Question

The reform reaction between steam and gaseous methane (CH4) produces "synthesis gas," a mixture of carbon monoxide gas and dihydrogen gas. Synthesis gas is one of the most widely used industrial chemicals, and is the major industrial source of hydrogen. Suppose a chemical engineer studying a new catalyst for the reform reaction finds that 924. liters per second of methane are consumed when the reaction is run at 261.°C and 0.96atm. Calculate the rate at which dihydrogen is being produced.

Answer 1

The answer is "
`= 0.078 \ kg \ H_2`".

Step-by-step explanation:

calculating the moles in
`CH_4 =(PV)/(RT)`

`=((0.58 \ atm) * (923 \ L) )/( (0.0821 (L \cdot atm)/(K \cdot mol))(232^(\circ) C +273))\\\\=((535.34 \ atm \cdot \ L) )/( (0.0821 (L \cdot atm)/(K \cdot mol))(505)K)\\\\=((535.34 \ atm \cdot \ L) )/( (41.4605 (L \cdot atm)/(mol)))\\\\= 12.9 \ mol`

Eqution:

`CH_4 +H_2O \to 3H_2+ CO \ (g)`

Calculating the amount of
`H_2` produced:

`= 12.9 \ mol CH_4 * (3 \ mol \ H_2 )/(1 \ mol \ CH_4)* (2.016 g H_2)/(1 \ mol \ H_2)\\\\= 78 \ g \ H_2 \\\\= 0.078 \ kg \ H_2`

So, the amount of dihydrogen produced =
`0.078 (kg)/(s)`