Acid-Base Reactions (College Board AP Chemistry)

Strong Acid - Strong Base

• When a strong acid and a strong base are mixed, they react quantitatively in a reaction represented by the equation:
  

H₃O⁺ (aq) + OH^– (aq) → 2H₂O (l)

OR

H⁺ (aq) + OH^– (aq) → H₂O (l)

 

• An example is HCl reacting with KOH
  • HCl + KOH → KCl + H₂O
• The pH of the resulting solution may be determined from the concentration of excess reagent

Weak base - Strong acid

• When a weak base and a strong acid are mixed, they will react quantitatively in a reaction represented by the equation:

B (aq) + H₃O⁺ (aq) ⇌ HB⁺ (aq) + H₂O (l)

• For example if NH₃ and HCl react
  • NH₃ (aq) + HCl (aq) → NH₄Cl (aq) 
  • NH₃ (aq) + H₃O⁺ (aq) NH₄⁺ (aq) + H₂O

• If the weak base is in excess, then a buffer solution is formed, and the pH can be determined from the Henderson–Hasselbalch equation equation:

pOH = pK_b + log 

• This is because a salt and acid will be in the remaining solution

• If the strong acid is in excess, then the pH can be determined from the moles of excess hydronium ion and the total volume of solution

pH = -log10[H₃O⁺] 

OR

pH = -log10[H⁺] 

• If they are equimolar, then the (slightly acidic) pH can be determined from the equilibrium represented by the equation:

HB⁺ (aq) + H₂O (l) ⇌ B (aq) + H₃O⁺ (aq)

• Acid and base reactions will favour the side with the weakest acids and bases

NH₄⁺ (aq) + H₂O NH₃ (aq) + H₃O+ (aq)

• NH₄⁺ is a weak acid, so will have a K_a value < 1 
• Therefore the equilibrium will shift to the left hand side

Strong Base - Weak Acid

• When a weak acid and a strong base are mixed, they react quantitatively in a reaction represented by the equation:
  

HA (aq) + OH^– (aq) ⇌ A^– (aq) + H₂O (l)

• If the weak acid is in excess, then a buffer solution is formed, and the pH can be determined from the Henderson-Hasselbalch equation

pH = pK_a + log

• If the strong base is in excess, then the pH can be determined from the moles of excess hydroxide ion and the total volume of solution

[H⁺] = 

• If they are equimolar, then the (slightly basic) pH can be determined from the equilibrium represented by the equation:

A^– (aq) + H₂O (l) ⇌ HA (aq) + OH^– (aq)

Weak Base - Weak Acid

• When a weak acid and a weak base are mixed, they will react to an equilibrium state whose reaction may be represented by the equation:
  

HA (aq) + B (aq) ⇌ A^– (aq) + HB⁺ (aq)

• In an acid base reaction, the equilibrium will favour the side with the weaker acid and weaker base
• If the equilibrium constant for the reaction is greater than one (K >1) the equilibrium will favour the formation of products
  • This means the weaker acid and weaker base are on the product side
• If the equilibrium constant for the reaction is less than one (K <1) the equilibrium will favour the formation of reactants
  • This means the weaker acid and weaker base are on the reactant side
• The reaction between HSO₄^- and CO₃^2- is an example of a weak acid - weak base reaction
• The products are weaker acids and bases than the reactants, so the equilibrium will favour the formation of the products

Weak acid - weak base reaction

The reaction favours the right hand side

• If the solutions are equimolar, then we need to consider the K_a and K_b values to determine where the final solution will be acidic, neutral or basic
• This can be shown by the reaction between HF and NH₃

Reaction between HF and NH₃

The reaction favours the right hand side

• As we can see the K_a and K_b values are greater for the reactants than the products
  • This means they are stronger acids and bases
  • Therefore the equilibrium will favour the products and K will be greater than 1
• The K_a value for NH₄⁺ is greater than the K_b value for F^–, therefore it is better at forming  H₃O⁺ ions than the F^– ion is at forming OH^– ions
• Therefore the solution will be acidic