Strong Acid/Base; Weak Acid/Base

Calculate ion concentrations in an aqueous solution of a strong acid

The concentration of hydronium ion for  a strong acid, is the same as the concentration of the acid.

An example of a monoprotic strong acid-  a 0.50 M solution of nitric acid:

HNO3(aq) + H2O(l) àH3O+(aq) + NO3–(aq).

The final concentrations of ions can be determined for strong acids easily since there is 100% dissociation.

HNO3(aq)    +H2O(l)          àH3O+(aq)      +NO3–(aq).
Initial Concentration0.50 M–00
Change-0.50 M–+ 0.50 M+ 0.50 M
Final Concentrations0 M–+ 0.50 M+ 0.50 M

There is one diprotic strong acid, however, which is sulfuric acid (H2SO4).  Only the first ionization is strong in sulfuric acid, and similarly to nitric acid, the ions formed from sulfuric acid will be the same as the concentration of the sulfuric acid. 

Let’s repeat the example, but with sulfuric acid:

For a sample of the only diprotic strong acid-  a 0.50 M solution of sulfuric acid:

H2SO4(aq) + H2O(l)  H3O+(aq) + HSO4–(aq).

The final concentrations of ions can be determined for strong acids easily since there is 100% dissociation.

There is one diprotic strong acid, however, which is sulfuric acid (H2SO4).  Only the first ionization is strong in sulfuric acid, and similarly to nitric acid, the ions formed from sulfuric acid will be the same as the concentration of the sulfuric acid. The HSO4– ion is also an acid, but it is weak and does not dissociate 100% or even to a large extent, and therefore only negligibly contributes to the concentration of ions. 

H2SO4(aq)    +H2O(l)          H3O+(aq)      +HSO4–(aq).
Initial Concentration0.50 M–00
Change-0.50 M–+ 0.50 M+ 0.50 M
Final Concentrations0 M–+ 0.50 M+ 0.50 M

Calculate the pH of an aqueous solution of a strong base

  • Completely dissociate in water
  • Examples :Hydroxides of Group I elements, Alkaline earth hydroxides (Group II)

Calculate the pH of a 5.0×10–2 M NaOH solution

The major species in this solution are Na+, OH–, and H2O

Although autoionization of water also produces OH– ions, the pH will be dominated by the OH– ions from the dissolved NaOHConcentration of H+ can be calculated from Kw (dissociation constant for water which is 1 X10-14 @ 25oC)

pH = 12.70

Weak Acids and Weak Bases – Determine the value of Ka from the pH of a weak acidic solution

We just learned that strong acids and bases completely dissociate.  Weak acids and bases clearly do not completely dissociate, and often only dissociate by 5% or less.  The calculation to find the hydronium ion or hydroxide ion is different, and makes use of the equilibrium constants, Ka for acids or Kb for bases. Strength of an acid is directly related to how much the products are favored over the reactants

HA(aq) + H2O(l)  H3O+(aq) + A–(aq)•If Ka > 1, products are favored, strong acid•If Ka < 1, reactants are favored, weak acid
Here are Tips on how to solve weak acid equilibrium problems: 

  1. List the major species in the solution
  2. Choose the species that can produce H+
  3. Write balanced equations for the reactions producing H+
  4. Using the values of the equilibrium constants for the reactions, decide which equilibrium will dominate in producing H+
  5. Write the equilibrium expression for the dominant equilibrium
  6. List the initial concentrations of the species participating in the dominant equilibrium
  7. Define the change needed to achieve equilibrium
  8. Define x
  9. Write the equilibrium concentrations in terms of x

Example:   

Propanoic acid, CH3CH2COOH, is a weak monoprotic acid that is used to inhibit mould formation in bread. A student prepared a 0.10M solution of propanoic acid and found that the pH was 2.96. What is the Ka for propanoic acid?

Answer: 

Calculate the pH of a weak acid solution

The Ka can be used to calculate the equilibrium concentrations and the pH in a solution of weak acidWrite the chemical equation and set up an ICE table.  To make the math simple, the 500 rule can be applied and this usually applies to dilute weak acids. –Let x represent the change in concentration–If [HA]/Ka > 500, x can be ignored–If [HA]/Ka < 500, use quadratic formula
Once the H3O+ concentration is calculated through an ICE table, the pH can then be calculated.  Example:

Formic acid, HCOOH, is present in the sting of certain ants. What is the pH of a 0.25 mol/L solution of formic acid? (Ka = 1.8 X 10-4)

Answer: 

 Calculate the pH of a weak basic solution

•Weak bases do not completely dissociate in water so [OH–] depends on the initial concentration of the base and the concentration of the conjugate acid

  1. List the major species in the solution
  2. Choose the species that can produce H+
  3. Write balanced equations for the reactions producing H+
  4. Using the values of the equilibrium constants for the reactions, decide which equilibrium will dominate in producing H+
  5. Write the equilibrium expression for the dominant equilibrium
  6. List the initial concentrations of the species participating in the dominant equilibrium
  7. Define the change needed to achieve equilibrium(Define x)
  8. Write the equilibrium concentrations in terms of x

Strength of the base is directly related to how much the products are favored over the reactants

•If Kb > 1, products favored, strong base•If Kb < 1, reactants favored, weak base

Example: 

The characteristic taste of tonic water is due to the addition of quinine. Quinine is a naturally occurring compound that is also used to treat malaria. The base dissociation constant, Kb for quinine is 3.3 × 10−6. Calculate [OH−] and the pH of a 1.7 × 10−3 M solution of quinine.

Answer: