In the electrolysis of acidulated water, it is desired to obtain 1.12 cc of hydrogen per second under STP condition. The current to be passed is:

Options

A. `1.93A`
B. `9.65`
C. `19.3A`
D. `0.965`

Explanation

To solve the problem of determining the current required to obtain 1.12 cc of hydrogen per second during the electrolysis of acidulated water, we can follow these steps:

### Step 1: Understand the volume of hydrogen produced
At standard temperature and pressure (STP), 1 mole of any gas occupies 22.4 liters (or 22,400 cc). We need to find out how many moles of hydrogen correspond to 1.12 cc.

\[ \text{Moles of } H_2 = \frac{1.12 \text{ cc}}{22400 \text{ cc/mol}} = \frac{1.12}{22400} \text{ mol} \]

### Step 2: Calculate the moles of hydrogen
Calculating the moles gives us:

\[ \text{Moles of } H_2 = \frac{1.12}{22400} = 5 \times 10^{-5} \text{ mol} \]

### Step 3: Determine the charge needed for the reaction
From electrolysis, we know that 1 mole of \( H_2 \) requires 2 moles of electrons (since the reaction involves the reduction of \( 2H^+ + 2e^- \rightarrow H_2 \)). Therefore, the moles of electrons needed can be calculated as:

\[ \text{Moles of electrons} = 2 \times \text{Moles of } H_2 = 2 \times 5 \times 10^{-5} = 1 \times 10^{-4} \text{ mol} \]

### Step 4: Calculate the total charge required
Using Faraday's constant (approximately 96500 C/mol), we can find the total charge (Q) required:

\[ Q = \text{Moles of electrons} \times \text{Faraday's constant} = 1 \times 10^{-4} \text{ mol} \times 96500 \text{ C/mol} \]

\[ Q = 9.65 \text{ C} \]

### Step 5: Calculate the current
Using the formula \( Q = I \times t \), where \( t = 1 \text{ second} \):

\[ I = \frac{Q}{t} = \frac{9.65 \text{ C}}{1 \text{ s}} = 9.65 \text{ A} \]

### Final Answer
Thus, the current that needs to be passed to obtain 1.12 cc of hydrogen per second under STP conditions is **9.65 A**.

---