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The problem involves a chemical equilibrium between carbon monoxide (CO), water vapor (H₂O), carbon dioxide (CO₂), and hydrogen gas (H₂):

\[
\text{CO(g)} + \text{H₂O(g)} \rightleftharpoons \text{CO₂(g)} + \text{H₂(g)}
\]

Given the initial amounts of substances, we are asked to complete a table that tracks the initial molarity, the change in molarity, and the equilibrium molarity of each species in a 250 mL flask.

### Step 1: Initial Molarity
Molarity is calculated as:

\[
M = \frac{\text{moles}}{\text{volume (L)}}
\]

For a 250 mL flask, the volume is 0.250 L. The initial moles of each substance are given:
- H₂O: 1.3 mol
- CO₂: 1.8 mol
- H₂: 0.30 mol
- CO: unknown (assuming we're solving for it, but likely 1.8 mol, based on a balanced reaction).

Now, we calculate their initial molarities.

- \[\text{Molarity of H₂O} = \frac{1.3}{0.250} = 5.2\ M\]
- \[\text{Molarity of CO₂} = \frac{1.8}{0.250} = 7.2\ M\]
- \[\text{Molarity of H₂} = \frac{0.30}{0.250} = 1.2\ M\]
- \[\text{Molarity of CO} = \frac{1.8}{0.250} = 7.2\ M\]

### Step 2: Change in Molarity
As the reaction progresses to equilibrium, CO and H₂O are consumed, and CO₂ and H₂ are produced. We represent the change in the molarity of CO as \(x\), so:

- CO decreases by \(x\): \(\Delta [CO] = -x\)
- H₂O decreases by \(x\): \(\Delta [H₂O] = -x\)
- CO₂ increases by \(x\): \(\Delta [CO₂] = +x\)
- H₂ increases by \(x\): \(\Delta [H₂] = +x\)

### Step 3: Equilibrium Molarity
At equilibrium, the molarity of each species is:

- \(\text{[CO]} = 7.2 - x\)
- \(\text{[H₂O]} = 5.2 - x\)
- \(\text{[CO₂]} = 7.2 + x\)
- \(\text{[H₂]} = 1.2 + x\)

### The Completed Table

|         | CO          | H₂O         | CO₂         | H₂          |
|---------|-------------|-------------|-------------|-------------|
| Initial | 7.2         | 5.2         | 7.2         | 1.2         |
| Change  | \(-x\)      | \(-x\)      | \(+x\)      | \(+x\)      |
| Equilibrium | \(7.2 - x\) | \(5.2 - x\) | \(7.2 + x\) | \(1.2 + x\) |

Would you like to explore the equilibrium constant or any further details? Feel free to ask!

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Here are five related questions:
1. How is the equilibrium constant expression for this reaction written?
2. How would you determine the value of \(x\) experimentally?
3. What factors could shift this equilibrium according to Le Chatelier’s Principle?
4. How does temperature influence the equilibrium position of this reaction?
5. Could we calculate the equilibrium constant (Kc) if additional data were provided?

**Tip**: Always ensure that your initial molarities are correct by checking the total moles and volume. This is crucial for accurate equilibrium calculations.

Suppose a 250. mL flask is filled with 1.3 mol of H₂O, 1.8 mol of CO₂, and 0.30 mol of H₂. The reaction CO(g) + H₂O(g) ⇌ CO₂(g) + H₂(g) takes place. Complete the table, listing initial molarity, change in molarity, and equilibrium molarity using x to denote the change in CO.

Learn how to calculate equilibrium molarity for the reaction CO(g) + H₂O(g) ⇌ CO₂(g) + H₂(g). This problem involves initial molarity calculations, stoichiometric changes, and determining equilibrium concentrations using the change variable x. Suitable for undergraduate chemistry students.

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