a' semi normal solution' of sodium acetate in water has [H+]

To understand the concept of a "semi-normal solution" of sodium acetate in water and its relation to the concentration of hydrogen ions, or [H+], we first need to clarify a few key points about sodium acetate and its behavior in solution.

What is Sodium Acetate?

Sodium acetate (NaC2H3O2) is the sodium salt of acetic acid. When dissolved in water, it dissociates into sodium ions (Na+) and acetate ions (C2H3O2-). The acetate ion can interact with water, leading to some interesting chemistry.

Understanding Semi-Normal Solutions

A semi-normal solution typically refers to a solution that is half the concentration of a normal solution. In the context of sodium acetate, if we consider a semi-normal solution to be 0.5 N, this means that the concentration of the acetate ions is 0.5 equivalents per liter. Since sodium acetate dissociates completely, this also means that the concentration of sodium ions will be 0.5 M, and the concentration of acetate ions will also be 0.5 M.

Acetate Ion and pH

The acetate ion can act as a weak base. It can accept a proton (H+) from water, leading to the formation of acetic acid (CH3COOH) and hydroxide ions (OH-). This reaction can be represented as:

• C2H3O2- + H2O ⇌ CH3COOH + OH-

As a result of this equilibrium, the presence of acetate ions in solution will affect the pH. The generation of hydroxide ions will make the solution slightly basic, which means that the concentration of hydrogen ions, [H+], will be lower than that in pure water.

Calculating [H+] in a Semi-Normal Sodium Acetate Solution

To find the [H+] in a semi-normal sodium acetate solution, we can use the relationship between pH and pOH. Since the solution is basic, we can calculate the pOH first:

• pOH = -log[OH-]

Assuming that the acetate ion concentration is 0.5 M, we can use the Kb (base dissociation constant) for acetate to find the concentration of hydroxide ions. The Kb for acetate is approximately 5.6 x 10^-10. Using the formula for Kb:

• Kb = [CH3COOH][OH-] / [C2H3O2-]

Assuming x is the concentration of OH- produced, we can set up the equation:

• 5.6 x 10^-10 = (x)(x) / (0.5 - x) ≈ (x^2) / 0.5

Solving for x gives us the concentration of OH-. Once we have [OH-], we can find [H+] using the relationship:

• [H+] = 1.0 x 10^-14 / [OH-]

Example Calculation

If we solve for x, we find that [OH-] is approximately 1.06 x 10^-5 M. Using this value:

• [H+] = 1.0 x 10^-14 / 1.06 x 10^-5 ≈ 9.43 x 10^-10 M

This means that in a semi-normal solution of sodium acetate, the concentration of hydrogen ions is significantly lower than in pure water, confirming that the solution is basic.

Final Thoughts

In summary, a semi-normal solution of sodium acetate in water leads to a basic environment due to the presence of acetate ions, which can accept protons from water. This results in a lower concentration of hydrogen ions compared to pure water, demonstrating the interplay between weak acids and bases in solution. Understanding these concepts is crucial for grasping the behavior of buffers and pH in various chemical contexts.