Acid and bases

• Acid-base equilibria involve the transfer of protons between substances
• Substances can be classified as acids or bases depending on their interaction with protons
• A Brønsted-Lowry acid is a proton donor, for example, Ammonium ions (NH4+)
• A Brønsted-Lowry base is a proton acceptor, for example, Hydroxide ions (OH-)
• Acid strength doesn’t refer to the concentration of a solution
• A strong acid completely dissociates to ions when in solution with pH 3-5
• A weak acid only slightly dissociates when in solution with pH 0-1
• The same definitions apply to strong and weak bases, with strong bases having pH 12-14 and weak bases pH 9-11
• pH is a measure of acidity and alkalinity on a logarithmic scale from 0 to 14
• 0 is an acidic solution with a high concentration of H+ ions, whereas 14 is a basic solution with a low concentration of H+ ions
• The concentration of H+ ions can be determined if the pH is known
• Water slightly dissociates to ions with its equilibrium constant, Kw, having a constant value of 1 x 10-14 at 25°C
• The forward reaction in the equilibrium of water is endothermic and favored with increased temperature, resulting in more H+ ions produced and the water becoming more acidic
• Weak acids and bases only slightly dissociate in solution to form an equilibrium mixture
• The reaction has an equilibrium dissociation constant, Ka, which can be found using pKa
• Different methods are used depending on the reaction and concentrations:

  • HA in excess: Use [HA] and [A-] along with Ka to find [H+], then pH
  • A- in excess: Use Kw to find [H+], then pH
  • HA = A-: pKa is equal to pH, therefore find pKa
• A pH titration curve shows how pH changes during an acid-base reaction, reaching a neutralization point identified as a large vertical section through the equivalence point
• For different acid-base combinations, the neutralization point varies:

  • Strong Acid + Strong Base = pH 7
  • Strong Acid + Weak Base = < pH 7 (more acidic)
  • Weak Acid + Strong Base = > pH 7 (more basic)
  • Weak Acid + Weak Base = normally pH 7 but hard to determine
• Specific indicators like methyl orange and phenolphthalein are used for specific reactions to indicate a pH change within a certain range
• Methyl Orange:

  • Used for reactions with a more acidic neutralization point
  • Orange in acids and turns yellow at the neutralization point

  Phenolphthalein:

  • Used for reactions with a more basic neutralization point
  • Pink in alkalis and turns colorless at the neutralization point
• Buffer solutions resist changes in pH when small volumes of acid or base are added
• Acidic buffer solutions contain a weak acid and the salt of that weak acid, while basic buffer solutions contain a weak base and the salt of that weak base
• Buffer calculations involve acid-base calculations with two types: Acid + Base and Acid + Salt
• The pH of a buffer solution changes in the order of 0.1 or 0.01 units when small volumes of acid or base are added
• Adding small amounts of acid increases the acidity of the solution, while adding small amounts of base decreases the acidity
• Buffer solutions are important in nature to keep systems regulated, especially in living organisms where enzymes or reactions require a specific pH
• Bronsted-Lowry acid: substance that can donate a proton
• Bronsted-Lowry base: substance that can accept a proton
• In the reaction HCl (g) + H2O (l) → H3O+ (aq) + Cl- (aq), HCl is the acid and H2O is the base
• pH = - log [H+] where [H+] is the concentration of hydrogen ions in the solution
• Strong acids completely dissociate in solution
• pH of strong acids can be calculated using pH = - log [acid concentration]
• To find [H+], use [H+] = 1 x 10^(-pH)
• In pure water, Kw = [H+] x [OH-] = 1 x 10^(-14) mol^2 dm^(-6)
• At 25°C, [H+] = [OH-] = 1 x 10^(-7) mol dm^(-3), so pH of pure water is 7
• Increasing temperature in water dissociation equilibrium shifts to the right, increasing [H+] and lowering pH
• For weak acids, they slightly dissociate in water forming an equilibrium mixture
• Ka is the acid dissociation constant, larger Ka means a stronger acid
• To calculate pH of weak acids, simplify the Ka expression and solve for [H+]
• For partially neutralized strong acid and base reactions, calculate moles of acid and base, determine excess ions, and find new concentration of H+ or OH- ions to calculate pH