Introduction: The definition of a mole is Avogadro’s number (6.02 x 1023) of particles (atoms, molecules, ions, electrons etc.). Moles are a very important part of chemistry especially in stoichiometry since it is part of many other calculation quantities and formulas including molar mass, solution calculations and gas volume calculations.
The mole is also used in chemical reactions and equations to calculate the amount of reactant needed to react completely with another reactant or to calculate the product produced from the amount of reactant provided and vice versa. This is done by using the ratio of the coefficients in a balanced equation. This ratio of coefficients is also known as the mole ratio.
In the following experiment, a simple displacement reaction would occur from the reaction of an aqueous solution of copper (II) sulphate and zinc powder.
Zn (s) + CuSO4 (aq) ï¿½ ZnSO4 (aq) + Cu (s)
This reaction would be set up to allow the zinc to be the limiting factor therefore react completely, in order for that to happen, copper (II) sulphate would be in excess. As zinc is the limiting factor, it will be used to calculate the expected amount of copper produced from the 1 to 1 mole ratio of zinc and copper from the balanced equation above.
Aim: To find the mole ratio of a reactant to a product in a chemical reaction .
Apparatus: – Balance – Bunsen Burner
– Two 150 cm3 Beaker – Glass Stirring Rod
– Tripod – Gauze
– 100 cm3 Graduated Cylinder – Goggles
– Heat Proof Mat – Pure Distilled Water
– Tongs – Balance to 0.01 g
Reagents: – Copper (II) Sulphate Crystals
– Zinc Powder
1. Weigh the mass of a clean, dry 150 cm3 beaker. Then weigh out 7.0 g of copper (II) sulphate using the beaker
2. Add 50.0 cm3 of pure distilled water into the beaker and heat the solution gently until all the copper (II) sulphate has dissolved
3. Determine the mass of the second clean, dry 150 cm3 beaker. Then weigh out as accurately as possible, 1.30 g of zinc powder using the beaker
4. Record the mass of the beaker and the zinc powder in the results table, nearest to 0.01 g
5. Slowly pour the copper (II) sulphate into the beaker containing the zinc. Stir continuously for 1-2 minutes
6. Leave the beaker for 10 minutes while the reaction continues. Record your observations
7. When the copper has settled, pour out the light blue liquid.
8. Add 10 cm3 of pure distilled water into the beaker.
9. Leave it for 10 minutes again and pour out the liquid again
10. Repeat steps 8 and 9, two times
11. Place the beaker in the oven to dry for 24 hours
12. Remove the beaker from the oven and determine the mass.
1. Mass of empty 250 cm3 beaker
2. Mass of the 250 cm3 beaker and copper (II) sulphate
3. Mass of copper (II) sulphate
4. Mass of empty 150 cm3 beaker
5. Mass of 150 cm3 beaker and zinc powder
6. Mass of zinc powder
7. Mass of the 150 cm3 beaker and copper formed (after drying overnight)
8. Mass of copper
1. Number of moles of copper produced
= 0.0222 mol
2. Number of zinc moles reacted
= 0.0199 mol
4. Mass of copper expected
n(Cu) = 1 x 0.0199
= 0.0199 mol
m = nM
= 0.0199 x 63.55
= 1.26 g
5. Percentage Yield:
Conclusion: The mole ratio from calculation 3 is approximately 1 to 1, same as the expected mole ratio from the balanced equation. The expected mass of copper is 1.26 g but 1.41g of copper was weighed out therefore, the percentage yield of the above experiment is 89.4%. This is mostly caused by the impurity of the copper since there might be a small amount of leftover zinc sulphate in the beaker.
Evaluation: From the experiment above, some things could have been done better to achieve a higher percentage yield. The glass rod should’ve been dipped into water before stirring the copper (II) sulphate and zinc in order for no copper to attach to the glass rod after stirring. More importantly, I could have improved and made this experiment more accurate by rinsing the copper more thoroughly so no zinc sulphate would be left in the beaker.