Acid- Bases Equilibria

• Acids are proton (H + ) donors, bases are proton (H + ) acceptors, and conjugate acids are formed by the addition of H + to a base, while conjugate bases are formed by the loss of H + from an acid.
• A strong acid is fully dissociated in water, a weak acid is only partially dissociated in water, and when a weak acid dissociates, an equilibrium is set up so it has an equilibrium constant.
• The higher the K a value, the stronger the acid, and pK a is the negative log of K a.
• To calculate the pH of any solution, the concentration of H+ ions must be found, and for a strong acid, HX, [H + ] = [HX], while for a strong dibasic acid, H 2 A, that fully dissociates, [H + ] = [2HA].
• To calculate hydrogen ion concentration from pH, use [H + ] = 10 –pH.
• For a weak acid, HA, since K a = [H + ][A – ] and [H + ] = [A – ], [H + ] = √Ka × [HA].
• In pure water, a small fraction of molecules dissociates into H + and OH – ions, forming an equilibrium.
• The amount of water that dissociates is tiny, so its concentration can be considered to be constant, and a new equilibrium constant, K w , is defined.
• In a weak acid-weak base titration, the graph has no vertical region so a pH probe must be used.
• The expression K a = ([H^+][A^-])/([HA]) is used to calculate the concentration of hydrogen ions ([H+]) in a solution when the concentration of salt ([A-]) and acid [HA] are known.
• Most indicators are suitable for this titration.
• The pH for buffers is all 0.1 mol dm-3, and the original volume is 25 cm3.
• Salts of a strong acid/weak base, such as NH4Cl, have a pH lower than 7 due to the formation of hydrogen ions when NH4+ ions react with water.
• There is a sudden increase in pH from about 3 to 8 around 25 cm3.
• If the concentration of salt and acid are equal, then the concentration of hydrogen ions is also equal, and pH is equal to -log K a.
• Salts of a weak acid/strong base, such as CH3COONa, have a pH higher than 7 due to the formation of hydroxide ions when CH3COO- ions react with water.
• In a strong acid-weak base titration, the graph starts at pH 1 and there is a slow increase in pH as the first 20 cm3 is added.
• In a strong acid-strong base titration, the graph starts at pH 1 and there is a slow increase in pH as the first 20 cm3 is added.
• There is a sudden increase in pH from about 3 to 11 around 25 cm3.
• For a solution containing 0.010 mol dm-3 ethanoic acid and 0.020 mol dm-3 sodium ethanoate, Ka for ethanoic acid is 1.78 × 10-5 mol dm-3.
• Phenolphthalein is a suitable indicator for this titration.
• In pure water, [H+] = [OH-] and Kw = [H+]2, Kw = [H+]2 = 1.00 × 10-14 and [H+] = 1.00 × 10-7 mol dm-3, pH = -log[H+] = 7.
• Basic buffers contain indicators, which are very weak acids that change colour as the pH changes.
• Acid buffers consist of a weak base and a salt of the same base, for example, ammonia solution and ammonium chloride.
• Buffers are used where pH is very important, such as regulating the pH of blood and storing biological molecules like enzymes.
• Kw = [H+][OH-] mol2 dm-6 at 25°C, Kw = 1.00 × 10-14 mol2 dm-6.
• Industrial examples of buffer use include fermentation and dye making.
• Buffer solutions consist of a weak acid and a salt of the same acid, for example, ethanoic acid and sodium ethanoate.
• To ensure that an indicator is suitable for a titration, the pH range over which it changes colour must lie within the vertical part of the titration curve.
• Adding an acid increases the amount of H+ ions and the equilibrium shifts to the left, removing the H+ ions by reaction with CH3COO-.
• The solution contains lots of NH4+ ions and NH3.
• Buffers are solutions whose pH stays relatively constant as a small amount of an acid or alkali is added.
• The solution contains lots of CH3COO- ions and CH3COOH.
• Adding a base increases the amount of OH- ions, causing the equilibrium to shift to the left, removing the OH- ions by reaction with NH4+.
• Adding an alkali increases the amount of OH- ions, causing the equilibrium to shift to the right, producing H+ ions from CH3COOH.
• The ionic equation for all neutralisation reactions is the reverse of the equilibrium above, i.e H+ + OH- ⇌ H2O.
• Adding an acid increases the amount of H+ ions, causing the equilibrium to shift to the right, producing more OH- ions from NH3 and H2O.