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Calculating the molar mass of a compound is the first step in many chemistry problems. Here, we focus on chloroacetic acid, represented as \(\mathrm{ClCH}_{2}\mathrm{CO}_{2}\mathrm{H}\). To find its molar mass, we sum the atomic masses of each constituent atom:

• Chlorine (Cl): 35.45 g/mol
• Hydrogen (H): 2 atoms at 1.01 g/mol each
• Oxygen (O): 2 atoms at 16.00 g/mol each
• Carbon (C): 12.01 g/mol

Adding these together gives us a total molar mass of 94.49 g/mol.
This value allows us to convert between grams and moles, a critical step in many chemical equilibrium and concentration calculations.

In chemistry, when a weak acid like chloroacetic acid dissolves in water, it only partially dissociates. This dissociation is represented by the equilibrium expression:
\[\mathrm{ClCH}_{2}\mathrm{CO}_{2}\mathrm{H} \rightleftharpoons \mathrm{ClCH}_{2}\mathrm{CO}_{2}^{-} + \mathrm{H}^{+}\]
To describe this equilibrium mathematically, we use the acid dissociation constant, \(K_a\), given by:
\[K_a = \frac{[\mathrm{ClCH}_2\mathrm{CO}_2^{-}][\mathrm{H}^+]}{[\mathrm{ClCH}_2\mathrm{CO}_2\mathrm{H}]}\]
This equation shows the balance between the dissociated ions and the undissociated acid. It's essential to understand how these concentrations relate to reactants and products in weak acid reactions.
The small \(K_a\) value indicates that the acid is indeed weak, meaning it does not completely ionize in solution.

pH is a measure of the acidity or basicity of a solution. For our weak acid, after dissolution and establishing equilibrium, we calculate \([\mathrm{H}^{+}]\) concentration.
This allows us to compute the solution's pH using the formula:
\[\mathrm{pH} = -\log_{10}[\mathrm{H}^+]\]
By substituting the calculated \([\mathrm{H}^+]\) value of 0.00335 M, we find:
\[\mathrm{pH} = -\log_{10}(0.00335) \approx 2.47\]
This value tells us that the solution is acidic, as expected from the dissociation of chloroacetic acid, even though it is considered a moderately weak acid.

The dissociation constant \(K_a\) is significant in understanding weak acids like chloroacetic acid. It quantifies the degree of ionization and provides insights into the acid's strength.
The given \(K_a\) for chloroacetic acid is \(1.40 \times 10^{-3}\), indicating partial dissociation in water. This small value shows that not all of the acid molecules give away their protons (\(\mathrm{H}^+\)) to form the ions
\([\mathrm{ClCH}_2\mathrm{CO}_2^{-}]\) and \([\mathrm{H}^+]\).

• A larger \(K_a\) would mean a stronger acid (more dissociation).
• A smaller \(K_a\) suggests a weaker acid (less dissociation).

Understanding \(K_a\) helps predict how much the weak acid will affect the pH of a solution.