Calorimetry

In theory, the specific heat of a given metal should be less than or equal to 0.5J/goC, and the energy of neutralization of hydrochloric and nitric acid with sodium hydroxide should be between -57 and -58kJ/mol. In order to verify the theory, a given mass of unknown metal was heated to thermal equilibrium and then combined with room temperature water in a calorimeter. The metal was supposed to change temperature, resulting in the transition of a certain amount of energy from the metal to the water. By using the masses of water and the metal, the specific heat of the metal can be calculated using the thermal equilibrium formula. The second part involves the determination of the enthalpy of neutralization of acid and base reaction. For one to test the hypothesis, each acid is combined with sodium hydroxide in a calorimeter, and two separate trials made. The changes in temperature for the hydrochloric acid and the nitric acid can then be used to calculate the enthalpy change. The same method applies to the determination of the heat of neutralization while dissolving salt.

Introduction

Calorimetry is a chemical process of determining the quantity of heat involved in a reaction or other methods. In the experiment, the changes of the amount depend on whether there is a loss or gain in heat energy, and the changes are measured using a laboratory thermometer. In definition, calorimetry is the spontaneous transfer of heat from a system of higher heat-energy to that with a lower strength; with the two systems in contact (Letitia and Kleppa). The transmission of heat stops after the two systems have attained an equilibrium temperature. During the changes, the quantity of heat energy is measured using the thermometer, with the calorimeter as the instrument of measuring transfer during the physical or chemical changes (Navrotsky). A calorimeter is a simple device which consists of two styrofoam coffee cups and loose lid with a hole at the top for inserting the thermometer.

Heat changes occur inside the calorimeter, but the thermometer through the lids measures the temperature changes. The heat losses to the surrounding are assumed to be negligible for a well-lagged calorimeter because the rate of heat transfer is faster than when the insulation is poor. The calorimeter is treated as an isolated system where the sum of heat energy for the whole system cancels out. For the measuring of the chemical or heat changes, the temperatures before and after the system reach the equilibrium is determined, with the small amount of energy being lost to the surrounding. The trend line will then be extrapolated to account for the heat small heat losses to the surrounding. For the case of the metal (part A); the amount of heat absorbed by water equals the heat lost by the metal. For the second case of determining the enthalpy of neutralization of acid and base reaction, HCL and nitric acid are mixed in different trials with sodium hydroxide in a calorimeter. A reaction takes place, and the amount of heat evolved by the reaction equals the quantity of heat absorbed by the mixture of HCl, nitric acid and the sodium hydroxide solution. The last part (section C) deals with determining the heat of neutralization while dissolving salt, which is based on the fact that an acid-base neutralization reactions produce heat and water. The neutralization process occurs according to the equation:

.

The quantity of heat produced in the process can be measured in the calorimeter

Procedure

Specific heat capacity of the metal

Water is heated to boiling point in a 400ml beaker, and then the unknown alloy of 20 g mass metal is prepared by transferring them to a dry test tube.

The test-tube is paced in the boiling water, with the metal inside.

The water, test tube with the metal inside are boiled together until and steady equilibrium temperature is reached for the beaker and its contents.

While the water is still boiling at the steady temperature, water at a lower temperature is prepared in the calorimeter at a room temperature.

The mass of water is recorded, followed by the mass of the calorimeter plus its contents.The initial temperature of the calorimeter and its contents is recorded.

The final step is to transfer the hot metal from the beaker to the calorimeter, stir the mixture and record the new thermometer reading after a given time interval.

The data is collected in a table, and the temperatures plotted as a function of time.

Enthalpy of Neutralization of acid-base reaction

For this section, the volume and temperature of a one molar hydrochloric acid are measured and recorded.

The same volume of 1 Molar sodium hydroxide solution is measured and recorded using a different clean beaker.

The temperature of the base is noted and recorded. The base is then mixed with the acid carefully at given time intervals, and the temperature captured after every second for the first one minute, and then after 30-45 seconds for the next 5 minutes.

The data of the temperature and time are recorded in the table and a graph of temperature as a function of time plotted. From the graphical representation, the maximum temperature can be determined through extrapolation.

The experiment is repeated to determine the total heat of neutralization.

Enthalpy of solution for the Dissolution of a Salt

For us to determine the heat of solution of salt, 5.0 g of a known salt is measured and the value recorded on the experiment sheet.

The mass of a dry calorimeter is measured, and 20ml of a deionized salt solution is added into the calorimeter.

The mass of the calorimeter and its contents is noted together with the final temperature of the mixture.

The temperature is read and recorded at a five-second interval for 1 minute, and after that every 30-45 seconds for 5 minutes.

A graph of temperature versus time is plotted, and minimum or maximum temperatures determined depending on whether the process is exothermic or endothermic.

Data and Results

Specific heat capacity of the metal

Measurement

Trial 1

Trial 2

Mass of metal (g)

10.582

19.113

Temperature of metal (boiling water) (0C)

98.9

98.6

Mass of calorimeter, mc (g)

39.85

38.25

Mass of calorimeter + water, M (g)

57.50

57.51

Mass of water, mw= M- mc

17.66

19.26

Temperature of water in the calorimeter, Tw (0C)

22.7

22.4

Maximum temperature of metal and water from graph

26.7

28.5

The following table shows the variation of temperature with time after adding the unknown metal at 98.9 0C to the calorimeter (Chieh).

Time (s)

Water temperature (0C)

Trendline dataset (0C)

0

22.7

36

35.0

50

35.8

70

35.6

35.8

94

35.5

35.6

115

35.5

35.5

135

35.4

35.4

155

35.3

35.4

170

35.0

35.3

205

34.9

35.0

230

34.9

34.9

240

34.9

34.9

270

34.8

34.8

300

34.6

34.6

For the specific heat capacity;

Measurement

Trial 1

Trial 2

Temperature change of water,

4

6.1

Heat gained by water (J)

295.108

491.167

Temperature change of the metal

72.2

70.1

Specific heat of the metal (

0.386

0.366

Average specific heat

0.376

The temperature changes are used to calculate the average specific heat capacity as follows (assuming no heat loses to the surrounding):

Specific heat capacity of the metal, (metal) = =

Enthalpy of Neutralization of acid-base reaction

Measurement

Trial 1

Trial 2

Volume of acid (ml)

50.0

50.0

Temperature of acid (0C)

23.2

23.7

Volume of NaOH (ml)

50.0

50.0

Temperature of NaOH (0C)

23.6

24.0

Exact molar concentration of NaOH (mol/L)

0.969

Maximum temperature from graph (0C)

29.5

30.0

Calculations

Measurement

Trial 1

Trial 2

Average initial temperature of acid and NaOH

23.4

23.3

Temperature change

6.1

6.7

Volume of the mixture

100

100

Mass of the final mixture (g)

100

100

Specific heat of the mixture (J/g 0C)

4.18

Heat evolved

2549. 8

2800.6

Moles of OH- reacted, the limiting reactant (mol)

0.04845

0.04845

Moles of water formed

0.04845

0.04845

(kJ/ mol of H2O)

52.303

57.803

Average (kJ/ mol of H2O)

55.053

The amount of heat evolved by a reaction is equal to the amount of heat absorbed by the mixture.

For this reaction, a graph of the temperature of the solution versus time is as shown below:

Enthalpy of the solution for the Dissolution of a Salt – Ammonium chloride (NH4Cl)

Measurement

Trial 1

Trial 2

Mass of salt (g)

5.023

5.397

Mass of calorimeter (g)

40.225

40.225

Mass of calorimeter + water (g)

59.834

74.997

Mass of water (g)

19.614

34.752

Initial temperature of water (oC)

22.2

22.3

Final temperature of the mixture from graph (oC)

11.4

10.1

Calculations

Measurement

Trial 1

Trial 2

Temperature change of solution (oC)

10.8

12.4

Heat change of water

885.43

1801.27

Heat change of salt (J)

85.17

105.06

Total enthalpy change (J)

970.62

1906.33

(J/ mol of salt)

10436.77

18874.55

Average (J/ mol of salt)

14655.66

In carrying out the calculations above, the calorimeter is assumed to be an isolated system with the heat lost to the surrounding being negligible. The graph of temperature as a function time is as shown below:

Discussion

The experiment was divided into three sections. The first part involved the determination of the specific heat of an unknown metal. When the heated metal was put into the calorimeter containing water at a room temperature, the two reached a thermal equilibrium after a while. The average thermal equilibrium temperature was recorded as 26.7 oC while the maximum value from the graph was 28.5 oC. The graph value is essential since it corrects the heat loss to the calorimeter before the attainment of the equilibrium value (Navrotsky). The heat loss is equal to the heat gained because the calorimeter is assumed to be isolated. The water acquired an average of 393.14J of heat while the metal lost the same quantity of heat. From these observations, the average specific heat capacity of the metal was found to be 0.376 . The value is below the 0.5 as expected.

When a mixture of the acids was formed, the H+ ions and the OH- formed water, and the energy was released. The heat lost by the reactants was absorbed by the mixture in the calorimeter, leading to an increase in temperature. Various peak temperatures were recorded for the acids, and the values were almost equal to the ones achieved through extrapolation. The graphical figures helped in correcting the heat lost to the calorimeter before the peak value as reached. Again, the calorimeter is assumed to be an isolated system, meaning that the heat lost by the reactants equals the heat gained by the mixture. For this experiment, the heat-energy of neutralization is found to be 55.05 kJ/mol which is slightly below the expected value (Beran). For the dissolution of ammonium chloride salt, the enthalpy is 14655.66 kJ/mole of salt. The error could be as a result of the inaccurate change in temperature estimate, and underestimations of the moles of water produced by the reaction. Apparently, the rate of cooling for the acids is also slower than as expected. The reason could be an overestimation of the amount of water produced during the process, or water spillage during the experiment.

Conclusion

Even though the experiment gives some empirically reliable results, the best outcomes could be achieved by limiting the error sources. The materials for use need to be in good condition, and the experiment needs to be done concisely to avoid spillages and under/over-estimations. The unknown metal has an average specific heat capacity of 0.367 J/goC, the heat of neutralization of the base-acid reaction is 55.06kJ/mole of water produced. The energy of dissolution of the ammonium chloride salt is 14655.66 kJ/mole of salt.

Work Cited

Beran, J. A. Laboratory Manual for Principles of General Chemistry. New Jersey: John Wiley & Sons, 2011.

Chieh, Chung. “Measuring Heats of Reactions.” 20 October 2015. Calorimetry. 25 November 2017 .

Letitia, Topor and O. J. Kleppa. “Enthalpies of formation of first-row transition-metal diborides by a new calorimetric method.” The Journal of Chemical Thermodynamics 17.11 (1985): 1003-1016.

Navrotsky, Alexandra. “Progress and new directions in high-temperature calorimetry revisited.” Physics and Chemistry of Minerals 24.3 (1997): 222-241.