Evaluating an enthalpy change that can not be measured directly

Essay by Spaceman1College, UndergraduateA, January 1996

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Chemistry Experiment.

Dr. Watson.

Evaluating An Enthalpy Change That Cannot

Be Measured Directly.


We were told that sodium hydrogencarbonate decomposes on heating to give sodium

carbonate, water and carbon dioxide as shown in the equation below:-

2NaHCO3(s)--------> Na2CO3 (s) + H2O (l) + CO2 (g) = DeltaH1

This was given as deltaH1 and we had to calculate as part of the experiment.

This however cannot be measured directly, but can be found using the enthalpy

changes from two other reactions. These being that of sodium hydrogencarbonate and

hydrochloric acid and also sodium carbonate and hydrochloric acid.

We were given a list of instructions in how to carry out the experiment, which are

given later.

List of Apparatus Used.

1 x 500ml Beaker.

1 x Thermometer(-10 to 50oC).

1 x Polystyrene Cup.

1 x Weighing Balance.

1 x Weighing Bottle.

10 grams of Sodium Hydrogencarbonate.

10 grams of Sodium Carbonate.

A bottle of 2 molar HCL.



Three grams of sodium hydrogen carbonate was weighted out accurately using a

weighting bottle and a balance. Then thirty centimetres cubed of 2 molar HCL was

measured using a measuring cylinder. The acid was then placed into the polystyrene

cup and its temperature was taken and recorded using the thermometer. The pre-

weighted sodium hydrogencarbonate was then added to the solution, and the final

temperature was recorded.

The contents of the cup were then emptied out and the cup was washed out with

water and then thoroughly dried. This was done three times for the sodium hydrogen

carbonate so that I could remove any anomalies that were obtained.

The experiment was then repeated in exactly the same manner except sodium

carbonate was used instead of sodium hydrogen carbonate.

The results were then tabulated, this table is shown below.

Results Table.

Results Table for Sodium Hydrogencarbonate.

Results Table for Sodium Carbonate.


From these results I had to calculate deltaH2 and deltaH3. DeltaH2 refers to the

enthalpy change when sodium hydrogencarbonate reacts with hydrochloric acid, and

deltaH3 is the enthalpy change when the sodium carbonate reacts with the acid.

Firstly however it is necessary to show the equations for the two reactions:-

DeltaH2= 2NaHCO3 (s) + 2HCl (aq)----> 2NaCl (aq) + 2H2O (l) + 2CO2 (g).

DeltaH3= Na2CO3 + 2HCl (aq) ----> 2NaCl (aq) + H2O (l) + CO2 (g)

The enthalpy changes of the two reactions can be worked out using the formula

shown below :-

Energy Exchanged between = Specific Heat Capacity x Mass of the x Temperature

Reactants and Surroundings of the Solution Solution Change.

Therefore the DeltaH2 of the reaction when fitted into the formula is :-

Energy Exchanged between = 4.18 x (84 x 2) x -11.1

Reactants and Surroundings.

This gives the enthalpy change for DeltaH2 to be = -7794.9 Joules per mole.

The same formula is used for DeltaH3:-

Energy Exchanged Between = 4.18 x 106 x 21.8

Reactants and Surroundings.

This gives the Enthalpy change for DeltaH3 to be = 9659.1 Joules per mole.

From these two results we are able to work out what DeltaH1 is likely to be even

though we have not done the experiment. This is done using the formula :-

DeltaH1 = DeltaH3 + DeltaH2 =>

DeltaH1 = 9659.1 + (-7794.9) =>

DeltaH1 = 1864.2 Joules per mole.


The result obtained will not be a very accurate due to the means by which the

experiment was done. The equipment used was not the most efficient for measuring

enthalpy changes, however it does give a rough estimate to work from. Some errors

of the equipment would have been heat lost through conduction from the reaction

vessel. Also heat may well have been lost through the open top of the container, even

though there was a lid this was not very secure some heat will have escaped through


In summation the experiment was very difficult to undertake as the enthalpy change

for DeltaH1 is hard to determine due to the fact that it thermally decomposes in the

air, causing great problems in calculating its enthalpy change with its surroundings.