Chem1A – Experiment F: Specific Heat 1

Name __________________________________ Date ____________________ Grade ____

Pre-Lab Questions: Specific Heat

MUST be completed before an experiment is started. Show all work and be sure to include units.

Q1. What is the difference between heat capacity and specific heat?

Q2. What is a calorimeter and why is it used in thermochemistry?

Q3. Why do you need to measure the specific heat of the calorimeter?

Q4. What happens to the temperature of the metal when it is placed in the water in the calorimeter? What happens to the temperature of the water?

Q5. What equation will be used to determine the specific heat of the metal?

Chem1A – Experiment F: Specific Heat 2

Experiment F: Specific Heat Purpose The purpose of this experiment is to understand heat exchange in physical and chemical processes using a closed-cup calorimeter.

Introduction Heat is a form of energy, often called thermal energy. Energy can be transformed from one form to another (for example, electric energy to heat and light in a light bulb), but it cannot be created or destroyed; rather, energy is conserved. The higher the temperature of a material, the more thermal energy it possesses. In addition, at a given temperature, the more amount of a given substance there is, the more total thermal energy the material possesses. On an atomic level, absorbed heat causes the atoms of a solid to vibrate, much as if they were bonded to one another through springs. As the temperature is raised, the energy of the vibrations increases. In a solid, this is the only motion possible. In a liquid or gas, absorbed heat causes the atoms in the molecule to vibrate, and the molecule to both rotate and move from place to place. Because there are more “storage” possibilities for energy in liquids and gases, their heat capacities are larger than solids. Heat capacity is the amount of heat (q) required to raise the temperature of a sample of that material by 1˚C. Heat capacity is an extensive property of matter, meaning that its value varies with sample size.

𝐻𝑒𝑎𝑡  𝑐𝑎𝑝𝑎𝑐𝑖𝑡𝑦 =  𝑞∆𝑇

In order to account for the sample size, specific heat of a substance is measured, which is a slight variation to the heat capacity. Specific heat, csp, is the amount of heat required to change the heat content of exactly 1 gram of a material by exactly 1°C. [Note: some textbooks use different conventions for abbreviating specific heat and heat capacity. In this lab, specific heat is a lowercase ‘c’ with subscript ‘sp’.]

𝑐!" =  𝑞

𝑚 • ∆𝑇

Table 1 gives specific heat values of some of the most common substances.

Table 1: Specific Heats of Some Common Substances

Substance Specific Heat in J/(g•K)

Water 4.184 Wood 1.76

Aluminum 0.902 Concrete 0.88

Glass 0.84 Iron 0.451

Nickel 0.444 Copper 0.385

Zinc 0.385 Cadmium 0.232

Tin 0.222 Lead 0.129

Compare the specific heat of iron and wood. Because the specific heat of wood is about four times as much as that of iron, it takes about four times as much heat to raise the temperature of wood than iron. This can be verified by leaving a piece of wood and a piece of iron out on a hot sunny day. The sun gives off heat, which will be absorbed by wood and iron equally. However, since wood has a greater specific heat value, it is able to absorb more heat before its temperature rises. Therefore, you will feel that the piece of iron is hotter than the piece of wood when it is left out in the sun for equal amount of time.

Chem1A – Experiment F: Specific Heat 3

To measure specific heat in the laboratory we focus on one area of thermodynamics, called calorimetry, which is a technique used to calculate heat flow into and out of materials. The basic idea of the technique is quite simple. The matter we are interested in studying is placed into a container called a calorimeter, which is a well-insulated container used in measuring energy changes. The calorimeter is insulated to reduce the loss or gain of energy to or from the surroundings. Heat always flows from materials at a higher temperature to materials at a lower temperature. When two materials, each initially at a different temperature, are placed in contact with one another, heat always flows from the warmer material into the colder material until the two are at the same temperature. If we assume no loss of heat to the surrounding environment, then the heat gained by the initially colder material must equal to the heat lost by the initially warmer material: Eqn 1: -(heat lost)hot object = (heat gained)cold object From the law of conservation of energy, Eqn 2: qhot + qcold = 0 The heat quantities in this equation would be calculated as q = (specific heat) x (mass of the object) x (the temperature change of the object) Substituting this expression for q in equation 1, we get: Eqn 3: -[csp•m•ΔT]hot = [csp•m•ΔT]cold When a measurement is made in a calorimeter, we must account for the fact that the calorimeter itself can release heat into or absorb heat out of the process while our measurement is being recorded. In our experiment, since same calorimeter will be used for all parts of the experiment, we can determine the amount of heat released or absorbed by the calorimeter once and use throughout the calculations. Temperature values will be recorded in ºC, but the calculations will be done in Kelvin. Note that the numerical value of ΔT in ºC and in K is the same. There are two parts to the experiment. In part A, you will determine the heat absorbed by the cup and specific heat of an unknown metal. In part B, you will determine heat of solution and the heat of reaction.

Chem1A – Experiment F: Specific Heat 4

Part A: Specific Heat

In Part A of this experiment, you will determine the specific heat of an unknown metal sample by using a calorimeter. To do this, first you must obtain the amount of heat that the cup and the thermometer will absorb per Kelvin. This will be done in step I of the specific heat determination. Part A - Step I: Measuring Heat Exchange to the Calorimeter In step I, you will pour a measured amount of hot water into a measured amount of cold water contained in a well-insulated coffee-cup calorimeter with a thermometer. The calorimeter and thermometer will absorb some heat. If the mass of calorimeter and the thermometer can be written as masscal, and the specific heat of calorimeter and the thermometer can be written as csp(cal), then equation (3) can be written as: Eqn 4: -[csp(water) x mass hot water x ΔThot water] = [(csp(water) x mass cold water x ΔTcold water) + (csp(cal) x mass cal x ΔTcold water)] Since the same calorimeter will be used for the entire experiment, we can use its heat capacity, Heat Capacitycup, for the heat absorbed by the cup and the thermometer. (csp(cal) x mass cal = Heat Capacitycup) Thus, equation (4) becomes: Eqn 5: -[csp•m hot water•ΔT]hot water = [(csp•m cold water + Heat Capacitycup)•ΔT] cold water For step I, since the specific heat of water is 4.184 J/(g·K), equation (5) becomes: Eqn 6: -[{4.184 J/(g·K)}•mhot water •ΔT]hot water = [({4.184 J/(g·K)}•mcold water + Heat Capacitycup)•ΔT]cold water Part A - Step II: Specific heat of an unknown metal In step II, the metal sample will be heated to a high temperature then placed into a calorimeter containing a known quantity of water at a lower temperature. Having measured the mass of the water in the calorimeter, the temperature change of the water (ΔT), and knowing the specific heat of water (4.184 J/g °C) the heat gained by the water (lost by the metal) can be calculated. Eqn 7: -[csp•mmetal•ΔT]hot = [{(4.184 J/(g·K)) •mwater + Heat Capacitycup}•ΔT]cold You will solve the equation (7) for csp, the specific heat of the unknown metal.

Chem1A – Experiment F: Specific Heat 5

Part B. Heat of reaction When a chemical reaction occurs in water, there is an exchange of heat between the reaction and the solvent, water. In other words, the reaction releases (or absorbs) heat into the ‘mixture’ containing the reactants and water. Eqn 8: qrxn + qmixture = 0 Equation (8) becomes Eqn 9: qrxn = -[{4.184 J/(g·K)} •mmixture + Heat Capacitycup]•ΔT [Note: the ‘mixture’ mass is the total mass.] Since the experiment is conducted at constant atmospheric pressure, the heat flow associated with the reaction mixture is also equal to the enthalpy change, ΔH, for the reaction. By measuring the mass of the reaction mixture, and by observing the temperature change that the mixture undergoes, you can find ΔHrxn. If the temperature of the mixture goes up (positive ΔT), heat has been given off by the reaction, so the reaction is exothermic, qrxn is negative and ΔHrxn is negative. If the temperature of the mixture goes down (negative ΔT), heat has been absorbed by the reaction, so the reaction is endothermic, qrxn is positive and ΔHrxn is positive.

Part B - Step I: Heat of solution: One of the most simple reactions that can be studied in solution occurs when a solid is dissolved in water. Example: the solvation of sodium hydroxide in water:

Eqn 10: NaOH (s) à Na+ (aq) + OH- (aq) ΔH = ΔHsoln

When this reaction occurs, the temperature of the solution becomes much higher than that of the NaOH (s) and water that were used. If we dissolve a known amount of NaOH in a measured amount of water in a calorimeter, and measure the temperature change that occurs, we can use equation (10) to find ΔHsoln. Noting that ΔHsoln is directly proportional to the amount of NaOH used, you can calculate ΔHsoln of either a gram or a mole of NaOH. In this step of the experiment you will measure ΔHsoln for an unknown ionic solid. Part B - Step II: Heat of reaction: When solutions are mixed together, chemical reactions can occur. A very common reaction is a neutralization reaction when an acid is mixed with a base. In Part B – Step II of this experiment, you will measure ΔHrxn for a neutralization reaction between HCl (aq) and NaOH (aq). The heat effect is quite large, and is the result of the reaction between H+ ions in the acidic solution with OH- ions in the basic solution. Eqn 11: H+ (aq) + OH- (aq) à H2O (l) ΔH = ΔHrxn

Chem1A – Experiment F: Specific Heat 6

Experimental Procedure Safety Considerations for this experiment: Do not fill water to the top in your hot water bath. Hot water will spill over and could cause severe burns. Depending upon the height of your beaker, check with your instructor about the safe level to work with. To avoid breakage, carefully drop the metal piece into the beaker containing test tubes. Do not touch the hot metal piece with bare hands. Use tongs. Do not leave the hotplate or the hot water bath unattended. Be sure to wear your safety goggles and lab apron at all times. Clean all glassware used for the experiment before leaving the laboratory. Waste must be disposed off in appropriate containers. Wash your hands before leaving the lab room. Experiment This experiment will involve: - Using a closed-cup calorimeter to measure specific heat. - Identify a piece of an unknown metal during a physical change by measuring its specific heat and comparing to the standard values. - Calculate heat of solution and heat of reaction for chemical reactions.

Procedure Read all temperatures to the nearest 0.1°C. Use the balances to obtain masses to ± 0.01 g. Part A. Specific heat 1. Set a 400-mL beaker containing about 350 mL of tap water as shown in Figure 1. This will serve as water bath for

your experiment. Turn on the hot plate to medium high so that the water bath will be heating while you are preparing the test tube for the experiment.

Figure 1: Setup of hot water bath

2. Get an unknown metal block. Record the number of the metal in Part A(II) of data table. Dry the metal block onto a paper towel. Immerse the metal block in the hot water bath as shown in Figure 1.

3. You will need two large test tubes and stoppers for Part A(I). Into each test tube, add 30 ml of deionized water. Loosely stopper them, and place them into the hot water bath.

4. When the water is boiling, turn the heat down to maintain a gentle boil in the beaker. The test tubes should reach the temperature in the water bath after 10 minutes of boiling. Assume that the temperature of the boiling water is 100.0°C.

Chem1A – Experiment F: Specific Heat 7

Part A. Specific heat - Step I: Measuring Heat Capacitycup

1. Place 70 ml of deionized water into the coffee cup. Weigh the cup with lid and record the mass in Part A(I) data table. Record the temperature of the water using a thermometer.

2. Using test tube holder, pour the hot water from the test tube into the cup and quickly cover it with the lid. 3. Insert the thermometer into the calorimeter through the wide opening on the lid. Do not try to push the thermometer

through the small opening in the middle as shown in Figures 2(a) and 2(b) below:

(a) (b)

Figure 2: Proper use of thermometer in a coffee cup. (a) Thermometers should go into the cup where there is the widest hole on top. (b) Never put a thermometer into the center hole.

4. Stir with the thermometer and record the highest temperature reached. Now weigh the cup and record the mass as the

final mass of cup and water. 5. You will run two trials for Part A (step I) to see how consistent your data is. For the second trial, repeat steps 1-4 in

this step. When done, empty the water from the cup into the sink.

Part A. Specific heat, Step II: Measuring csp of the Metal

1. You will do just one trial for Part A(II). Add 100 ml of deionized water to the cup. Weigh the cup with the water and lid. Record the mass in Part A(II) of data table.

2. Record the temperature of the water in the cup. 3. Quickly transfer the metal block from the water bath to the cup, and cover the cup immediately. Stir the cup with the

thermometer and record the highest temperature reached. 4. Now weigh the cup with water and metal block. 5. Pour the water out of the cup into the sink. Blot the metal dry with paper towel, and place it in its container before

returning the metal block to the instructor.

Correct   Incorrect  

Chem1A – Experiment F: Specific Heat 8

Part B: Heat of Reaction

In this part, energy changes in chemical reactions will be observed and the associated temperature changes will be recorded for calculation of heat of solution and heat of reaction for the reactions.

Step I: Heat of Solution:

1. Place 50 mL deionized water in the cup. Weigh the cup with the water and lid. Record the mass in Part B (part I) of

data table. Record the temperature of the water in the cup.

2. Get an unknown ionic solid sample. Record the number of the unknown in Part B (part I) of data table. In a small

beaker, weigh about 2.5 g of the solid compound. Quickly transfer the compound to the cup, and cover the cup

immediately. Stir the cup continuously with the thermometer with occasional swirling.

3. Record the highest temperature reached. Check to make sure that all of the solid is dissolved.

4. Weigh the cup with lid and solution, and record the mass in Part B (step I) of data table.

5. Pour this solution in the WASTE Container.

6. Return the unknown sample bottle to the instructor.

Step II: Heat of Neutralization Reaction:

1. In a graduated cylinder, measure 12.5 mL of 1.00 M HCl (aq). Pour this acidic solution into the calorimetric cup.

Measure the temperature of the solution and record it in Part B (step II) of data table. Rinse and dry the thermometer.

2. Rinse the cylinder with distilled water. Then measure 12.5 mL of 1.00 M NaOH (aq). Measure the temperature of

this basic solution and record it in Part B (step II) data table.

3. Place the clean and dry thermometer in the cup. Quickly pour the NaOH solution into the cup containing HCl

solution. Cover the cup immediately. Stir the cup continuously with the thermometer with occasional swirling.

Record the highest temperature reached.

4. When you have completed the experiment, pour the solution in the WASTE container. Rinse the calorimetric cup and

thermometer with water, and return them to the instructor.

Chem1A – Experiment F: Specific Heat 9

Name ________________________________________ Date _____________________ Grade ____

Data Sheet for Experiment: Specific Heat PART A: Specific Heat Step 1: Measuring Heat Capacitycup Trial 1 Trial 2 1) Mass of the coffee cup and lid ___________ g ___________ g (assumed to be constant during experiment) 2) Mass after addition of 70 mL water ___________ g ___________ g

3) Initial temperature of water in the cup ___________ °C ___________ °C

4) Final temperature of water in the cup after hot water is added ___________ °C ___________ °C

5) Final mass of cup, lid, and water ___________ g ___________ g

Calculations

6) Mass of hot water ___________ g ___________ g

7) Initial temperature of the hot water ___________ °C ___________ °C

8) ΔT for the hot water (a negative value) ___________ °C ___________ °C

9) Mass of cold water ___________ g ___________ g

10) ΔT for the cold water ___________ °C ___________ °C

11) Heat Capacity of the calorimeter ___________ J/K ___________ J/K

Show your setup, and include appropriate units and significant digits.

12) Average value of Heat Capacitycup: ___________ J/K

Chem1A – Experiment F: Specific Heat 10

Step II: csp of unknown metal 1) Mass of the coffee cup, lid, and 100 mL water ___________ g 2) Initial temperature of metal (same as hot water bath) ___________ °C 3) Initial temperature of water in the cup ___________ °C 4) Final temperature of water in the cup after hot metal is added ___________ °C 5) Final mass of cup, lid, water, and metal ___________ g Calculations 6) Mass of cold water in the cup ___________ g 7) ΔT for the water ___________ °C 8) Mass of metal used ___________ g 9) ΔT for the hot metal (a negative value) ___________ °C 10) csp of the unknown metal ___________ J/(g·K) Show your work for the csp calculation, and include appropriate units and significant digits. (In this calculation, use the average value of Heat Capacitycup from part I.)

11) Metal ___________ csp (book) ___________ J/(g·K)

(Find this value in a reference or textbook.)

12) % Deviation = csp(experimental) - csp(actual)

csp(actual)

x 100 = __________ %

Chem1A – Experiment F: Specific Heat 11

PART B: Heat of Reaction

Step I: Heat of solution of an ionic salt 1) Mass of the coffee cup, lid, and 50 mL water ___________ g 2) Initial temperature of water in the cup ___________ °C 3) Final temperature of solution in the cup after ALL solid is dissolved ___________ °C 4) Final mass of cup, lid, and solution ___________ g Calculations

5) Mass of solution in the cup ___________ g 6) ΔT for the solution ___________ °C 7) ΔHsol (Eq.9 and Eq. 10) ___________ J Show your work 8) Mass of unknown solid in solution ___________ g 9) ΔHsol per gram of solid sample ___________ J/g 10) The solution reaction is endothermic exothermic. (Circle the correct answer and explain). Solid unknown #: ____________________

Step II: Heat of neutralization reaction

1) Initial temperature of HCl solution ___________ °C 2) Initial temperature of NaOH solution ___________ °C 3) Average initial temperature of the two solutions ___________ °C 4) Final temperature of neutralized solution ___________ °C 5) Final mass of cup, lid, and solution ___________ g Calculations 6) Mass of solution in the cup ___________ g 7) ΔT for the solution = Tf – Average Ti ___________ °C 8) ΔHrxn (Eq.9 and Eq. 11) ___________ J 9) ΔHrxn per mole of H+ and OH- ions reacting ___________ J/mol 10) The solution reaction is endothermic exothermic. (Circle the correct answer and explain).

Chem1A – Experiment F: Specific Heat 12

Post-Lab Questions: Specific Heat

Please answer the following questions and show all work and units. Express all answers to the correct number of significant digits. All completed questions must be submitted with your lab report. Q1. Explain what the value of Heat Capacitycup in Part A(I) describes, and why it should be a small positive number.

Q2. The Law of Dulong and Petit states that for a metal, Atomic  Mass = !"  !/(!"#.!)!!"

. Calculate the atomic mass of the

metal from your experimental value of csp of the metal. Look up the book value for the atomic mass of the metal from the Periodic Table. Then calculate

% Deviation = 100)(

)((exp) xAM

AMAM

book

book− = __________ %

Q3. Recalculate the atomic mass of the metal, using the reference value of csp in the Dulong and Petit relation. Comment on the accuracy of the Law of Dulong and Petit.

Q4. When 2.0 g of NaOH were dissolved in 49.0 g water in a calorimeter at 24.0 °C, the temperature of the solution went up to 34.5 °C. Is this solution reaction exothermic? Explain. Calculate ΔHrxn of 1.00 mol NaOH in water.