pH of strong acid and alkali | Part-2

Calculation of pH of Strong Acid and Alkali of Particular Strength

The pH scale is a numeric scale used to indicate the concentration of hydrogen ions (H⁺) in an aqueous solution. The scale spans from 0 to 14 and is used to identify if a solution is acidic, neutral, or basic in nature.

pH is defined as:

pH = -log₁₀[H⁺]

Here, [H⁺] represents the concentration of hydrogen ions (in moles per liter) present in the solution.

What are Strong Acids and Strong Alkalis?

Strong acids like HCl, HNO₃, and H₂SO₄ fully ionize in water, releasing one or more H⁺ ions per molecule. Similarly, strong alkalis such as NaOH and KOH completely dissociate in water, providing OH⁻ ions.

This complete dissociation allows for direct calculation of [H⁺] or [OH⁻] from the given molarity.

Calculating pH of a Strong Acid

Because strong acids fully ionize in solution, the hydrogen ion concentration directly corresponds to the acid’s molarity.

HCl → H⁺ + Cl⁻

For an HCl solution with a molarity of 0.01 M:

[H⁺] = 0.01 M  
pH = -log₁₀(0.01) = 2

For 0.0001 M HNO₃:

pH = -log₁₀(0.0001) = 4

Calculating pH of a Strong Alkali

Strong bases fully ionize, so the concentration of OH⁻ is the same as the base's molarity. To calculate pH, first determine pOH:

NaOH → Na⁺ + OH⁻

Then use:

pOH = -log₁₀[OH⁻]  
pH = 14 - pOH

Example: For 0.001 M NaOH

pOH = -log₁₀(0.001) = 3  
pH = 14 - 3 = 11

For 0.00001 M KOH:

pOH = 5  
pH = 14 - 5 = 9

Important Points for JEE Aspirants

• Strong acids and bases dissociate 100% in water.
• Use pH = –log[H⁺] for acids and pH = 14 – pOH for alkalis.
• Always check units: concentrations must be in mol/L.
• Familiarity with logarithms is essential: e.g., log(10⁻³) = –3

Knowing how to calculate the pH of strong acid and alkali solutions based on their molarity is a fundamental skill in chemistry. These calculations are straightforward due to complete dissociation, and mastering them gives you a strong foundation for topics like titration, salt hydrolysis, and buffer solutions.