Two bulbs are connected by a stopcock. The large bulb, with a volume of 6.00 L, contains nitric oxide at a pressure of 0.850 atm, and the small bulb, with a volume of 1.50 L, contains oxygen at a pressure of 2.50 atm. The temperature at the beginning and the end of the experiment is 22∘C . After the stopcock is opened, the gases mix and react.

2NO(g)+O2(g)→2NO2(g)

1. Which gases are present at the end of the experiment?
2. What are the partial pressures of the gases? If the gas was consumed completely, put 0 for the answer.

(I). The gases are present at the end of the experiment  are O₂ and NO₂

(II).  The pressure of O₂ is 0.158 atm.

The pressure of NO₂ is 0.681 atm

The Pressure of NO is zero.

Explanation:

Given that,

Volume of large bulb = 6.00 L

Pressure = 0.850 atm

Volume of small bulb = 1.50 L

Pressure = 2.50 atm

Temperature = 22°C = 295 K

We need to calculate the moles in NO

Using formula of moles

Put the value into the formula

We need to calculate the moles in O

Using formula of moles

Put the value into the formula

The balance equation for the reaction is

So, The gases are present at the end of the experiment  are O₂ and NO₂

We need to calculate the remaining moles of O₂

Using formula for remaining moles

Moles of  O₂ remaining =

Moles of NO₂ = 0.211 moles

Total volume

Put the value into the formula

(2). If the gas was consumed completely

We need to calculate the pressure of O₂

Using formula of pressure

Put the value into the formula

We need to calculate the pressure of NO₂

Using formula of pressure

Put the value into the formula

If the gas was consumed completely

Then, Pressure of NO is zero.

Hence, (I). The gases are present at the end of the experiment  are O₂ and NO₂

(II).  The pressure of O₂ is 0.158 atm.

The pressure of NO₂ is 0.681 atm

The Pressure of NO is zero.

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