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Lab 7 - Measuring Enthalpy Changes and Gas Laws

Purpose

Goals

Introduction

Chemical processes occur spontaneously when they lower the free energy of the system. The free energy at constant temperature and pressure is the Gibbs free energy,
ΔG,
 which is defined as follows:
( 1 )
ΔG = ΔH − TΔS,
where
ΔG
 is the change in free energy of the system,
ΔH
 is the change in enthalpy of the system and
ΔS
 is the change in entropy of the system.
ΔG,
 the change in free energy, must be less than zero for a spontaneous process. This follows from the definition of a spontaneous process; it lowers the energy of the system. Free energy, given the symbol G in honor of the American chemist, J. Willard Gibbs, is the energy available from a spontaneous process to do work. For a nonspontaneous process, it is the amount of energy that must be supplied to cause the process to occur.
ΔH,
 the enthalpy change, is the amount of heat energy absorbed by a process that occurs at constant pressure (–
ΔH
 is the heat given off).
ΔH
 is often proportional to
ΔT,
 the temperature change caused by the process. The proportionality constant between
ΔH
 and
ΔT
 is called the specific heat (s). The specific heat is a property of the substance being heated or cooled. For example, the specific heat of water is 4.18 J/(g
·
 °C) meaning that 4.18 joules of heat energy are required to raise the temperature of 1 gram of water by 1.0°C. Not all processes that involve a change in enthalpy are accompanied by a change in temperature. Phase changes, e.g., melting, boiling, and sublimation all absorb heat. However, the substance undergoing the phase change maintains a constant temperature. All the energy supplied to the substance is consumed in the phase change. Processes that give off heat energy (
ΔH
 < 0) are exothermic. Combustion is obviously an exothermic chemical reaction. Condensation of steam to liquid water is also exothermic, although water maintains a constant temperature during the process. Processes that absorb heat energy (
ΔH
 > 0) are endothermic. Sweat, liquid water that evaporates from your skin, cools you by absorbing heat from your body.
ΔS,
 the change in entropy, is related to the number of ways the energy of the system can be distributed. Entropy is commonly defined as disorder or randomness. Entropy is high in gases, because the molecules are free to move about in all directions, fill the space available, and adopt any orientation relative to each other. Entropy is low in solids, because the molecules cannot move much relative to each other, and their orientation is fixed. With this information, one can see that a chemical process will certainly be spontaneous if it is exothermic (
ΔH
 < 0) and its entropy increases (
ΔS
 > 0). If the opposite holds (
ΔH
 > 0 and
ΔS
 < 0), the process cannot be spontaneous. What if both have the same sign? If the process is exothermic, the process will be spontaneous below the temperature that satisfies the condition
ΔH
 =
TΔS.
 It turns out that changes in entropy are small relative to changes in enthalpy in most processes. It is usually not possible to tell whether entropy has increased or decreased in spontaneous exothermic processes. The exceptions occur when gases are produced or consumed. Water will condense spontaneously if the temperature is below 100°C; the process is sufficiently exothermic to offset the loss in entropy that accompanies the phase change. If the process is endothermic, the process will be spontaneous above the temperature that satisfies the condition
ΔH = TΔS.
 As mentioned above, entropy changes are small relative to enthalpy changes when gases are not involved. Therefore, spontaneous endothermic processes are relatively uncommon. However, they must involve an increase in entropy. When working with
ΔG,
ΔH,
ΔT,
ΔS,
 and in fact all chemical terms involving a
Δ
 (change in), the quantity is calculated as final – initial. Thus, if a beaker containing water was originally at 22.5°C and a chemical process caused its temperature to fall to 17.3°C, the change in temperature would be
( 2 )
ΔT = Tfinal – Tinitial = 17.3°C – 22.5°C = –5.2°C.
It is important to remember final – initial in order to get signs right. This can sometimes be confusing. It is important to distinguish between system and surroundings, the two parts of the thermodynamic universe, when dealing with reaction energetics. The first line of this document stated that spontaneous processes lower the free energy of the system. The system is the part of the universe being studied. The first law of thermodynamics states that energy is neither created nor destroyed. The energy released by a system must be converted to some other form or transferred to some other place. Often, the energy is transferred to the surroundings, the part of the universe that interacts with the system being studied. In such studies, it is often practical to isolate a rather small "universe" of system and surroundings in which nearly all the energy is retained. In Parts A–C of this lab, we will assume that the thermodynamic universe is the test tube or beaker in which the experiment is being conducted. Of course, some heat is transferred outside the test tube (you will be able to feel beakers heat up or cool down) but most is retained. The energy released or absorbed by the systems being studied (chemicals dissolving or reacting) will be transferred to or from the water in which the processes occur. Thus, water is a major component of the surroundings. Returning to Equation 2
ΔT = Tfinal – Tinitial = 17.3°C – 22.5°C = –5.2°C.
 and sign changes. In that example, the temperature of the water (the surroundings) decreased and heat was released. This means that a chemical process (the system) absorbed heat from the water. Since the process absorbed energy, it was endothermic. Now consider an exothermic example. If a chemical process (the system) releases heat to the water (the surroundings), the temperature of the water will increase as the water absorbs heat energy and the
ΔT
 of the water will be positive. In Part A of the lab, you will observe temperature changes that occur as compounds dissolve in water. In Parts B and C, you will observe temperature changes that occur during chemical reactions. The reaction in Part C is a neutralization. Neutralization reactions (also known as acid-base reactions, which you will study in more detail later in the course) usually involve transfer of H+ (a proton) from one chemical species to another. For example, the reaction of acetic acid and sodium hydroxide is shown below.
( 3 )
HC2H3O2(aq) + NaOH(aq) → NaC2H3O2(aq) + H2O
Strong acids, such as HCl and HNO3, react with water as shown below.
( 4 )
HCl(aq) + H2O → Cl−(aq) + H3O+(aq)
( 5 )
HNO3(aq) + H2O → NO3−(aq) + H3O+(aq)
H3O+ is called the hydronium ion. Whenever strong acids are placed in water, this is the species that reacts. For example, the reaction with sodium hydroxide is shown below.
( 6 )
H3O+(aq) + NaOH(aq) → Na+(aq) + 2 H2O
In these reactions, species are followed by "(aq)" to indicate that they are present in aqueous solution. This notation will be used frequently in the rest of this course. In Part D of the lab, you will observe temperature changes that occur as energy is added to water in its solid and liquid phases. Part E of this lab will involve investigating the relationship of the volume and pressure of gases. Gases can be characterized by a number of variables, which include the amount, volume, pressure, and temperature. The ideal gas law summarizes this relationship.
( 7 )
PV = nRT
where P = pressure in atmospheres, V = volume in liters, n = moles of gas, R = 0.0821 liter · atmospheres/Kelvin · mole, and
T = temperature in Kelvin. In this lab, you will determine the relationship between the pressure and volume of a gas using the MicroLab interface and syringe assembly. Two different volume and pressure relationships will be studied. The temperature of the system will simply be the temperature of the laboratory. However, the different volumes will allow students to investigate the relationship of volume and pressure. The volume will serve as the independent variable, which is changed by the researcher. The pressure will be the dependent variable, whose response is measured. The temperature should remain mostly unchanged during the experiment and the moles should remain unchanged as the system is sealed.

Equipment

  • 1 30 mL beaker
  • 1 100 mL beaker
  • 1 10 mL graduated cylinder
  • 1 Microlab interface
  • 1 MicroLab thermistor instruction sheet
  • 1 thermistor
  • 4 small test tubes
  • 1 test tube rack
  • 1 ring stand and clamp
  • 1 hot plate/stirrer
  • 1 magnetic stir bar
  • 1 60 mL syringe
  • 1 20 mL syringe
  • connecting adapter
  • 1 deionized water squirt bottle

Reagents

  • ~2 g CaCl2(s, anhydrous)
  • ~2 g NH4NO3(s)
  • 1 mL 1 M NaOH(aq)
  • 1 mL 1 M FeCl3(aq)
  • 8 mL 0.1 M NaOH(aq)
  • 3 mL 0.1 M HCl(aq)
  • 3 mL 0.1 HNO3(aq)
  • 3 mL 0.1 M HC2H3O2(aq)
  • 4 drops phenolphthalein indicator
  • 60 g ice
  • 13 mL deionized water

Safety

The solutions used in this experiment are dilute, but corrosive. They can attack the skin and cause permanent damage to the eyes. If any of the solutions splash into your eyes, flush them in the eyewash. Hold your eyes open or have someone assist you. If you spill any of the solutions on your skin or clothing, flush the area immediately with water. Have your lab partner notify your instructor about the accident. You will be working with a hot plate. Keep flammable materials (papers, hair, clothing) away from it. It will stay hot for a long time after you have turned it off. Do not touch it until you have felt the air near it and are sure it will not burn you. When using the syringes, a moderate amount of force will be needed to hold the syringe plunger at certain volumes against reduced pressure. Care should be taken to grip the syringe firmly to prevent pinched fingers.

Waste Disposal

The reaction mixture from Part B of the experiment should be put in the waste container provided. Use water in a squeeze bottle to rinse the solid material out of the test tube and then wash the tube with soap and water. Solutions from Parts A, C and D can be washed down the sink.

Prior to Class

Please read the following section of Lab Safety and Practices: Preparing Graphs. Please complete your WebAssign prelab assignment. Check your WebAssign Account for due dates. Students who do not complete the WebAssign prelab are required to bring and hand in the prelab worksheet.

PDF file

Lab Procedure

Please print the worksheet for this lab. You will need this sheet to record your data.

PDF file

Part A: Heat of Solution - CaCl2 and NH4NO3

1
Open the MicroLab program. Calibrate the thermistor as described in the MicroLab instructions provided in the lab.
2
After the calibration is complete, set the MicroLab collection increment to 2 seconds using the instructions provided.
3
Use a spoon-type spatula to dispense anhydrous solid CaCl2 into a dry 30 mL beaker. One spoonful, not heaping, is plenty.
4
Obtain 10 mL of deionized water in a graduated cylinder.
5
Measure the initial temperature of the water in the graduated cylinder by placing the thermistor in the water and pressing "start" on the MicroLab data collection program. Record the initial temperature in Data Table A.
6
Pour the water into the beaker containing the CaCl2.
7
Place the thermistor in the beaker. Watch the temperature change on the monitor. Record the maximum or minimum temperature reached in Data Table A. Stop the MicroLab data collection program.
8
When finished, dispose of your solution in the sink, rinse the beaker with deionized water and dry.
9
Repeat steps 3 through 8 with solid NH4NO3 in place of the solid CaCl2.
10
Calculate the
ΔT
 for each reaction and record them in Data Table A.

Part B: Heat of Reaction - FeCl3(aq) + NaOH(aq)

1
Take two test tubes to the side shelf and obtain about 1 mL of 1 M FeCl3 in one and about 1 mL of 1 M NaOH in the other.
2
Start the MicroLab data collection program and use the thermistor to measure the temperature of the FeCl3 solution. Also, note its color. Enter this information in Data Table B.
3
Add the 1 M NaOH to the FeCl3 test tube.
4
Place the thermistor in the test tube containing the mixture. Watch the temperature change on the monitor. Record the maximum or minimum temperature reached in Data Table B. Stop the MicroLab data collection program.
5
Record the appearance of the reaction mixture in Data Table B.
6
Calculate the
ΔT
 for the reaction and record it in Data Table B.
7
When finished, dispose of your solution in the waste bottle provided. Use water in a squeeze bottle to rinse the solid material out of the test tube and then wash the tube with soap and water.

Part C: Heat of Neutralization - NaOH + HCl, NaOH + HNO3, NaOH + HC2H3O2

1
Add 2 mL of 0.1 M NaOH to each of 4 test tubes. Add 1 drop of phenolphthalein to each tube.
2
Start the MicroLab data collection program and measure the temperature of one of the NaOH solutions. Record it in Data Table C. Note the appearance of the solutions in Data Table C. Assume that all the NaOH solutions in this experiment start out at this temperature.
3
Measure 3 mL of deionized water in a graduated cylinder and add it to the first test tube. Record the maximum or minimum temperature reached and any change in appearance in Data Table C.
4
Measure 3 mL of 0.1 M HCl in a graduated cylinder and add it to the second test tube. Record the maximum or minimum temperature reached and any change in appearance in Data Table C.
5
Measure 3 mL of 0.1 M HNO3 in a graduated cylinder and add it to the third test tube. Record the maximum or minimum temperature reached and any change in appearance in Data Table C.
6
Measure 3 mL of 0.1 M HC2H3O2 in a graduated cylinder and add it to the fourth test tube. Record the maximum or minimum temperature reached and any change in appearance in Data Table C. Stop the MicroLab data collection program.
7
Calculate the
ΔT
 for each reaction and record them in Data Table C.
8
When finished, dispose of your solution in the sink, rinse the beaker with deionized water and dry.

Part D: Enthalpy and Phase Changes

1
Change the MicroLab collection increment to 30 seconds.
2
Place a magnetic stir bar in a 100 mL beaker. Fill the beaker to the 60 mL mark with crushed ice and add tap water to the 60 mL mark.
3
Clamp your thermistor in the ring stand with a hot plate/stirrer on the base.
4
Place the beaker on the hot plate/stirrer and lower the thermistor into the ice slurry. The tip of the thermistor should be about a third of the way down in the slurry.
5
Start the stirrer; set the dial to about 6 for gentle stirring.
6
Initiate the MicroLab data collection program.
7
Turn the hot plate to 100 and continue allowing MicroLab to record the temperature.
8
After two minutes, turn the hotplate to 200.
9
Note the temperature when all the ice has melted in Data Table D and turn the heater to 400. Continue allowing MicroLab to record the temperature.
10
As the water heats, record the temperature at which you observe bubbles, steam, and true boiling in Data Table D.
11
Once boiling begins, read the temperature every 30 seconds for three minutes. Then stop the MicroLab data collection program.
12
Make a sketch of the MicroLab graph on your worksheet, noting where the phase transitions occur. Make sure to give your graph a title. Label the axes with the quantities being plotted and the units in which they are measured.
13
When you have completed your measurements, wash and dry all your equipment and return it neatly to the set-up area where you found it. Make sure the hot plate is turned off.

Part E: Gas Laws

1
Open the MicroLab program and select New Experiment.
2
Enter your experiment name, "Volume vs. Pressure," and select OK.
3
Click ADD SENSOR, and use the pulldown menu to select PRESSURE. The program will also request the input. Choose the only input available with a red square for the pressure. Enter "torr" for the label. Then select NEXT and make sure the toggle button is selected for "Factor calibrated torr," followed by FINISH.
4
Click ADD SENSOR and use the pulldown menu to select KEYBOARD. The label will be VOLUME. Enter the prompt as "volume?" and the units as mL, followed by FINISH.
5
To view data as it is collected, click on the keyboard sensor and drag the icon to the x-axis of the graph, and column A of the spreadsheet. Click on the pressure sensor and drag the icon to the y-axis of the graph, column B of the spreadsheet, and the lower right display area.
6
Compress the 60 mL syringe to the 0 mL mark and attach the hose to the MicroLab pressure sensor.
7
Press start on the MicroLab console.
8
Hold the syringe to the 0 mL mark, which is actually 4 mL of gas, as the connector hose is 4 mL. Request that your lab partner enter 4 in the keyboard and take a reading.
9
Take additional readings at the 10, 20, 30, and 40 mL syringe marks. (Please note: this is a two-person operation. One person will need to hold the syringe open against a vacuum, while the other enters volumes using the keyboard and takes pressure readings. Don't forget to add 4 mL to your volumes!) Enter your data in Table E. You will be asked to plot the data in different ways in your WebAssign In-Lab and postlab assignments.
10
Remove the 60 mL syringe and replace with a 20 mL syringe, compressed to the 0 mL mark.
11
Start a new experiment on the MicroLab console.
12
Hold the 20 mL syringe to the 0 mL mark. Request that your lab partner enter 4 in the keyboard and take a reading.
13
Take additional readings at the 5, 10, 15 and 20 mL syringe marks. Enter your data in Table E.
14
Before leaving, go to a computer in the laboratory and enter your results in the In-Lab assignment. If all results are scored as correct, log out. If not all results are correct, try to find the error or consult with your lab instructor. When all results are correct, note them and log out of WebAssign. The In-Lab assignment must be completed by the end of the lab period. If additional time is required, please consult with your lab instructor.
Reference for Part E: MicroLAB Data Collection Technology.

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