Complete reaction of 2.60g of chromium metal with 50.00mL of 1.200M H2SO 4 in ab

Question

Complete reaction of 2.60g of chromium metal with 50.00mL of 1.200M H2SO4 in absence of air gave a blue solution and a colorless gas that was collected at 25C and a pressure of 735 mmHg.

a. Write a balanced net ionic equation for the reaction.

b. How many liters of gas were produced?

c. What is the pH of the solution?

d. Describe the bonding in the blue colored ion using both the crystal field theory and the valence bond theory. Include the appropriate crystal field d-orbital energy level diagram and the valence bond orbital diagram. Identify the hybrid orbitals used in the valence bond description.

e. When an excess of KCN is added to the solution, the color changes and the paramagnetism of the solution decreases. Explain.

Explanation / Answer

a) 2 Cr + 3 H2So4 = Cr2(So4)3 + 3 H2

b) 2.60g of chromium metal with 50.00mL of 1.200M H2SO4 , 25C and a pressure of 735 mmHg.

As per the stoichiometry 2 moles of Cr will react with 3 moles of H2SO4 to give three moles of Hydrogen

atomic weight of Cr = 52

104g of chromium will react with 196g H2SO4 to give 6g of hydrogen

so 2.6g of chromium will react with 4.9 g H2SO4 to give 0.15g of hydrogen

so moles of hydrogen formed = mass / molecular weight = 0.15 /2 = 0.075 moles

So volume of hydrogen formed will be

PV = nRT

Volume = 0.075 X 62.363 X 298 / 735 mm Hg

Volume = 1.89 L

c) Moles of H2SO4 reacted = mass of H2SO4 / molecular weight = 4.9 / 98 = 0.05 moles

Initial moles of H2SO4 = 1.2 X 50 / 1000 = 0.06 moles

so moles of H2SO4 left = 0.01 moles

concentration of H2SO4 = 0.01 moles / 0.05 L = 0.2 M

so concentration of H+ = 0.4 M

so pH = -log [H+] = -log [0.4] = 0.39

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