Lab 7 - Measuring Enthalpy Changes and Gas Laws

Purpose

To observe changes in enthalpy in chemical processes. To determine the relationship between the pressure and volume of a gas.

Goals

• 1
  To identify exothermic and endothermic processes.
• 2
  To relate enthalpy changes and entropy changes to changes in free energy.
• 3
  To learn to use the MicroLab syringe assembly and interface.
• 4
  To graphically display the relationship between pressure and volume of a gas.

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, for instance, 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 

=

T ΔS. 

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 = T ΔS. 

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 = T_final – T_initial = 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 and B 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 = T_final – T_initial = 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 Part B, you will observe a temperature change that occurs during a chemical reaction. In Part C of the lab, you will observe temperature changes that occur as energy is added to water in its solid and liquid phases. Part D of this lab will involve investigating the relationship between 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. 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

Reagents

• ~2 g CaCl₂ (s, anhydrous)
• ~2 g NH₄NO₃ (s)
• 1 mL 1 M NaOH (aq)
• 1 mL 1 M FeCl₃ (aq)
• 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 with water at the eyewash station for at least 15 minutes. 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 for at least 15 minutes. Have your lab partner notify your instructor about the accident. The chemicals used in this experiment are dilute and in low quantity so gloves will not be available. The NaOH solution has an alkalinity similar to that of oven cleaner or bleach solution. Remember to wash your hands with soap and water when the experiment is completed. 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 and C can be disposed of down the sink drain followed by washing with plenty of water.

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.

Lab Procedure

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

Part A: Heat of Solution - CaCl₂ and NH₄NO₃

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

Part C: Enthalpy and Phase Changes

Part D: Gas Laws

Reference for Part D

• •
  MicroLAB Data Collection Technology. https://microlabinfo.com/wp-content/uploads/2021/09/6.1.BOYLE_S-LAW.instructor-res-1.pdf (accessed June 16, 2022).

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