If `CO_(2)` is allowed to escape from the following reaction at equilibrium
`CO_(2)+H_(2)O+H_(2)CO_(3)hArr 2H^(+) + 2HCO_(3)^(-)`

Options

A. pH will decrease
B. pH will remain constant
C. pH will increase
D. forward reaction will be favoured

Explanation

To solve the problem, we need to analyze the given equilibrium reaction and understand how the escape of CO₂ affects the equilibrium position. The reaction is:

\[ \text{CO}_2 + \text{H}_2\text{O} + \text{H}_2\text{CO}_3 \rightleftharpoons 2\text{H}^+ + 2\text{HCO}_3^- \]

### Step 1: Identify the Equilibrium Condition
At equilibrium, the rates of the forward and reverse reactions are equal. This means that the concentrations of the reactants and products remain constant over time.

### Step 2: Understand the Effect of CO₂ Escape
When CO₂ is allowed to escape from the system, the concentration of CO₂ decreases. According to Le Chatelier's principle, if a change is made to a system at equilibrium, the system will adjust to counteract that change and restore a new equilibrium.

### Step 3: Apply Le Chatelier's Principle
Since CO₂ is a reactant in the equilibrium reaction, removing it (by allowing it to escape) will shift the equilibrium to the left (towards the reactants) to compensate for the loss of CO₂. This means that the reaction will favor the formation of more CO₂ and H₂CO₃.

### Step 4: Write the Direction of Shift
The shift in equilibrium can be summarized as follows:
- The removal of CO₂ causes the equilibrium to shift to the left.
- This results in an increase in the concentration of CO₂ and H₂CO₃ while decreasing the concentrations of H⁺ and HCO₃⁻.

### Final Conclusion
Thus, if CO₂ is allowed to escape from the equilibrium reaction, the reaction will shift to the left.

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