Last Updated 28 Jan 2021

Experiment 1 Calorimetry

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All chemical reactions involve energy. By understanding the behavior and connection of energy flow within a chemical reaction, we can understand and manipulate them to our advantage. The most common form of energy observed during chemical reactions is heat. The reaction may absorb (endothermic) or release (exothermic) heat, depending on the reacting substances. Calorimetry is the process of measuring the heat flow between a system and its environment. The device used to measure this heat transfer is called a Calorimeter.

The measurement of this heat is called the enthalpy of the reaction (? H). There are two types of calorimeter. The first is a bomb calorimeter where the reaction takes place at constant volume. The other type is the coffee cup calorimeter, wherein the pressure is held constant while the reaction takes place. In the experiment, a modified coffee cup calorimeter is used. It is made up of a Styrofoam ball for insulation, a six-inch test tube, a cork and a thermometer. The setup used is adiabatic, which means, the system is isolated from the surroundings so there is no heat flow.

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The first step is to calibrate the calorimeter. Every calorimeter has its own specific heat constant (CCal). NaOH is first poured into the test tube, and then the Ti is measured. Then, HCl is made to react with the base and then the Tf is measured. The net ionic equation of the neutralization reaction is show below. H+(aq) + OH-(aq) H2O(l) ? H = -55. 85KJ Using the data obtained (? T) and the knowledge of the enthalpy of reaction of neutralization of an acid by a base, the CCal can be computed.

Equation (1) is an application of the Law of Conservation of Energy, where the heat released by the reaction will be absorbed by the surroundings. Equation (2), which is the formula for computing CCal can be derived from the thermochemical equations. -qreaction= qcalorimeter (1) Ccal= nLR ? ?H? T (2) Table 1 shows the average Ccal values of each group. The values can then be used to compute for the ? H of other reactions.

Each calorimeter has its own Ccal value, which will be used in the following computations.

1 1. 520
2 1. 117
3 1. 047
4 1. 100
5 1. 894
6 0. 698
7 0. 960
8 2. 304
9 1. 224

After the calorimeter is calibrated and the CCal is computed, the ? H of the other reactions can now be obtained. The same procedure is repeated using the other specified pairs of reagents. After the ? T is obtained, the ? H of the reaction can be calculated using the previously obtained CCal value through Equation (3) derived formulas from the thermochemical equations.

Table 2 shows the computed average ? H values of the different reactions in the experiment, theoretical values of ? H calculated from the data in the lab manual using Equation (4), and the computed percent error. All reactions were assumed exothermic based on the results. The amount of heat released by the reactions varied. The reactions between active metals and acids released the most amount of heat while the neutralization reactions and the synthesis reactions released fair amounts of heat. ?Hrxn= -Ccal? ?TnLR (3) H= vp? Hproducts-vp? H(reactants) (4)

Table 2. Tabulated results.

NH3 + HCl -52. 22 KJ/mol -121. 6 KJ/mol Exothermic 132. 9%
NaOH + CH3COOH -52. 47 KJ/mol -39. 09 KJ/mol Exothermic 25. 5%
NH3 + CH3COOH -56. 09 KJ/mol -68. 055 KJ/mol Exothermic 21. 3%
NaOH + HNO3 -55. 85 KJ/mol -4. 4 KJ/mol Exothermic 92. 1%
HCl + Mg -466. 85 KJ/mol -478. 05 KJ/mol Exothermic 2. 4%
CH3COOH + Mg -953. 1 KJ/mol -237. 15 KJ/mol Exothermic 75. 1
CuSO4 + Zn -426. 65 KJ/mol Exothermic 75. 1
Na2CO3 + CaCl3 -13. 07 KJ/mol -276. 48 KJ/mol Exothermic 95. 1%
Na2CO3 + CaCl3 -13. 07 KJ/mol -1. 224 KJ/mol Exothermic 87. 1%

The results show that each reaction releases or absorbs different amounts of heat depending on the nature of the reactants and type of reaction. Displacement reactions by active metals release significantly higher amounts of heat than neutralization reactions while the synthesis reactions yielded the least amount of heat.

The results of the experiments support the fact that energy flows in and/or out of a system during a chemical reaction.


  1. Petrucci, R. Appendix D Data Tables. Gen. Chem. 2011, 10th Ed. , 242 - 244
  2. Brown, W. Doc Brown’s Chemistry. http://www. docbrown. info/page07/delta1He. htm (accessed November 17, 2012)
  3. Petrucci, R. Appendix D Data Tables. Gen. Chem. 2011, 10th Ed. , 271
  4. University of the Philippines Diliman Institute of Chemistry. Experiment 1 Calorimetry. General Chemistry II Laboratory Manual. 2011. 1- 8
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