Acids and Bases

• The Bronstead-Lowry theory states that an acid is a proton donor; hydrochloric acid , for example, can donate H⁺ ions to water forming the oxonium ion:
  HCl_(aq) + H₂O_(l) H₃O⁺_(aq) + Cl^-_(aq)
  
• In any acid-base reaction, a conjugate base and conjugate acid is formed:
• Hydrochloric acid is a strong acid and so the position of equilibrium is far to the right; the HCl is fully dissociated into its ions when added to water; it is a strong acid.
• The strength of an acid refers to the power with which it donates H⁺ ions to an aqueous solution; strong acids are powerful H⁺ donors (the donation of H⁺ ions is almost complete). Conversely, weak acids are less powerful at donating H⁺ ions, and the reaction with water forming oxonium ions is incomplete.
• The Bronstead-Lowry theory defines a base as a proton acceptor. Ammonia accepts H⁺ ions when dissolved in water:
  NH_3(aq) + H₂O_(l) NH₄⁺_(aq) + OH^-_(aq)
  
• In this reversible reaction, the position of equilibrium is to the left; ammonia is a weak base.

Water acting as an acid and a base

• As can be seen from the previous reactions, water can act as both an acid and a base; it is amphoteric.
• The following reaction can occur between water molecules:
  H₂O_(l) + H₂O_(l) H₃O⁺_(aq) + OH^-_(aq)
  
  Or more simply:
  
  H₂O_(l) H⁺_(aq) + OH^-_(aq)
  
• The position of equilibrium is far to the left in the above reaction; we know that this is the case, as pure water does not conduct electricity; if the position of equilibrium was to the right, then the ions would allow it to conduct electricity.
• As it is an equilibrium reaction, an equilibrium constant can be established:
• As the position of equilibrium is so far to the left, [H₂O] is virtually constant, it can therefore be brought next to the K value:
• This equilibrium constant is known as the dissociation of water constant, K_w.
• At 298K, this constant has a value of 10^-14 mol² dm^-6.
• When an acid or base is added to water, the equilibrium between H⁺ and OH^- is disrupted. The concentrations of OH^- or H⁺ ions in the solution can be calculated using the dissociation constant. For example, in a 0.1 mol dm^-3 solution of HCl, the concentration of OH^- ions can be calculated:
• In a 0.02 mol dm^-3 solution of NaOH, the concentration of H⁺ ions can be calculated:
• Rather than using very small values, the negative logarithm of the concentration of H⁺ ions is taken; this gives us the pH of the solution:
• Therefore, the pH of the 0.1 mol dm^-3 solution of HCl equates to:
• The pH of the 0.02 mol dm^-3 solution of NaOH is equal to:
• The pH scale allows the strengths of acids and bases to be compared; a strong acid has a low pH; as the pH increases, the solutions become less acidic, then basic and finally become strongly basic.

Calculating pH values

Strong acids

• This is a reasonable straight forward calculation; the reaction with water goes to completion, therefore the number of moles of H₃O⁺ ions is equal to the number of moles of acid in the solution.
• Therefore, a 1 mol dm^-3 solution of HX (acid) solution will contain 1 mol of H⁺ ions, therefore its pH will be equal to:
• It is important to note the difference with a diprotic acid, such as H₂SO₄. When added to an aqueous solution, sulphuric acid dissociates as follows:
  H₂SO_4(aq) 2H⁺ _(aq) + SO₄^2-_(aq)
  
• Therefore, a 0.01 mol dm^-3 solution of H₂SO₄ will actually dissociate to form 0.02 mols of H⁺ ions.

Weak acids

• Weak acids do not fully dissociate in a solution, and so the calculation is more complex.
• Imagine that a HA represents a weak acid; the following equation would illustrates what happens when the acid is added to water:
  HA_(aq) + H₂O_(l) H₃O⁺_(aq) + A^- _(aq)
  
• Once again, an equilibrium constant can be established for this equilibrium reaction:
• Water can be neglected from the equilibrium constant as it is acting as the solvent.
• The equilibrium constant is called the acidity (dissociation) constant and is given the symbol K_a.
• The table below shows the K_a for some common weak acids:
• Two assumptions can be made when calculating the [H⁺]:
  1. [H⁺] = [A^-], this may seem obvious as both are formed when HA dissociates; however, water itself is also a source of H⁺ ions. The amount of H⁺ ions formed by water is insignificant, thus it is neglected.
  2. The amount of HA at equilibrium is equal to the amount of HA put into the solution; in other words, the fraction of HA that has lost an H⁺ ion can be ignored. This can be done, as we are dealing with weak acids and so this fraction is only very small.
• With these assumptions, it is possible to calculate the pH of a weak acid. For example, the pH of a 0.01 mol dm^-3 solution of ethanoic acid can be worked out as follows:
  

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