Question

A dichromate solution is standardized by titrating it with a0.1340 M Fe2+ solution. The balanced net ionic equation for the reaction is: Cr2O72-(aq) + 6Fe2+(aq)+14H3O+(aq)ebab9d75-6e9a-4c80-9bcd-b491b668f097.gif2Cr3+(aq) + 6Fe3+(aq)+21H2O(l)
If 15.15 mL of the 0.1340 M Fe2+solution is required to react completely with 40.00 mL of the dichromate solution, calculate the concentration of the dichromate solution. _____M

please show how you started and the steps! thank you!

Answer 1

Answer:

8,459x10⁻³ M of Cr₂O₇²⁻

Explanation:

The equation for the reaction of Fe²⁺ with Cr₂O₇²⁻ is:

Cr₂O₇²⁻(aq) + 6Fe²⁺(aq)+14H₃O⁺(aq) → 2Cr³⁺(aq) + 6Fe³⁺(aq)+21H₂O(l)

The moles of Fe²⁺ that you required for a complete reaction of Cr₂O₇²⁻ are:

0,01515 L ×$\frac{0,1340moles}{L}$ = 2,0301x10⁻³ moles of Fe²⁺

By the equation of the reaction, 1 mol of Cr₂O₇²⁻ reacts with 6 moles of Fe²⁺, thus, moles of Cr₂O₇²⁻ are:

2,0301x10⁻³ moles of Fe²⁺×$\frac{1molCr_{2}O_{7}^{2-}}{6molesFe^{2+}}$ = 3,3835x10⁻⁴ moles of Cr₂O₇²⁻

The molarity is:

$\frac{3,3835x10^{-4}moles}{0,04L}$ = 8,459x10⁻³ M of Cr₂O₇²⁻

I hope it helps!