Introduction to pH and Strong Bases

The pH scale is a logarithmic scale used to specify the acidity or basicity of a solution. Strong bases, such as sodium hydroxide (NaOH) or potassium hydroxide (KOH), completely dissociate in water, making their pH calculations straightforward. In this article, we will explore three practical examples of calculating the pH of strong bases.

Example 1: Calculating the pH of a 0.1 M NaOH Solution

In a laboratory setting, a chemist prepares a sodium hydroxide solution with a concentration of 0.1 M to conduct a titration experiment. To understand the strength of the base, the chemist needs to calculate the pH of the solution.

To calculate the pH, we first determine the pOH using the formula:

pOH = -log[OH⁻]
Since NaOH is a strong base, it fully dissociates in water:
NaOH → Na⁺ + OH⁻
Thus, [OH⁻] = 0.1 M.

Calculating pOH:
pOH = -log(0.1) = 1

To find pH, we can use the relationship between pH and pOH:

pH + pOH = 14

Therefore:
pH = 14 - pOH
pH = 14 - 1 = 13

The pH of a 0.1 M NaOH solution is 13, indicating a strongly basic solution.

Notes:

• The concentration of the strong base directly influences the pH.
• For every tenfold increase in concentration, the pH increases by 1 unit.

Example 2: pH Calculation for a 0.5 M KOH Solution

An industrial facility uses potassium hydroxide (KOH) for various chemical processes. They prepare a 0.5 M KOH solution and want to evaluate its pH to ensure safety and compliance with regulations.

Similar to the previous example, we can calculate the pOH first:

pOH = -log[OH⁻]
Since KOH is a strong base, it fully dissociates in water:
KOH → K⁺ + OH⁻
Thus, [OH⁻] = 0.5 M.

Calculating pOH:
pOH = -log(0.5) ≈ 0.30

Using the relationship between pH and pOH:

pH + pOH = 14

Therefore:
pH = 14 - pOH
pH = 14 - 0.30 ≈ 13.70

The pH of this 0.5 M KOH solution is approximately 13.70, confirming its strong basicity.

Notes:

• KOH is often used in industrial applications, so knowing the pH is crucial for safety.
• The high pH indicates a significant concentration of hydroxide ions, which can be corrosive.

Example 3: Determining pH of a Dilute 0.01 M Ba(OH)₂ Solution

In a classroom experiment, students prepare a dilute barium hydroxide (Ba(OH)₂) solution with a concentration of 0.01 M. The instructor asks them to calculate the pH of this solution to understand the properties of strong bases.

Barium hydroxide is also a strong base and dissociates completely in water:
Ba(OH)₂ → Ba²⁺ + 2OH⁻
Thus, for every mole of Ba(OH)₂, there are two moles of OH⁻ produced, leading to:
[OH⁻] = 2 × 0.01 M = 0.02 M.

Calculating pOH:
pOH = -log(0.02) ≈ 1.70

Using the relationship between pH and pOH:

pH + pOH = 14

Therefore:
pH = 14 - pOH
pH = 14 - 1.70 ≈ 12.30

The pH of a 0.01 M Ba(OH)₂ solution is approximately 12.30, showing it is still a basic solution, albeit less so than the previous examples.

Notes:

• Barium hydroxide is often used in laboratories and industrial settings, making pH calculations essential.
• The stoichiometry of dissociation must be considered when calculating hydroxide ion concentration.