Chapter 5 Equilibra

• is a species that can donate a proton • For example, hydrogen chloride (HCl) is a Brønsted acid as it can lose a proton to form a hydrogen (H+) and chloride (Cl-) ion HCl (aq) → H+ (aq) + Cl- (aq)

• is a species that can accept a proton • For example, a hydroxide (OH-) ion is a Brønsted base as it can accept a proton to form water OH- (aq) + H+ (aq) → H2O (l)

• same rate as the reactants are used • This means that at equilibrium, both reactants and products are present in the solution

is two species that are different from each other by an H+ ion • Conjugate here means related • In other words, the acid and base are related to each other by one proton difference

transfer of a proton

• acidity or basicity of an acid or alkali • The pH scale goes from 0 to 14 • Acids have pH between 0-7 • Pure water is neutral and has a pH of 7 • Bases and alkalis have pH between 7-14

• The pH can be calculated using: pH = -log10 [H+] where [H+] = concentration of H+ ions (mol dm-3) • The pH can also be used to calculate the concentration of H+ ions in solution by rearranging the equation to: [H+] = 10-pH

• acidic dissociation constant • It is the equilibrium constant for the dissociation of a weak acid at 298 K • For the partial ionisation of a weak acid HA the equilibrium expression to find Ka is as follows: HA (aq) ⇌ H+ (aq) + A- (aq)

• The concentration of hydrogen ions due to the ionisation of water is negligible • The dissociation of the weak acid is so small that the concentration of HA is approximately the same as the concentration of A-

dissociation • A high value of Ka means that: • The equilibrium position lies to the right • The acid is almost completely ionised • The acid is strongly acidic • A low value of Ka means that: • The equilibrium position lies to the left • The acid is only slightly ionised (there are mainly HA and only a few H+ and A- ions) • The acid is weakly acidic

• very low, pKa values are used instead to compare the strengths of weak acids with each other pKa= -log10Ka • The less positive the pKa value the more acidic the acid is

• ionic product of water • It is the equilibrium constant for the dissociation of water at 298 K • Its value is 1.00 x 10-14 mol2 dm-6 • For the ionisation of water the equilibrium expression to find Kw is as follows: H2O (l) ⇌ H+ (aq) + OH- (aq)

• is very low, only small amounts of H+ and OH- ions are formed • The concentration of H2O can therefore be regarded as constant and removed from the Kw expression • The equilibrium expression therefore becomes: • Kw* = [H+] [OH-] • As the [H+] = [OH+] in pure water, the equilibrium expression can be further simplified to: • Kw* = [H+]2

• If the concentration of H+ of an acid or alkali is known, the pH can be calculated using the equation: pH = -log [H+] • Similarly, the concentration of H+ of a solution can be calculated if the pH is known by rearranging the above equation to: [H+] = 10-pH

• ionised in solution HA (aq) → H+ (aq) + A- (aq) • Therefore, the concentration of hydrogen ions ([H+]) is equal to the concentration of acid ([HA]) • The number of hydrogen ions ([H+]) formed from the ionisation of water is very small relative to the [H+] due to ionisation of the strong acid and can therefore be neglected • The total [H+] is therefore the same as the [HA]

• completely ionised in solution BOH (aq) → B+ (aq) + OH- (aq) • Therefore, the concentration of hydroxide ions ([OH-]) is equal to the concentration of base ([BOH]) • Even strong alkalis have small amounts of H+ in solution which is due to the ionisation of water

Kw = [H+] [OH-] • Since Kw is 1.00 x 10-14 mol2 dm-6 • Once the [H+] has been determined, the pH of the strong alkali can be founding using pH = -log[H+] • Similarly, the ionic product of water can be used to find the concentration of OH- ions in solution if [H+] is known

solution in which the pH does not change a lot when small amounts of acids or alkalis are added • A buffer solution is used to keep the pH almost constant • A buffer can consists of weak acid - conjugate base or weak base - conjugate acid

aqueous mixture of ethanoic acid and sodium ethanoate

• weak acid and partially ionises in solution to form a relatively low concentration of ethanoate ions

• The equilibrium position shifts to the left as H+ ions react with CH3COO- ions to form more CH3COOH until equilibrium is re-established • As there is a large reserve supply of CH3COO- the concentration of CH3COO- in solution doesn’t change much as it reacts with the added H+ ions • As there is a large reserve supply of CH3COOH the concentration of CH3COOH in solution doesn’t change much as CH3COOH is formed from the reaction of CH3COO- with H+ • As a result, the pH remains reasonable constant

• The OH- reacts with H+ to form water OH- (aq) + H+ (aq) → H2O (l) • The H+ concentration decreases • The equilibrium position shifts to the right and more CH3COOH molecules ionise to form more H+ and CH3COO- until equilibrium is re-established CH3COOH (aq) → H+ (aq) + CH3COO- (aq) • As there is a large reserve supply of CH3COOH the concentration of CH3COOH in solution doesn’t change much when CH3COOH dissociates to form more H+ ions • As there is a large reserve supply of CH3COO- the concentration of CH3COO- in solution doesn’t change much • As a result, the pH remains reasonable constant

• In humans, HCO3- ions act as a buffer to keep the blood pH between 7.35 and 7.45 • Body cells produce CO2 during aerobic respiration • This CO2 will combine with water in blood to form a solution containing H+ ions CO2 (g) + H2O (l) ⇌ H+ (aq) + HCO3- (aq)