Standard Enthalpy Change in Reaction Between Zinc and Copper Sulphate

Raw data

Time/s(±0.5s)

Temperature/°C(±0.5°C)

0 19.0

30 19.0

60 19.0

90 19.0

120 19.0

150 19.0

180 19.0

210 (zinc added)

29.0

240 28.0

270 32.0

300 37.0

330 38.0

360 41.0

390 42.0

420 44.0

Time/s(±0.5s)

Temperature/°C(±0.5°C)

450 46.0

480 48.0

510 48.0

540 49.0

570 48.0

600 48.0

630 47.0

660 46.0

690 45.0

720 45.0

750 44.0

780 43.0

810 42.5

840 42.0

870 41.0

900 40.0

ObservationsAt zero seconds, the copper sulphate was a bright blue aqueous solution, while the zinc was a fine black powder. When the reactants were added (at 210 seconds), the solution violently effervesced for a few seconds before becoming still again, and the polystyrene cup calorimeter became warm to the touch. At 900 seconds the cup’s contents began to cool at a steady rate. The products were a colorless solution (which I infer to be zinc sulphate) and a black precipitate (which I infer to be excess zinc powder, since solid copper is brown).

The reaction can be described by the following chemical equation:

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

AnalysisThe data can be presented in the form of a graph (see attached sheet), with time along the x-axis (independent variable) and temperature along the y-axis (dependent variable).

The formula for standard enthalpy change of a reaction is as follows:

Q=mCΔT

Here, Q is enthalpy change (J), m is mass (g), C is specific heat capacity (J °C-1 g-1), and ΔT is the change in temperature (°C).

In this investigation, Q is the variable being found. m is the mass of the substance being heated (25.0cm3 ± 0.05cm3 of 1M copper sulphate solution). Though the atoms of copper sulphate were also heated, the primary substance being heated was the water, so it is assumed that the mass of water is equal to its volume (as the density of water is 1g/cm3). Thus, m is 25.0g ± 0.05g = 25.0g ± 0.200%.

C is the literature value for the specific heat capacity of water, 4.18 J °C-1 g-1. It is assumed to be exact and without error because it is a literature value.

The minimum temperature recorded was 19°C (±0.5°C) and the maximum temperature was 49°C (±0.5°C). Thus,

ΔT = 49.0 – 19.0 = 30.0°C

The percentage uncertainty is given by:

∑ ( ∆A ÷ A ) =0.549.0

+0.519.0

= 0.0365198… ≈ 3.65%

Thus,

ΔT = 30.0°C ± 3.65%

The change in enthalpy for the reaction between zinc and copper sulphate is therefore

Q = mC∆T = 25.0 g × 4.18 J ° C-1 g-1 × 30.0 °C = 3135 J ± 3.85% = 3135 J ± 121

The reaction is exothermic (energy is released from the reactants), so the enthalpy change should be negative. Thus,

Q = -3135 J ± 121

The reference value for this reaction is in J mol-1, so to compare my value I must calculate the molar enthalpy change of the reaction.

Molar enthalpy change = joules ÷ number of moles = -3135 J ± 3.85% ÷ (concentration × volume)

= -3135 J ± 3.85% ÷ (1 mol dm-3 × 25cm3 ± 0.2%)

= -3135 J ± 3.85% ÷ 0.001 mol cm3 ÷ (25cm3 ± 0.2%)

= -125400 J mol-1 ± 4.05% = -125.4 kJ mol-1 ± 4.05% = -125.4 kJ mol-1 ± 5.08 kJ mol-1

The reference value of the molar enthalpy change for this reaction is -217 kJ mol -1, which is a substantial amount higher than my experimental value. My percentage error is given by

(-217 )- (-125.4 )-217

= 0.4221198157 … = 42.2%

This is far from being within my supposed uncertainty range of ± 5.08 kJ mol -1. This could suggest that the source of error lies in my methodology.

EvaluationPerhaps the largest source of systematic error was the heat energy lost to the surroundings by conduction or convection. In my observations, I stated that the polystyrene cup calorimeter was warm to the touch, and because it was a poor insulator, it lost large amounts of heat to the surroundings through convection, conduction, and radiation. Losing heat to the environment instead of the thermometer would lower the value of the maximum temperature, decreasing the change in temperature, lowering the enthalpy and the molar enthalpy. Though this error could have been slightly mitigated by the plastic lid I placed over the cup, I do not believe it sufficiently reduced the error as intended. The heat lost to the surroundings could have been reduced by insulating the cup through various means (such as wrapping it with a thick insulating material such as furry fabric) or placing another plastic cup around it.

This systematic error can be seen on the graph; extrapolating to the point where the line of best has the same x-value as the zinc was added (210 seconds) gives a large area above the plotted points. This can be seen as the heat energy lost to the environment that was thus excluded from my measurements. Furthermore, the y-value of the line of best fit at the corresponding x-value of when the zinc was added (210 seconds) is the theoretical maximum temperature. I should have used this value to calculate an experimental and a theoretical value for the reaction, giving me a clearer idea of where my value was in relation to the literature value.

Another source of error was the fact that I only carried out the experiment with small quantities of copper sulphate and zinc. I could have varied the amount of copper sulphate and zinc (in fixed ratios), calculated the molar enthalpies of each one, and averaged to give a more suitable value. Using larger quantities would also reduce the impact of random errors on my result by reducing the percentage uncertainty of measurements. Finally, I only conducted the experiment once, without any repeats. Thus, the effect of random errors on my calculation of molar enthalpy would not be reduced, making it less accurate.