In: Chemistry
I'm a little stumped of this one...
IV) 2CrO4 –2(aq) + 2H+ (aq) Cr2O7 –2(aq) + H2O(l)
(7) *Add 2 drops of 6M HNO3 to about 3mL of 0.1M K2CrO4.
(8) Add 10% NaOH to (7) dropwise until the original color of K2CrO4 is restored.
(9) *Add 2 drops of 6M H2SO4 to about 3mL of 0.1M K2CrO4.
(10) Add 10% NaOH to (9) dropwise until the original color of K2CrO4 is restored.
Question:
In step (10), which direction DID the equilibrium shift? Using balanced equations, explain the direction of this shift.
2CrO42-(aq) + 2H+ (aq) ⇌ Cr2O72-(aq) + H2 O(l)
In step (10), the equilibrium shift towards the left (reactants).
Explanation- In step (9) 2 drops of 6M H2SO4 is added to 3mL of 0.1M K2CrO4. Originally the solution is yellow (more CrO42-) as H+ are added to the solution, the equilibrium shifts towards the products, resulting in the solution turning orange (more Cr2O72-). When NaOH is added to the solution in step (10), it neutralizes the acid; removing H+ ions and equilibrium sifted towards left (the reactants) and the solutions turns yellow (more CrO42-) again.
Balanced equations for step (10)
Cr2O72- + 2 OH- <====> 2 CrO42- + H2O