Acids, Bases, and K_w

• Acids, Bases, and K_w are fundamental concepts in chemistry.

Acids and Their Properties

• Acids react with water to form H⁺(aq) ions.
• Described by the equation: HX → H⁺(aq) + X^-(aq).
• H⁺ ions are essentially single protons, they are tiny and charged.
• H⁺ ions do not exist alone in water; they interact with H₂O molecules.
• The simplified representation of H+(aq) or H3O+ is often used.
• Acids are proton donors, according to the Bronsted-Lowry Model.
• Strong acids fully ionize in a solution with all molecules ionised, described by the equation: HX → H+(aq) + X-(aq).
• Weak acids only partially ionize in a solution with a small fraction of molecules ionising, described by the equation: HX ⇌ H+(aq) + X-(aq).
• Concentrated acids have a high number of moles of acid per dm³ of solution.
• Dilute acids have a low number of moles of acid per dm³ of solution.
• It is important note: strong/weak and concentrated/dilute acids are not interchangeable terms.
• Monobasic, dibasic, and tribasic acids refer to the total number of replaceable hydrogen ions.
  • H3PO4 + 3OH- → 3H2O + PO43- (tribasic)
  • H2SO4 + 2OH- → 2H2O + SO42- (dibasic)
  • HCl + OH- → H2O + Cl- (monobasic)
• Terms like monoprotic, diprotic, and triprotic are interchangeable with monobasic, dibasic, and tribasic respectively.

Measuring pH

• pH is measured accurately using a pH meter.
• pH meters convert the electrode potential difference into a pH reading.
• Universal indicators are used to estimate pH values.
• The pH scale was introduced in 1909 by Danish chemist Soren Sorenson.
• It replaced color-based acidity/basicity measurements with numbers.
• Sorenson measured hydrogen ion concentration using an electrochemical cell.
• pH is a negative logarithm to the base 10 of the H⁺(aq) concentration.
• Aqueous hydrogen ion concentrations are converted onto a scale from just below 0 to 14.
• The higher the H⁺(aq) concentration, the more acidic the solution.
• It is calculated by: pH = -log₁₀[H⁺(aq)], where [H⁺] is in mol dm^-3.
• Each pH value is 10 times the value below it.
  • pH 5 is 10 times more acidic than pH 6
  • pH 4 is 10 times less acidic than pH 3

Calculating pH

• To calculate the pH of a strong acid with a known concentration: pH = -log₁₀[H⁺].
• The 'log' button has a default base of log₁₀ on calculators.
• If the pH is given, the concentration can be calculated using: [H⁺] = 10^-pH.
• For diprotic acids like H₂SO₄ remember to account for the fact they release two H⁺ ions per molecule.
  • For example: H₂SO₄(aq) → 2H⁺(aq) + SO₄²⁻(aq)

Ionic Product of Water (K_w)

• Water is amphoteric, acting as both an acid and a base.
• Hydronium and hydroxide ions are formed in the dissociation of water, its a very weak process
  • 2H₂O(l) ⇌ H₃O⁺(aq) + OH^-(aq)
• The net effect is that an equilibrium is established .
• [H₂O(l)] is virtually constant, which will bring about a new equilibrium constant, K_w.
• K_w = [H⁺(aq)][OH⁻(aq)].
• K_w = 1.00 × 10^-14 mol² dm^-6.
• In neutral water, [H⁺(aq)] = [OH⁻(aq)].
• For every hydrogen ion formed, there is a hydroxide ion formed as well.
• The dissociation of water is an endothermic reaction
• As a result K_w increases with temperature.
• [H⁺] increases and pH decreases.
• Water remains neutral because [H⁺] = [OH^-].
• Different at different temperatures.
• pH is 6.52 at 60°C and 7.27 at 10°C.
• It is still neutral as concentrations of H+ and OH- ions are still equal.
• For whole number pH values, the indices for [H⁺(aq)] and [OH^-(aq)] add up to -14.
  • acid rain solution with a pH of 3 gives [H⁺(aq)] = 10^-3 mol dm^-3, and [OH⁻⁽aq)] = 10^-11 mol dm^-3.

Calculating pH using K_w

• A solution is acidic if [H⁺(aq)] > [OH⁻(aq)].
• A solution is neutral if [H⁺(aq)] = [OH⁻(aq)].
• A solution is alkaline if [H⁺(aq)] < [OH⁻(aq)].
• Acidic solutions still have OH⁻(aq) ions, but there are more H⁺(aq) ions
• K_w value controls the concentrations of H⁺(aq) and OH⁻(aq).