unknown you re using. 4.) Put your metal in a large test tube, and put the test tube in a beaker of water. Heat the water on a hot plate to approximately 90. Be careful to place the test tube in the beaker in such a manner that the water cannot splash into the test tube. Be certain that the metal sample is in the hot water for at least 5 minutes. 5.) Extend the time of the run by pressing ontrol D and putting in 500 seconds.
6.) Proceed to Part III Part III: Heat of Solution 7.) Nest two clean, dry Styrofoam cups together inside a clean, dry 400 ml beaker. This is your calorimeter. 8.) Put about 50 g of deionized water into the calorimeter. Record the actual mass of water in the calorimeter. 9.) Let the filled calorimeter stand for at least 4 minutes, to allow everything to come to room temperature. 10.) Weigh out approximately 2 g of potassium nitrate (KNO 3 ) and record its mass to the nearest 0.0001 g. Grind the salt finely so it will dissolve uniformly. 11.) Observe the temperature of the calorimeter carefully. Run at least 60 seconds of baseline for the water in the calorimeter. For the dissolution of salts, one of two temperature trends can occur: the temperature may either rise quickly and then fall slowly or fall quickly and then rise slowly. Record the maximum or minimum observed temperature according to the trend observed. 12.) Repeat steps (7.) 11.), using approximately two grams of anhydrous sodium carbonate salt (Na 2 O 3 ) and a fresh portion of water. Part IV: Heat of Neutralization 13.) Assemble a clean, dry calorimeter. 14.) Add 25 ml of sodium hydroxide solution (NaOH) to the calorimeter using a graduated cylinder. over the calorimeter and let it sit for at least 2 minutes with its temperature probe in place. An excess of NaOH is being used in this experiment 15.) Using a graduated cylinder, transfer about 25 ml of hydrochloric acid (Hl) to a dry 100 or 150 ml beaker. Record the exact concentration of the acid, and exactly how much you use. 16.) Run at least 60 seconds of baseline of the temperature of the NaOH. 17.) Use a thermometer to record the temperature of the acid in the beaker to the nearest 0.1. 18.) Quickly, but carefully, transfer all the acid from the beaker into the NaOH in the calorimeter. Gently stir the contents, being careful to hold only the rim of the calorimeter. 19.) arefully observe the temperature of the calorimeter. The temperature may either rise quickly and then fall slowly or it may fall quickly and then rise slowly. Record the maximum or minimum observed temperature. Again, do this by extrapolation. 20.) Repeat steps 1 6, adding 25 ml of acetic acid (H 3 OOH) to 25 ml of 1.10 M NaOH previously placed in the calorimeter. Be sure to thoroughly rinse and dry the calorimeter between runs. Make sure you know how much of each solution you use, as well as the concentration of each solution.
21.) Write the balanced equations for both reactions, and the net ionic equations for both reactions. Part II: Specific Heat apacity of a Metal (continued) 20. You are now ready to add the hot metal to the water in the calorimeter; first you need to record the temperature of the water bath (This will be the T im for the hot metal); pour the hot metal into the calorimeter and allow the system to equilibrate for at least 2-4 more minutes. Do not delay while transferring the metal to your calorimeter. There are consequences for delay, which you will figure out in the post laboratory questions. As soon as the metal is added, swirl the calorimeter continuously to allow through mixing of the water. 21. You may need to fit a straight line to the final linear part of your data. lick and drag the mouse over a linear segment of the data from the end of your run to allow this section of the line to be measured for slope and intercept. With the line segment selected, click on Analyze and select Linear Fit or click on the Linear Fit icon (ask your TA). 22. To Print, click on File and then on Print Graph. lick on the footer box and add the names and other pertinent information. It is recommended that you add the experimental run (e.g. Run #4) and initials of yourself and your partner. DO NOT THROW YOUR OFFEE UPS AWAY! RINSE OUT AND DRY. DATA ANALYSIS Part II: Specific Heat apacity of a Metal Analyze the calorimeter temperature versus time data as you did in Part II. Again extrapolate to find T f and obtain T, the temperature increase. Use your value for to calculate the heat transfer (q ) to the calorimeter: q m s H2 O Your TA will provide you with the c value. Remember that the heat that the metal lost is equal in magnitude to the heat the calorimeter gained, but opposite in sign. q q H Use T f to find T H, the temperature decrease of the metal sample. The heat transfer (q H ) from the metal sample of mass m M is exactly equal in magnitude (but opposite in sign) to the heat gained by the calorimeter: q H m M s M H Insert your experimental values and solve for the specific heat capacity of the metal.
The molar heat capacity (c) of a substance is related to its specific heat capacity through the molar mass: c s M Use the Dulong and Petit value of 26 J/(mol * K) for c and your experimental value for s to obtain the molar mass, and hence the atomic mass M of your unknown metal. Part III: Heat of Solution The temperature change of the calorimeter is obtained by comparing its initial temperature (T i ) with the final temperature (T f ) obtained by extrapolation of your data: T f i The amount of heat absorbed during the dissolution can be calculated from this temperature change using: q m s solution solution where is the previously calculated heat capacity of the calorimeter and the mass of the solution is given by m (H2O) + m (salt). The specific heat capacities of the KNO 3 and Na 2 O 3 Remember that the amount of heat absorbed by the calorimeter and water is equal in magnitude to the amount of heat generated by the dissolution of the salt, but opposite in sign. q q rxn Divide q rxn by the mass of the salt to obtain the specific heat of solution; divide q rxn by the number of moles of salt to get the molar heat of solution. Part IV: Heat of Neutralization Find the total temperature change () of the calorimeter as in Part III. Assume a density of 1.00 g/ml to find the mass (m) of the solution in the calorimeter. Assume the specific heat capacity of the solutions to be equal to that of water. Use the known heat capacity ( ) of the calorimeter to calculate the amount of heat evolved during the reaction: q m solution s solution alculate the number of moles of acid contained in 25 ml of the 1.00 M acid solution. Divide the heat of neutralization by this number. This is the molar heat of neutralization of the acid. alculate this quantity for both acids, and include them with your lab report.