hess exemplar lab with comments

5/28/2018 Hess Exemplar Lab With Comments

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1 Dr. Geoffrey Neuss,

SL & HL Chemistry practical on Hesss Law (with comments)

Aim

The purpose of this experiment is to determine the indirect enthalpy change for the hydration of

copper(II) sulphate:

CuSO4(s) + 5H2O(aq) CuSO4.5H2O(s)

[anhydrous] [hydrated]

A minor point but the state symbol for water should be (l) not (aq).

If anhydrous copper(II) powder is left in the atmosphere it slowly absorbs water vapour giving the

hydrated solid (see above equation).

Heat is evolved in the reaction but is difficult to measure by the direct method. So instead it is

possible to measure the heat changes directly when both anhydrous and hydrated copper(II)

sulphate are separately dissolved in water and then construct an energy cycle (see below) to

determine the required H value directly.(The aim, as stated above, procedure and an overview of

this whole experiment can be found on p15 of Higher Level Chemistry Practical Book and p29 -

"Chemistry for the IB Diploma".)

Energy cycle

Results


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See graph for exothermic reaction of anhydrous copper(II) sulphate (H1) :

T2 = Temperature that would have been reached if no heat lost to surroundingT

1= Highest temperature actually reached

To= Initial temperature of reactant

The student has not recorded raw data in a table but just presented it in a graph. She has not

included the temperature for the two minutes before the copper(II) sulphate was added to the

water so there is no way of knowing whether the initial solution was at room temperature. On the

graph she has used subscripts for T1 and T2 but in the text she has used superscripts. The graph has

a title but neither of the two axes have labels of time or temperature and no units are given.

Change in temperature for reaction = T2To= 25.620.0

= 5.6oC

Thus we must consider the heat is lost by water to the environment during the exothermicreaction so to compensate we can plot a graph of temperature against time: then by

extrapolating the graph, the temperature rise that would have taken place had the reaction

been instantaneous can be shown. Working out the change in temperature is necessary to

calculate the enthalpy change in reaction.


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For the endothermic reaction of hydrated copper(II) sulphate (H2) :

It is not necessary for a graph to be plotted as we only need the temperature change fromthe initial temperature to the lowest point.

So the change in temperature for reaction

= Initial temperaturelowest temperature

= 20.0 - 17.0

= 3.0oC

The student is correct that no graph is needed as the temperature change is so small. Normally it is

only about one degree Celsius so she has a slightly larger drop in temperature than is usually

obtained.

Processing data

Reaction(H1) :

1). We have already found the change in temperature (above) to be : 5.6oC

2). Next we can find the heat evolved in the experiment for 0.025 mol of copper(II) sulphate:

Heat evolved = (50.0 / 1000) x 4.18 x 5.6oC = 1.17 KJ

3). Finally we can calculate the enthalpy change for the reaction H:

=1.17 x 1 / 0.025 = 46.8 KJ mol-1

Nowhere has she recorded her mass of salt taken or the uncertainty associated with it or how she

measured the volume of water and its uncertainty. She has wrongly given the units of heat as KJ

instead of kJ. The use of significant figures is inconsistent.

Reaction(H2) :

1). We have already found the change in the temperature (above) to be: 3.0oC

2). Next we can find the heat evolved in the experiment for 0.025 mol of copper(II) sulphate:

Heat evolved = 50.0 [although we used (18/1 x 0.025) x 5 = 2.25

= 50.02.25 = 47.75 cm3

This is to compensate for the water which comes from the hydrated salt]

Thus (50.0 / 1000) x 4.18 x 3.0oC = 0.63 KJ

3). Finally we can calculate the enthalpy change for the reaction H:

= 0.63 x 1/0.025 = 25.2 KJ mol-1

She talks about 'heat evolved' when in fact it was taken in although higher up she does state

correctly that it is an endothermic reaction. There is some confusion in the way the amount of water

taken has been calculated in the midst of calculating the heat change. Again significant figures are

not consistent.


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So overall:Hxf= H1H2

=46.825.2

=72.0 KJ mol-1

Qualitative observations of data:

The anhydrous copper(II) sulphate powder was initially white and then, when dissolved inwater turns light blue.

The hydrated copper(II) sulphate was initially light blue, and continues to stay light bluethroughout the reaction.

Conclusion

We can calculate theoretically the enthalpy change for the hydration of anhydrous copper(II)

sulphate by Hess' law using an energy cycle and the values:

H fCuSO4 = 770 KJ mol-1

H fCuSO4.5H2O = 2278 KJ mol-1

H fH2O = 286 KJ mol-1

Therefore Hx= 5(286) + 770 2278

=78 KJ mol-1

The student has used the data from the data book to correctly work out the literature value and

then goes on to use it to calculate correctly the percentage error. She should probably have

referenced the source of her data.

I can now look at my results and find the final value for the reaction H for the hydration of

anhydrous copper(II) sulphate to be: 72.0 KJ mol-1

, which considering the theoretical value is

reasonably accurate. But I can find the precise accuracy by calculating the percentage error:


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= 72.0 78.0 x 100

78.0

= 7.69%

This shows that the result I achieved (i.e. the experimental value) is actually quite similar to the

theoretical value and although there is always going to be a certain amount of uncertainty in an

experiment my result was accurate.

We can see that after the reactants are mixed (i.e. the copper(II) sulphate added) in H1, the initial

reaction gives out heat and therefore energy; although after the original rise in temperature for the

reaction we see a gradual loss in heat. This initial heat energy loss is what we were aiming to

calculate for this exothermic reaction, as heat is always given out to the surroundings during

exothermic reactions because the bonds in the products are stronger than the bonds in the

reactants. The original heat gain in the experiment is the enthalpy change for the reaction and

ultimately necessary to continue our calculations for the enthalpy change for the hydration ofcopper(II) sulphate indirectly.

Oppositely in the H2reaction, the initial reaction takes in heat and therefore energy, thus initiating

a decrease in temperature. This is because it is an endothermic reaction and needs to absorb heat

from the surroundings because the bonds in the reactants are stronger than the bonds in the

products. So we measure the decrease in temperature and use this to calculate the enthalpy change

in the reaction.

Finally, having calculated the enthalpy change values from both reaction H1 andH2,we could

calculate the overall enthalpy change by indirect means; hence using an energy cycle and Hess' Law

experimentally.

Evaluation

There are obviously going to be ways to improve any uncertainties in my experiment, despite being

only 7.69% percentage error and ultimately develop my ideas and conclusion. For example, we do

not know if the substances were completely pure and this could affect the accuracy, even if it is only

slightly and so we have to take this into account and try to make them as pure as possible. For

example the anhydrous copper(II) sulphate, before starting should be completely white to have total

pureness. Another uncertainty we have to included is the difference between the water content of

anhydrous and hydrated copper(II) sulphate, I did try to calculate as accurately as possible how

much less water we needed to use that was already compensated for in the hydrated copper(II)sulphate but it is still not completely precise. We used a data logger which would have processed the

data directly as soon as the readings were taken, and also an automatic stirrer to keep it equal

throughout so in a sense the experiment is very accurate but there are other factors we have to

acknowledge.

There is always going to be heat loss, even through the insulation of the polystyrene cup; so in order

to advance my investigation we could experiment with different materials as insulators. Thus I would

be testing if heat loss could be reduced by another more insulating material and hopefully creating a

decrease in percentage error. But overall my experiment was successful and I achieved the

experimental results necessary for my investigation to match the calculated theoretical value to a

reasonable degree.


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The conclusion does not relate the error with the uncertainties as nowhere have the uncertainties

been given or estimated. The evaluation does recognise that the anhydrous salt had probably

already absorbed some water from the atmosphere but she should have gone on to say how this

would have affected the temperature rise obtained. She repeats herself a lot without really saying

much and has missed the fact that it has been assumed that the solution has the same specific heatcapacity as pure water. There are no realistic suggestions, such as drying the anhydrous salt in an

oven beforehand, to improve the experimental method.

Marks awarded:

DCP: N (Although some qualitative raw data was given no appropriatequantitative raw data was

given.)C (The quantitative raw data has been processed correctly.)

N (The value found does not contain any uncertainties and nowhere have these been

presented.)

Total 2

CE: C (The percentage error has been correctly given. It has not been compared with the

uncertainties

but the lack of these has already been penalised in aspect 3 of DCP.)

P (Some weaknesses were identified)

P (Could be almost N as only very superficial improvements have been suggested)

Total 4

hess exemplar lab with comments

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