Later on, the graph was plotted.
Graph 1: Graph showing the course of the measurement
Using informations from the table and the extrapolations from graph, I am able to calculate the enthalpy change.
∆T= temp.max – temp.min
∆T= 59.8 – 30.5
∆T= 29.3ºC
Q= m·c·∆T where m mass of solution (g)
c specific heat capacity of H2O (J g–1 K–1)
∆T change of temperature (K)
Mass of water=100 cm3 * 1 g cm3= 100 g
Qwater= m·c·∆T
Qwater=100g* 4.18 J g–1 K–1* 29.3 K=12247.4 J
From the first law of thermodynamics we get the equation:
Qwater= Qice
So the Q of ice is equal to 12247.4 J, so 12.2474 kJ.
What is more, we know that the heat gained by ice was used to melt the ice first and then to raise the temperature of the whole solution:
Qice=Qfus + (∆Tice ·c ·mice)
Qfus=Qice  (∆Tice ·c ·mice)
Qfus=12247.4 – 30.5 K* 4.18 J g–1 K–1 * 17.9g
Qfus=  14529.71 J
1 mole H2O ice — x
0.99 moles of H2O ice —  14529.71 J
x= 14676.47 J mol1
The enthalpy change for this reaction is 14676.47 J mol1.
Part B: enthalpy change for displacement reaction
First step in the experiment was to create 1M copper sulphate. To do this solution, some calculations were needed:
mCuSO4 · 5H2O = =250 g
to create 50 cm3 1M solution I need:
x=12.5g of copper sulphate
So I dissolve 12.5 g copper sulphate in 50 cm3 of water.
Also, I calculated how many moles of zinc I am going to add to the solution.
MZn=65g
1 mole — 65 g
x— 5 g
x=0.077 moles of Zn
Since I don't know, I decided to calculate how many moles of zinc are actually going to react with copper sulphate.
The equation for the reaction looks as follow:
CuSO4 (aq) + Zn (s) → ZnSO4 (aq) + Cu (s)
or
Cu2+(aq) + Zn (s) → Zn2+ (aq) + Cu(s)
1 mole — 250 g
x— 12.5 g
x=0.05 moles of CuSO4
As the proportion between Zn and CuSO4 is 1:1,
1 mole CuSO4 — 1 mole Zn
0.05 moles — x
x=0.05 moles of Zn will actually react
Which means that zinc is in excess.
Next step was to recordfirstlythe temperature of the solution, then add zinc.
First two minutes of the temperature was measured for the solution of CuSO4 only.
Then, the powdered zinc was added and the solution was stired rapidly until the maximum temperature has been reached ( which happened in 3.5 minute).
Later on, the graph was plotted.
Graph2: Graph showing the course of the measurement
Using information from the table and extrapolations, I am able to calculate the enthalpy change.
∆T= temp.max – temp.of CuSO4
∆T= 70.018.8
∆T= 51.2ºC
Q= m·c·∆T where m mass of solution (g)
c specific heat capacity of H2O (J g–1 K–1)
∆T change of temperature (K)
m of solution=50 cm3 * 1 g cm3= 50 g
∆H=50 g* 4.18 J g–1 K–1 * 51.2 K =  10700 J
1 mole CuSO4 — x
0.05 moles of CuSO4 — 10.7 kJ
x= 214.0 kJ mol1
The enthalpy change for this reaction is 214.0 kJ mol1.
Also, I calculated the total error that might occur in the experiment due to uncertainities of the equipment and measurements:
Error= 0.5+0.05+0.01= ± 0.56
CE:
The first part of the experiment was to determine the heat required for the change of state, commonly known as the latent heat. Using the first law of thermodynamics and the Laplace's law which say that "the heat exchange accompanying a transformation is equal and opposite to the heat exchange accompanying the reverse transformation" and was indicated in the first part of experiment as Qwater= Qice. It means that the heat gained by ice from the water was firstly used to melt the ice and raise the temperature of water. Using the equations and the Hess's law I calculated the enthalpy change and determined the latent heat, the results that satisfy me.
As for the second part of experiment which usage of zinc and copper (II) sulphate, measurements made allowed me to create the graphs and calculate the enthalpy changes, which was the aim of the investigation. Knowing the commonly accepted value of enthalpy change for the displacement reaction (taken from Chemistry for IB 2nd edition Sadru Damji) is 217 kJ mol1 and comparing it with result of the experiment I calculated the error, which in fact was very small as its value was 1%.
error =
error=
error=1.38%
This lead me to conclusion that the method of collecting data and the way I processed it was accurate and very satisfying. Extrapolations on the graph allowed me to use precise values of measured datas and therefore to calculate the enthalpy changes.
Errors occured during the experiment and might have affected the final results:
 wrong or error in calculations to create the solution
 balance with only 2 decimal places
 the uncertainity of cylinders
 human reaction during measurements [0.15s]
 adding powder in not exact time
 polystyrene cup might not have been isolated all the time, so loss of energy could occur
 some zinc powder stayed on the Vernier coach thermometer

the value of heat capacity differs; I assumed it was 4.18 g–1 K–1
To improve the experiment, I might make some changes:
 make more readings to reduce the uncertainity
 more stirring during the experiment, to make sure all substances dissolved
 use very icy ice (taken straight from the fridge) so it wouldn't melt too soon
 be more accurate while measuring add all substances at the exact time
 use more accurate balance and weigh all substances twice

use thermometer and Vernier coach to make sure that temperature recorded is the same as measured
 have more time to do more measurements, prelong the time of measurements