25 ml of 0.50 M H2O2 solution is added to 50 ml of 0.20 M KMnO4 in acidic solution. Which of the following statement(s) is/are true?

To answer this question, we first need to understand the reaction that is taking place. The reaction between hydrogen peroxide (H2O2) and potassium permanganate (KMnO4) in acidic solution is a redox reaction. The balanced chemical equation for this reaction is:

2 KMnO4 + 5 H2O2 + 6 H+ → 2 Mn2+ + 5 O2 + 8 H2O + 2 K+

Now, let's calculate the number of moles for each reactant:

• For H2O2: Moles = Molarity x Volume = 0.50 M x 25 ml = 12.5 mmol
• For KMnO4: Moles = Molarity x Volume = 0.20 M x 50 ml = 10 mmol

From the balanced chemical equation, we can see that the stoichiometric ratio between KMnO4 and H2O2 is 2:5. This means that for every 2 moles of KMnO4, we need 5 moles of H2O2.

However, in our case, we have 10 mmol of KMnO4 and 12.5 mmol of H2O2. If we divide these amounts by their respective stoichiometric coefficients, we get:

• For H2O2: 12.5 mmol / 5 = 2.5
• For KMnO4: 10 mmol / 2 = 5

Since the value for H2O2 is smaller, H2O2 is the limiting reactant. This means that all of the H2O2 will be consumed in the reaction, and there will be some KMnO4 left over.

So, the statement(s) that is/are true would depend on the options given, but based on the information above, any statement that says H2O2 is the limiting reactant or that there will be leftover KMnO4 would be true.