How many ml of 1 (M) `H_(2)SO_(4)`is required neutralise 10 ml of 1 (M) NaOH solution?

Options

A. (a) 2.5
B. (b)5.0
C. (c )10.0
D. (d)20.0

Explanation

To determine how many milliliters of 1 M H₂SO₄ are required to neutralize 10 ml of 1 M NaOH solution, we can follow these steps:

### Step-by-Step Solution:

1. **Understand the Neutralization Reaction**:
The neutralization reaction between sulfuric acid (H₂SO₄) and sodium hydroxide (NaOH) can be represented as:
\[ H₂SO₄ + 2 NaOH \rightarrow Na₂SO₄ + 2 H₂O \]
From the equation, we see that 1 mole of H₂SO₄ reacts with 2 moles of NaOH.

2. **Determine the Number of Equivalents**:
For neutralization, the number of equivalents of acid must equal the number of equivalents of base. The number of equivalents can be calculated using the formula:
\[ \text{Number of equivalents} = \text{Normality} \times \text{Volume (in L)} \]

3. **Calculate Normality of H₂SO₄**:
The normality (N) of H₂SO₄ can be calculated as:
\[ \text{Normality of H₂SO₄} = \text{N factor} \times \text{Molarity} \]
Here, the N factor for H₂SO₄ is 2 (since it can donate 2 protons). Therefore:
\[ \text{Normality of H₂SO₄} = 2 \times 1 \, \text{M} = 2 \, \text{N} \]

4. **Calculate Normality of NaOH**:
The normality of NaOH is equal to its molarity because NaOH has an N factor of 1:
\[ \text{Normality of NaOH} = 1 \, \text{N} \]

5. **Set Up the Equation**:
Now we can set up the equation for the equivalents:
\[ \text{Normality of H₂SO₄} \times \text{Volume of H₂SO₄} = \text{Normality of NaOH} \times \text{Volume of NaOH} \]
Plugging in the values we have:
\[ 2 \, \text{N} \times \text{Volume of H₂SO₄} = 1 \, \text{N} \times 10 \, \text{ml} \]

6. **Solve for Volume of H₂SO₄**:
Rearranging the equation to find the volume of H₂SO₄:
\[ \text{Volume of H₂SO₄} = \frac{1 \, \text{N} \times 10 \, \text{ml}}{2 \, \text{N}} = \frac{10}{2} \, \text{ml} = 5 \, \text{ml} \]

### Final Answer:
5 ml of 1 M H₂SO₄ is required to neutralize 10 ml of 1 M NaOH solution.