Strong & Weak Acids (DP IB Chemistry)

Strong & Weak Acids

Strong acids

• A strong acid is an acid that dissociates almost completely in aqueous solutions

  • Examples include HCl (hydrochloric acid), HNO₃ (nitric acid) and H₂SO₄ (sulfuric acid)

  • The position of the equilibrium is so far over to the right that you can represent the reaction as an irreversible reaction

Diagram to show the dissociation of a strong acid

The diagram shows the complete dissociation of a strong acid in aqueous solution

• The solution formed is highly acidic due to the high concentration of the H⁺/H₃O⁺ ions

• Since the pH depends on the concentration of H⁺/H₃O⁺ ions, the pH can be calculated if the concentration of the strong acid is known

  • pH = -log₁₀[H⁺ (aq)]

  • [H+ (aq)] = concentration of H⁺ / H₃O⁺ ions

  • pH is the negative log of the concentration of H⁺ / H₃O⁺ ions and can be calculated if the concentration of the strong acid is known using the stoichiometry of the reaction

Weak acids

• A weak acid is an acid that partially (or incompletely) dissociates in aqueous solutions

  • E.g. most organic acids (ethanoic acid), HCN (hydrocyanic acid), H₂S (hydrogen sulfide) and H₂CO₃ (carbonic acid)

  • The position of the equilibrium is more towards the left and an equilibrium is established

 Diagram to show the dissociation of a weak acid

The diagram shows the partial dissociation of a weak acid in aqueous solution

• The solution is less acidic due to the lower concentration of H⁺ / H₃O⁺ ions

Acid & Equilibrium Position Table

• The strength of a Brønsted-Lowry acid depends on the ease with which it dissociates to release H⁺ ions

  • This depends upon the strength of the bond that has to be broken to release H+

  • For example, for hydrogen halides, the size of the halogen atom increases in size going down Group 17 which increases the length of the H–X bond

  • As longer bonds are weaker they need less energy to break 

  • The acid strength of the hydrogen halides increases down Group 17

    • HF < HCl < HBr < HI

Strong bases

• A strong base is a base that dissociates almost completely in aqueous solutions

  • E.g. group 1 metal hydroxides such as NaOH (sodium hydroxide)

  • The position of the equilibrium is so far over to the right that you can represent the reaction as an irreversible reaction

 Diagram to show the dissociation of a strong base

The diagram shows the complete dissociation of a strong base in aqueous solution

• The solution formed is highly basic due to the high concentration of the OH^– ions

Weak bases

• A weak base is a base that partially (or incompletely) dissociates in aqueous solutions

  • NH₃ (ammonia), amines and some hydroxides of transition metals

  • The position of the equilibrium is more to the left and an equilibrium is established

 Diagram to show the dissociation of a weak base

The diagram shows the partial dissociation of a weak base in aqueous solution

•  The solution is less basic due to the lower concentration of OH^– ions

Base & Equilibrium Position Table

Strength of conjugate acids and bases

• The conjugate base of HCl is the chloride ion, Cl^–,

• However, since the reverse reaction is virtually non-existent the chloride ion must be a very weak conjugate base

HCl (g) → H⁺ (aq)  +   Cl^– (aq)

acid                          conjugate base

• In general, strong acids produce weak conjugate bases and weak acids produce strong conjugate bases

• A strong base is also fully ionised and is a good proton acceptor

• For example, the hydroxide ion is a strong base and readily accepts protons:

   OH^– (aq) +  H⁺ (aq)  ⇌  H₂O (l)

• The conjugate acid of the hydroxide ion is water, which is a weak conjugate acid

• In general strong bases produce weak conjugate acids

How to distinguish between strong and weak acid

• Strong and weak acids can be distinguished from each other by their:

  • pH value (using a pH meter or universal indicator)

  • Electrical conductivity

  • Reactivity

pH value

• An acid dissociates into H⁺ in solution according to

HA → H⁺ + A^-

pH value of a Strong Acid & Weak Acid Table

• The stronger the acid, the greater the concentration of H⁺ and therefore the lower the pH

Electrical conductivity

• Since a stronger acid has a higher concentration of H⁺ it conducts electricity better

• Stronger acids therefore have a greater electrical conductivity

• The electrical conductivity can be determined by using a conductivity meter

• Like the pH meter, the conductivity meter is connected to an electrode

• The conductivity of the solution can be read off the meter

Diagram to show how to measure the electrical conductivity of an acid

A digital conductivity meter measures the electrical conductivity of a solution using an electrode

Reactivity

• Strong and weak acids of the same concentrations react differently with reactive metals

• This is because the concentration of H⁺ is greater in strong acids compared to weak acids

• The greater H⁺ concentration means that more H₂ gas is produced in a shorter time

Diagram to show how a strong acid reacts with magnesium

The diagram shows the reaction of 0.1 mol dm^-3 of a strong acid (HCl) with Mg. The reaction produces a lot of bubbles and hydrogen gas due to the high concentration of H⁺ present in solution

Diagram to show how a weak acid reacts with magnesium

The diagram shows the reaction of 0.1 mol dm^-3 of a weak acid (CH₃COOH) with Mg. The reaction produces fewer bubbles of hydrogen gas due to the lower concentration of H⁺ present in solution

• Similar observations would be made in the reaction between strong and weak acids with carbonates and hydrogencarbonates, although the gas given off this time is carbon dioxide

• With oxides and hydroxides, there may not be a lot of visible changes although it is likely that they would dissolve faster in a strong acid than in a weak acid

• These reactions are also likely to produce larger enthalpy changes which could be seen in higher temperature rises