2 HEAT EFFECTS AND CALORIMETRY

MA, PAULINE Date Performed: 10 SEPTEMBER 2014
NACIONGAYO, DANIELLE Date Submitted: 17 SEPTEMBER 2014
TEDERA, YVES

HEAT EFFECTS AND CALORIMETRY
Experiment No. 2

I. RESULTS
A. Determination of Heat Capacity
In this experiment, an improvised calorimeter was used to determine the heat capacity. The calorimeter weighed 4.47 grams prior to the addition of water. Tap water, 40 mL to be exact, was added to the calorimeter which increased the weight to 43.87 grams. The water was measured using a thermometer thrice (at room temperature, boiling, and addition of boiling water to room temperature water). The first temperature that was recorded was the constant temperature and it had a value of 28.5°C. The next measurement was after 40 mL of tap water was heated in a beaker, and. After a 3 minute observation, the temperature reading was 63.0°C. The heated tap water was immediately poured into the calorimeter and its temperature was also observed. The water initially obtained a temperature of 44.5°C and after 3 minutes, the temperature of the water reduced to 43.5°C. The calorimeter was weighed and its mass increased to 78.85 grams. A graph of the summary of observations is shown in Figure 1.

During the experiment, it was observed that heat was both gained and lost. The hot water lost 2,853.95 Joules (J) while the tap water gained 2,472.74 Joules (J). In general, the calorimeter has a heat capacity of 25.414 J/°C. All the data that was recorded is shown in Table 1.

Table 1. Determination of Heat Capacity of calorimeter
Mass of empty calorimeter
4.47 grams
Mass of calorimeter with tap water
43.87 grams
Mass of calorimeter + hot water
78.85 grams
Initial temperature of tap water in the calorimeter
28.5°C
Initial temperature of hot water (in the beaker)
63.0°C
Final temperature of water in the calorimeter + hot water
43.5°C
ΔT of the tap water and calorimeter
15.0°C
ΔT of the hot water
-19.5°C
Heat lost by the hot water
-2853.95°C
Heat gained by the tap water
2472.74°C
Heat gained by the calorimeter
381.21 Joules
Heat capacity of the calorimeter
25.414 J/°C

B. Heat of Neutralization The same calorimeter used in Part A was used in this part of the experiment. NaOH and HCl were measured separately and placed in the calorimeter and a beaker. The temperatures were measured and NaOH was observed to have a temperature of 28.5°C, while HCl had a temperature of 30.5°C. HCl was then poured into the calorimeter containing NaOH and the temperature was recorded. After 3 minutes, the temperature of the solution was measured and was observed to be 34.5°C. A graph of the summary of observations is shown in Figure 2. A few drops of phenolphthalein indicator were used to test the mixture. It was observed that the substance in excess was a base because the mixture changed color from colorless to light pink.

The solution produced an exothermic reaction wherein both the calorimeter and the solution gained heat. The solution gained 2,560.61 Joules and the calorimeter gained 152.48 Joules. The total heat of the reaction was -2713.09 Joules. The heat of neutralization per mole water formed was -67827.35 J/mol. Table 2 shows the summary of all the results.

Table 2. Heat of Neutralization
Temperature of 1.0M HCl
28.5°C
Temperature of 1.0M NaOH
30.5°C
Mole limiting reagent
0.04
Temperature of the equilibrated mixture
34.5°C
Heat gained by the mixed solution
2560.61 J
Heat gained by the calorimeter
152.484 J
Total heat of the reaction
-2713.094 J
Heat of neutralization per mole of water formed
-67827.35 J/mol

C. Exothermic Reaction The same improvised calorimeter was still used in this experiment and was filled with 50 mL distilled water. NaOH flakes were obtained and mixed into the water. Prior to the addition of NaOH flakes, the temperature of the water was observed to 28.0°C. After the dissolving the flakes, the temperature was also observed and recorded. The temperature of the solution was observed to be 39.0°C after solvation. After 3 minutes of observation, the solution increased to a constant temperature of 46.0°C. The heat absorbed by the calorimeter with water was 3720.41 Joules. The addition of NaOH produced an exothermic reaction, therefore the heat of the solution was
-4177.86 Joules. The heat of the solution per gram solid was -835.57 J/g while the heat of solution per mole was -33422.95 J/mol. Table 3 shows a summary of all the results.

Table 3. Heat of Solution (Exothermic)
Temperature of water
28.0°C
Temperature of equilibrated solution
46.0°C
Mass of calorimeter + water
53.87 grams
Mass of NaOH
5.0 grams
Heat absorbed by the calorimeter + water
3720.41 J
Heat of solution
-4177.86 J
Heat of solution per gram solid
-835.57 J/g
Heat of solution per mole
-33422.95 J/mol

D. Endothermic Reaction The same improvised calorimeter was still used in this experiment and was filled with 50 mL distilled water. Na2S2O3 flakes were obtained and mixed into the water. Prior to the addition of Na2S2O3 flakes, the temperature of the water was observed to 28.0°C. After the dissolving the flakes, the temperature was also observed and recorded. After 3 minutes of observation, the solution decreased to a constant temperature of 23.0°C. The heat absorbed by the calorimeter with water was -653.33 Joules. The addition of Na2S2O3 flakes produced an endothermic reaction; therefore the heat of the solution was 780.40 Joules. The heat of the solution per gram solid was 156.08 J/g while the heat of solution per mole was 24696.25 J/mol. Table 4 shows a summary of all the results.

Table 4. Heat of Solution (Endothermic)
Temperature of water
28.0°C
Temperature of equilibrated solution
23.0°C
Mass of calorimeter + water
35.7 grams
Mass of sodium thiosulfate (Na2S2O3)
5.0 grams
Heat absorbed by the calorimeter + water
-653.33 J
Heat of solution
780.40 J
Heat of solution per gram solid
156.08 J/g
Heat of solution per mole
24696.25 J/mol

II. DISCUSSION
A. Determination of Heat Capacity
The heat capacity is the amount of heat required to raise the temperature of an object or substance one degree. The temperature change is the difference between the final temperature ( Tf ) and the initial temperature ( Ti ).
A calorimeter is an experimental device in which a chemical reaction or physical process takes place. The calorimeter is well-insulated so that, ideally, no heat enters or leaves the calorimeter from the surroundings. For this reason, any heat liberated by the reaction or process being studied must be picked up by the calorimeter and other substances in the calorimeter. In this experiment, a Styrofoam cup is used as a calorimeter, because it has good insulated walls that can prevent heat exchange with the environment.
A thermometer is typically inserted in the calorimeter to measure the change in temperature that results from the reaction or physical process. A stirrer is employed to keep the contents of the calorimeter well-mixed and to ensure uniform heating.
In a calorimetric experiment, it is important to measure the heat capacity. This is because the heat capacity of the calorimeter and the temperature change it undergoes can be a significant part of the calorimetric calculation. If not considered, if not computed, this heat effect could drastically influence the results of a heat determination.

B. Heat of Neutralization
The standard enthalpy change of neutralization is the enthalpy change when solutions of an acid and an alkali react together under standard conditions to produce 1 mole of water.
Enthalpy changes of neutralization are always negative because heat is released when an acid and an alkali react. For reactions involving strong acids and alkalis, the values are always very closely similar, with values between -57 and -58 kJ mol-1. That varies slightly depending on the acid-alkali combination. In this experiment, the calculated value of the heat of neutralization of NaOH and HCl was -67827.35 J/mol. When converted to Kilojoules (kJ), the value is -67.83 kJ/mol, which is close to the theoretical value.

C. & D. Heat of Solution
Calorimetry is the science of measuring heat flow, and heat is defined as thermal energy flowing from an object at a higher temperature to one at a lower temperature. Ideally, the heat changes resulting from physical and chemical phenomena can be harnessed to do work. Most physical and chemical changes are either exothermic or endothermic. Exothermic reactions release energy or heat to increase the temperature of the surroundings; thus, the surroundings are hotter after an exothermic change. For example, nitroglycerine exploding is an extremely exothermic reaction. Endothermic reactions absorb energy or heat to decrease the temperature of the surroundings; thus, the surroundings are colder after an endothermic change. When a reaction is carried out under constant pressure, the heat of a reaction is defined as the enthalpy change for the reaction (H). Since most reactions occur under constant atmospheric pressure, the heat of a reaction is equal to H, which is generally reported in units of kilojoules (kJ) per mole of the reactants and products as shown in the balanced thermochemical equation.
The enthalpy of solution depends on the intermolecular forces of the solute and solvent. If the solution is ideal, and ΔHsolution = 0, then that means ΔH1 added to ΔH2 is equal to ΔH3.
In the experiment, the dissolution of NaOH flakes in water produces a reaction that releases heat to the environment. Thus, it produces an exothermic reaction. Both ΔH and q are negative in this reaction. However, if sodium thiosulfate is dissolved in water it absorbs heat from the environment. Since it is an endothermic reaction, ΔH and q are both positive.

III. ANSWERS
1. Plotting the temperature as a function of time ensures that the temperatures obtained are constant, which proves that the system is in thermal equilibrium.

2. In a calorimetric experiment, it is important to measure the heat capacity. This is because the heat capacity of the calorimeter and the temperature change it undergoes can be a significant part of the calorimetric calculation.
3. For an isolated system, one component of a system has to absorb the heat released by another component. The signs are just conventions used to designate release (released) and gain (positive).
4. When using Styrofoam as a cover for the calorimeter, it’s possible for the material to lose some of its parts, resulting to a change in weight and may also cause some leaks in the set-up where heat can enter or escape. Inaccuracies in thermometer readings cannot be ruled out and can also be a source of error.

IV. CONCLUSION
Calorimetry is the science of measuring heat flow, and heat is defined as thermal energy flowing from an object at a higher temperature to one at a lower temperature. Most physical and chemical changes are either exothermic or endothermic. Exothermic reactions release energy or heat to increase the temperature of the surroundings. Endothermic reactions absorb energy or heat to decrease the temperature of the surroundings. When a reaction is carried out under constant pressure, the heat of a reaction is defined as the enthalpy change for the reaction (ΔH). Since most reactions occur under constant atmospheric pressure, the heat of a reaction is equal to ΔH, which is generally reported in units of kilojoules (kJ) per mole of the reactants and products as shown in the balanced thermochemical equation.
In this experiment, the calorimeter used will consist of two nested Styrofoam coffee cups with a cover and a thermometer. Ideally, the calorimeter is so well insulated, that all of the heat gained or lost during the reaction is completely contained within the calorimeter. In reality, the coffee-cup calorimeter is not a perfect insulator, so it will actually absorb and lose heat.
Heat transfer does not occur instantaneously, so using a calorimeter to determine the enthalpy change requires measuring the temperature of the calorimeter contents as the physical or chemical change occurs and for several minutes afterwards. The maximum temperature reached by the calorimeter contents can rarely be measured directly because of drastic temperature fluctuations near the probe or heat lost at the time of mixing. If the contents are mixed uniformly while data is recorded, the contents should equilibrate at a consistent rate. Thus, the time-temperature data can be plotted, and the resulting regression line can be extrapolated to the time of mixing to get the maximum or final temperature.

V. CALCULATIONS
A. Determination of Heat Capacity qcal = ? qcal = - ( qhot + qtap ) qcal = - {[(78.85-43.87)g(4.184 J/g°C)(-19.5°C)] + [(43.87-4.47)g(4.184 J/g°C) (15°C)]} qcal = - (-381.21 J) qcal = 381.21 J

Ccal = ?
Ccal = qtap / ΔTtap
Ccal = 381.21 J / 15°C
Ccal = 25.414 J/°C
B. Heat of Neutralization
Moles of limiting reagent = ?
HCl limiting reagent

HCl + NaOH NaCl + H2O i 0.04 0.06 c -0.04 -0.04 +0.04 e 0 0.02 0.04

Heat gained by mixture = ? qmixture = mCsΔT qmixture = [(100 ml)(1.02 g/ml)][(4.184 J/g°C)(34.5 – 28.5)°C] qmixture = 2560.61 J

Heat gained by calorimeter = ? qcal = Ccal ΔT qcal = (25.414 J/g°C)(6°C) qcal = 152.484 J

Total Heat = ?
Total Heat = - (qcal + qsoln)
Total Heat = - (152.484 + 2560.61) J
Total Heat = - 2713.094 J

Heat of Neutralization = ?
Heat of Neutralization = Total heat / moles of limiting reagent
Heat of Neutralization = -2713.094 J / 0.04 mol
Heat of Neutralization = - 67827.35 J/mol

C. Heat of Solution (Exothermic)

Heat absorbed by the calorimeter + water = ? qcal+water = mCsΔT qcal+water = (53.87-4.47)g(4.184 J/g°C)(46-28)°C qcal+water = 3720.4128 J

Heat of solution = ? qsoln = - (qwater + qcal) qsoln = - [ 3720.4128 J + (25.414 J/°C)(18°C) ] qsoln = - 4177.8648 J
Heat of solution per gram solid = ?
Heat of solution per gram solid = Heat of solution/ g NaOH
Heat of solution per gram solid = -4177.8648 J / 5.0 g
Heat of solution per gram solid = -835.57296 J/g

Heat of solution per mole = ?
Heat of solution per mole = Heat of solution/ mol NaOH
Heat of solution per mole = -4177.8648 J / (5.0 g)(1 mol/ 39.99714 g)
Heat of solution per mole = -4177.8648 J / 0.125 mol
Heat of solution per mole = -33422.9472 J/mol

D. Heat of Solution (Endothermic)

Heat absorbed by the calorimeter + water = ? qcal+water = mCsΔT qcal+water = (35.7-4.47)g(4.184 J/g°C)(23-28)°C qcal+water = -653.3316 J

Heat of solution = ? qsoln = - (qwater + qcal) qsoln = - [ (-653.3316) J + (25.414 J/°C)(-5°C) ] qsoln = - 780.4016 J

Heat of solution per gram solid = ?
Heat of solution per gram solid = Heat of solution/ g NaOH
Heat of solution per gram solid = 780.4016 / 5.0 g
Heat of solution per gram solid = 156.0803 J/g

Heat of solution per mole = ?
Heat of solution per mole = Heat of solution/ mol NaOH
Heat of solution per mole = 780.4016 J / (5.0 g)(1 mol/ 158.1478 g)
Heat of solution per mole = 780.4016 J / 0.0316 mol
Heat of solution per mole = 24696.25 J/mol

VI. REFERENCES

Petrucci, Harwood, Herring, Madura. General Chemistry: Principles & Modern Applications. (9th ed). Pearson Education, Inc.
McMurray, Fay. Chemistry, (3rd ed.). Prentice-Hall, Inc.

References: Petrucci, Harwood, Herring, Madura. General Chemistry: Principles & Modern Applications. (9th ed). Pearson Education, Inc. McMurray, Fay. Chemistry, (3rd ed.). Prentice-Hall, Inc.