Acids and Bases in Water

• Acids and bases react when in contact with water by releasing protons (H⁺) or hydroxide (OH⁻) ions.
• These reactions can be described by dissociation equations (splitting into ions) or protolysis equations (proton transfer).
• These reactions are equilibrium reactions that occur quickly.
• Ion concentrations are present at equilibrium, which is referred to as ionic equilibrium.
• Dissociation and protolysis equations:
  • Dissociation: HNO₃ → H⁺(aq) + NO₃⁻(aq),
  • Protolysis: HNO₃ + H₂O → H₃O⁺ + NO₃⁻.
• According to Brönsted's acid-base theory, acids and bases react in pairs, so there is no free "H⁺".
• Dissociation always occurs in water, the proton (H⁺) binds to a water molecule, forming an oxonium/hydronium ion (H₃O⁺).
• The H⁺ concentration is used analogously to the H₃O⁺ concentration.
• The acidic/basic nature is determined by the number of protons/hydroxide ions in an aqueous solution.
• The higher the concentration, the more acidic/basic the solution is.
• The negative decadic logarithm of the hydronium/hydroxide concentration is used to measure acidity/basicity, denoted as the pH/pOH value (pH = potentia hydrogenii, the power of hydrogen).

pH/pOH Equations

• pH = -lg([H₃O⁺] / (1 mol/L))
• pOH = -lg([OH⁻] / (1 mol/L))
• [H₃O⁺] = 10⁻pH mol/L
• [OH⁻] = 10⁻pOH mol/L
• pH and pOH are related because acids and bases only have acidic/basic properties in aqueous solutions, which relies on the autoprotolysis of water.

Autoprotolysis and the Ion Product of Water

• Pure water can conduct electricity slightly because it can form ions.
• This is due to water's amphoteric nature, which means it can react as both a base and an acid.
• The reaction is called autoprotolysis, which is proton transfer between like molecules.
• Autoprotolysis equation: H₂O + H₂O ⇌ H₃O⁺ + OH⁻
• Mass action law (MWG): Kc = [H₃O⁺][OH⁻] / [H₂O]²
• At equilibrium (room temperature), only a few water molecules are protolyzed, so the equilibrium lies on the reactant side.
• The water concentration remains mostly constant and can be included in the equilibrium constant:
• Kw = [H₃O⁺][OH⁻] = 10⁻¹⁴ mol²/L² (ion product of water)
• Since oxonium and hydroxide ions exist in a 1:1 ratio, the equilibrium concentrations are:
  • [H₃O⁺] = [OH⁻] = √Kw = 10⁻⁷ mol/L
• Negative decadic logarithm applied to the ion product of water:
  • pKw = -lg(Kw / (1 mol²/L²)) = -lg([H₃O⁺][OH⁻] / (1 mol²/L²)) = pH + pOH = 14
• In pure water: pH = pOH = 7.
• The ion product of water applies to diluted aqueous solutions because water is in excess.

More on pH

• For aqueous solutions of acids/bases: pH + pOH = 14.
• Solutions with a pH of 7 are neutral because the H₃O⁺ and OH⁻ concentrations are equal.
• Solutions with a pH less than 7 are acidic, and solutions with a pH greater than 7 are basic.
• Indicators can make the pH visible.
• An indicator is an organic dye that changes color depending on the pH value.
• Each indicator has a characteristic color change range and spectrum.
• Neutralization reactions (acid-base titrations) determine the endpoint of neutralization using the color change of the indicator.
• The endpoint is the equivalence point, where the amounts of acids and bases are identical.

Strength of Acids and Bases

• The position of the protolysis equilibrium differs depending on the strength of the acid or base.
• Strong acids and bases differ from weak acids and bases in calculating the [H₃O⁺] or [OH⁻] concentration at equilibrium.
• Strong acids and bases undergo complete protolysis, shifting the equilibrium to the right.
• The [H₃O⁺] or [OH⁻] concentration matches the initial concentration of the acid or base, but also accounts for stoichiometry.
  • Strong acids/bases: [H₃O⁺] = y * [HA]₀, [OH⁻] = y * [B]₀ (y = valency)
• Weak monoprotic acids and bases undergo incomplete protolysis, with the equilibrium on the reactants' side.
• The [H₃O⁺] or [OH⁻] concentration is found using the mass action law (MWG).
  • Equation: HA + H₂O ⇌ H₃O⁺ + A⁻ and B + H₂O ⇌ BH⁺ + OH⁻ Assuming the water concentration is constant, the equilibrium constant Kc is:
• Acid constant: Ks = [H₃O⁺][A⁻] / [HA], Ks = Kc · [H₂O]
• Base constant: KB = [BH⁺][OH⁻] / [B], KB = KC · [H₂O]
• High Ks/KB values indicate strong acids/bases, whereas low values indicate weak acids/bases.
• Acid exponent: pKs = −lg(Ks).
• Base exponent: pKB = -lg(KB).
• In terms of exponents, small pKs/pKB values mean a stronger acid/base and vice versa.
• Because the MWG describes the equilibrium position, equilibrium concentrations must be used; for weak monoprotic acids and bases:
  • Only a few acid/base molecules are protolyzed, so the concentration is at the initial value.
  • From the reaction equation, there is a 1:1 ratio of H₃O⁺ to A⁻ or BH⁺ to OH⁻.
  • Therefore: Ks = [H₃O⁺]² / [HA]₀ and KB = [OH⁻]² / [B]₀
  • Transposed to [H₃O⁺] = √Ks · [HA]₀ and [OH⁻] = √KB · [B]₀

Calculating pH and pOH

• For pH/pOH calculations: pH = ½ * (pKs - lg([HA]₀ / (1 mol/L))) pOH = ½ * (pKB - lg([B]₀ / (1 mol/L)))
• For acids/Bases with a Ks/KB value of < 10⁻⁴ mol/L (or pKs/pKB value > 4) are weak acids/bases.

Polyprotic weak acids and bases

• The valency of an acid/base is the number of protons that can be released/absorbed.
• Polyprotic acids/bases are protolyzed stepwise.
• Each protolysis step has its own acid/base constant, so generally, Ks(I) > Ks(II) > Ks(III) or KB(I) > KB(II) > KB(III)
• For polyprotic weak acids/bases, if the Ks/KB values vary by a factor of 1000, only the acid/base constant of the first protolysis step needs to be considered.
• Otherwise, only a few acid/base molecules are protolyzed in the first step, with smaller amounts in the second.

Degree of Protolysis

• The degree of protolysis a is the ratio of protolyzed acid/base species to the initial concentration of the acid/base expressed as a value between 0 and 1.
• This can be converted to percentage if the value is multiplied by 100%.
• The degree of protolysis:
  • α = [H₃O⁺] / [HA]₀ or α = [OH⁻] / [B]₀
  • The larger the value of α, the more strongly (protolyzed) is the acid/base.