Chemical Energetics (5) Measuring the enthalpy change of reaction (ii)

This new blog looks at a common experiment that is often used to show you how to calculate an enthalpy of reaction.

The reaction often used is the redox reaction between zinc powder and copper sulphate solution. 

Zn(s)  +   CuSO₄ (aq)    =    Cu(s)  +   ZnSO₄ (aq)

Zinc powder is added gradually to copper sulphate solution in a polystyrene cup (calorimeter).

Polystyrene minimises heat loss from the reaction mixture especially if a lid is added to the cup.

Zinc is added gradually to prevent excess frothing of the mixture and possibly frothing over of the reaction mixture. 

The temperature of the mixture is taken every 10-20s then the zinc is added and a graph of temperature against time is plotted. (see below)   

A colour change from blue solution (due to Cu²⁺ ions) to colourless solution (due to Zn²⁺ ions) is observed since excess zinc is used. 

The reaction mixture is stirred during the addition of the zinc to ensure all the zinc that can react does so and no zinc remains coated in copper.

A brown coloured finely divided solid appears in the solution at the end of the reaction.  

Let’s have a look at a typical set of results from this experiment that you can find in the internet.

Here it is:

So for the Brits among you “styrofoam” is Americanese for expanded polystyrene. 

A “single replacement reaction” is, I think, another Americanese for the displacement or better redox reaction that occurs since zinc being a more reactive metal than copper displaces it from the solution. 

Experimental Analysis

Let’s examine the quantities used to see which reagent is in excess. 

Above you can see that in the chemical equation 1 mole of zinc combines exactly with 1 mole of copper ions. 

Reacting quantities

How many moles of zinc is used in this experiment?

We are told the mass of powdered zinc was 1.3g so the number of moles was 1.3/65 or 0.02moles.

How many moles of copper ions were used in this experiment?

We are told again that 100ml of 0.1M copper sulphate solution was used.

So moles copper ions =  0.1  *100  =  0.01 moles copper ions

                                            1000

So in this reaction the zinc is in excess.

With the zinc in excess we can be sure that all the copper ions will be reduced to copper atoms. 

The enthalpy change will be determined by the amount of copper ions available in the reaction i.e. the limiting molar quantity.

Calculation of enthalpy change

How much energy was evolved during the reaction?

Here we use the equation with which you ought to be familiar that is:

E = m c ΔT

Therefore E  =  mass of solution used *  specific heat capacity of water * temp change

Therefore  E  =  100g   *  4.18 J/g/oC  *   11^oC 

[I don’t think we can get a temp change better than that (2 sig figs) given the lack of fine detail in the graph above]

So    E   =   4598J

But this number of Joules of energy are produced from 0.1moles copper ions

Therefore the enthalpy change of reaction is 4598J   =   45980J 

                                                                                      0.1mol

or –46kJ/mol

This is a poor result when you consider that the enthalpy change is accepted as  –217kJ/mol

You need to ask why this value for the enthalpy change is so low!!

The answer you ought to come up with is probably to do with poor insulation of the reaction in the cup.  May be no lid was used.

And may be the zinc was poor quality and some was already oxidised so not enough actual zinc atoms went into the reaction with copper ions. 

Here is another set of actual results for you to determine the enthalpy of reaction for yourself:

25.00cm³ of 1.00M copper(II) sulphate solution was placed in a polystyrene cup.  About 6g of powdered zinc were added which is an excess of zinc.  A temperature rise of 50.6^oC was recorded.

Can you show that the enthalpy change measured in this experiment is

 –212kJ/mol?

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