Topic 9 and 18 – Acids and Bases
Amphoteric: Can have the properties of both a base and an acid, depending on whether it is reacting with a base or an acid.
Brønsted-Lowry: An acid is defined as a proton donator, while a base is a proton acceptor.
Buffer: A solution that resists changes in pH when small amounts of acid or alkali are added to it. When a small amount of acid is added, the excess of H+ ions causes the equilibrium to shift to the left balances the difference. When a small amount of alkali is added, the OH- ions react with the H+ ions to form water. The decrease in [H+] is compensated for by an equilibrium shift to the right. Vice versa for alkali buffers. Buffer solutions are made by several means:
**strong base + excess weak acid; **strong acid + excess weak base; **weak acid + same acid’s salt; **weak base + same base’s salt.
Charge density: Charge divided by size.
Concentrated: High number of moles of solute per dm3 of solution.
Conductivity: The more a solution is dissociated into its ions, the greater its conductivity.
Conjugate: The species remaining after an acid has lost a proton (conjugate base) or a base has gained one (conjugate acid). pKa + pKb = pKw
Diprotic: Where one mole of sulphuric acid produces two moles of hydrogen ions, e.g. H2SO4.
End point: The point at which the indicator changes color most rapidly.
Equimolar: Containing equal concentrations.
Equivalence point: Where the acid and base are in equimolar quantities. Exactly enough to react with each other.
Indicator: A weak acid or base in which the dissociated form is a different color to the undissociated form. The end point occurs when the pH is approximately equal to the pKin value, assuming that the color changes when [In-] [HIn]. Ideally, the end point corresponds to the equivalence point in a titration.
Lewis theory: An acid is defined as an e- acceptor (e.g. BF3) and a base is an e- donator (e.g. NH3) by dative bond.
Monoprotic: Where one mole of the acid produces one mole of hydrogen ions, e.g. HCl.
pH: Power of hydrogen. – log [H+]
Salt hydrolysis: The process by which a salt is broken down by water. The acidity of the salt depends on:
1) its derivations (salts derived from a strong acid and weak base will be acidic in solution, considering that the ions of the weak base will combine, leaving an excess of H+. The strong acid will not combine as it is completely dissociated, as per definition. Vice versa with strong acid and weak base); 2) charge density of cation (high charge density results in strong attraction to the lone pair of one of six water molecules surrounding the ion, a process by which the water molecule loses a hydrogen ion, i.e. leaving the solution acidic).
Strong: An acid or a base that dissociates completely into its ions. Ka >> 1. Some strong acids: hydrochloric, sulphuric, nitric (weaker than other two). Strong bases: hydroxides of alkali metals.
Titration: Technique for quantitative measure of concentration of a solution. Consider when 90% of the required base has been added to a 1M strong acid that is to be neutralized. Only 10% of the acid remains, meaning 0.01M, giving a pH of 2. When 99% has been added, the pH is 3, etc.
Water, ionic product of: The equilibrium constant for the dissociation of water into its ions, where [H2O] is taken to be constant. Value of Kw increases as temperature is increased, as the dissociation is an endothermic process. pKw
Weak: An acid or base that only slightly dissociates into its ions. Ka << 1. Some weak acids: ethanoic, carbonic. *Weak bases: ammonia, aminoethane.