Proton transfer reactions (acids and bases)

bronsted lowry acids and bases

• bronsted lowry acid: gives away a proton

• bronsted lowry base: accepts a proton using its lone pair of electrons

conjugate acids and bases

• conjugate base of an acid: species remaining after the acid donates a proton

• conjugate acid of a base: species remaining after the base accepts a proton

amphiprotic species

• species that can act as both proton donors and proton acceptors

• e.g water

the pH scale

• the acidity of an aqueous solution depends on the number of H+ ions in the solution

• pH = – log₁₀[H⁺]

• [H+] = 10^–pH

  • pH usually given to 2 d.p

  • logarithmic scale, each pH value is 10 times the value below it

• the lower pH, the more acidic, the higher pH, the more alkaline

pH of acids, bases and water explained

• acidic solutions

  • always have more H+ than OH- ions

  • [H+] always greater than 10^-7 mol dm-3, so pH always below 7

• basic solutions

  • always have more OH- ions than H+ ions

  • [H+] is always smaller than 10^-7 mol dm-3, so pH always above 7

• pH of water

  • at 298K, water has equal amounts of OH- and H+ ions with concs of 10^-7 mol dm-3

how to measure pH

• pH meter:

  • most accurate

  • connected to pH electrode which shows the pH value of solution

• universal indicator

  • less accurate

  • dipped into solution of acid and changes colour

  • colour is compared to a chart which shows different pH values

the ion product of water (Kw)

• an equilibrium exists in water, where a few water molecules dissociate into H+ and OH- ions

  • H₂O (l) ⇌ H⁺ (aq) + OH^– (aq)

• equilibrium constant for this reaction is:

  • Kc = [H+][OH-] / [H2O]

  • Kc x [H2O] = [H+][OH-]

• since the conc of H+ and OH- ions is very small, conc of water is considered to be a constant

• so K_w = [H⁺] [OH^-]

• K_w (ion product of water) =  K_c x [H₂O] =  1.00 x 10^-14 at 298K

how does temp affect Kw

• ionisation of water is an endothermic process, so due to le chateliers principle, an increase in temp will favour the forward reaction

  • [H+][OH-] increases

  • magnitude of Kw increases

  • pH decreases

• increasing temp decreases pH of water (more acidic)

• decrease temp increases pH of water (more basic)

strong and weak acids

• strong acid: dissociates almost completely in aqueous solutions

  • position of equilibrium extremely right, irreversible

  • e.g HCl, HBR, HI, HNO3, H2SO4

  • solution formed is highly acidic due to high conc of H+ ions

  • pH can be calculated if concentration of strong acid is known

• weak acids: acids that partially dissociate in aqueous solutions

  • position of equilibrium is more to the left equilibrium is established (reversible reaction)

  • e.g ethanoic acid

  • solution formed is less acidic due to lower concentration of H+ ions

• strength of bronsted lowry acid depends on the ease with which it dissociates to released H+ ions

strong and weak bases

• strong base: base that dissociates almost completely in aqueous solution

  • position of equilibrium extremely right, irreversible

  • e.g NaOH

  • solution formed highly basic due to high conc of OH- ions

• weak base: base that partially dissociates in aqueous solution

  • position of equilibrium more left, equilibrium is established (reversible)

  • e.g NH3

  • solution formed less basic due to lower conc of OH- ions

strength of conjugate acids and bases

• strong acids produce weak conjugate bases

  • because reverse reaction is virtually non existent

• weak acids produce strong conjugate bases

• strong bases produce weak conjugate acids

• weak bases produce strong conjugate acids

distinction between strong and weak acid

• pH value

  • the stronger the acid, the greater the H+ conc, the lower the pH

• electrical conductivity

  • stronger acids give higher readings conductivity meter because they have more ions in solution

• reactivity

  • as conc of H+ is much greater in strong acids the rate of reaction with metals and metal compounds is greater than of weak acids of the same conc

neutralisation reaction

• acid + base —> salt + water

• H+ + OH- —> H2O

• spectator ions not involved in formation of water form the the salt

summary of neutralisation reactions

• acid + metal —> salt + hydrogen gas

  • extent of reaction depends on reactivity of metal and strength of acid

• acid + metal oxide —> salt + water

• acid + metal carbonate —> salt + water + carbon dioxide

• acid + metal hydrogencarbonate —> salt + water + carbon dioxide

SA + SB pH curve

• e.g NaOH + HCl

• shows how the pH of a solution changes as the acid/base is added in a strong acid/strong base titration

  • pH of acid: where the curve starts on y-axis, roughly 1

  • equivalence point: “stoichiometric end of reaction”, when the acid/base neutralisation is complete

• calculating pH after x volume of base added:

  1. calculate moles of acid and base (conc x volume of base added)

  2. determine moles of excess acid

  3. determine new volume (vol of acid + added volume of base)

  4. [H+]= moles of excess acid / new volume

  5. -log[H+] = pH

pOH scale

• basicity of aqueous solution depends on number of hydroxide ions

• pOH = -log [OH^-]

• pH = 14 - pOH

relationship between H+ OH- pH and pOH

dissociation of a weak acid

• weak acids HA dissociate partially in water

  • HA (aq) ⇌  A^- (aq) + H⁺ (aq)

• at equilibrium, majority of HA molecules remain unreacted

• position of equilibrium is more left, equilibrium is established

• acid dissociation constant, Ka = [A-][H+] / [HA]

  • values of Ka are very small

dissociation of a weak base

• weak bases (B) ionise in water

  • B (aq) + H₂O (l) ⇌  BH⁺ (aq) + OH^- (aq)

• equilibrium is established

• base dissociation constant Kb = [BH-][OH-] / [B]

assumptions made when calculating pKb, Kb, pKa, Ka

• initial conc of acid is about the equilibrium conc of acid

• [A^-] = [H⁺]

• negligible ionisation of the water, so [H⁺] is not affected

• The temperature is 298 K

pKa and pKb

• pKa = -logKa

• pKb = -logKb

Ka and Kb of conjugate acid-base pairs

• K_a x K_b = K_w

salt hydrolysis

• ionic salt is formed from neutralisation of acid and base

  • HA (parent acid) + MOH (parent base) —> M+A- (salt) + H2O

• ionic salt formed will dissociate in water

  • hydrolysis is where water is used to break a bond within compound, resulting in aqueous ions

• pH of salt depends on the strength of parent acid and base used

summary of salt hydrolysis reactions

• SA + SB

  • neutral

  • weak conjugate acid and base ions therefore no hydrolysis occurs

  • equivalence point= pH 7

• SA + WB

  • acidic

  • strong conjugate acid of WB is formed, which reacts with water to form H+ ions, therefore solution becomes more acidic

    • M+ + H2O —> MOH + H+

  • equivalence point = <7 pH

• WA + SB

  • alkaline

  • strong conjugate base of WA is formed, which reacts with water to form OH- ions, therefore solution becomes more basic

    • A- + H2O —> HA + OH-

  • equivalence point = >7 pH

• WA + WB

  • depends on relative Ka and Kb

WA + SB pH curve

• e.g NaOH + CH3COOH

• starting pH roughly 3 because WA

• initial rise in pH steep as neutralisation of WA by strong base is rapid

• ethanoate ions (conjugate base to weak acid) are formed, creating a buffer

• buffer formed will resist changes in pH so pH rises gradually in buffer region

• pKa = pH [H+] = Ka at half equivalence point

• equivalence point = >7 pH

SA + WB pH curve

• e.g NH3 + HCl

• starting pH roughly 11 (WB)

• pH falls as WB begins to be neutralised and forms conjugate acid

  • creates buffer region so pH falls slowly

• half equivalence point (where half the amnt of WB has been neutralised)

• pKb = pOH [OH-] = Kb at half equivalence point

• equivalence point = <7 pH

acid-base indicators

• weak acid which dissociates to make anion of a different colour

• e. g HInd ⇌ H⁺ (aq) + Ind^– (aq)

• HIn and its conjugate base Ind- are different colours

• if solution acidic, equilibrium shifts left and HInd colour is shown

• if solution basic, equilibrium shifts right, and Ind- colour is shown

• universal indicator is a mix of many indicators with a wide pH range of colour

indicator calculations

• pH at which transitions between colours occur:

  • Ka

• endpoint of reaction is when there is balance of [HIn] and [In-]

  • Ka = [H+][Ind-] / [HInd] = [H+]

• take negative logs of both sides

  • pKa = pH

  • pKa of an indicator is the same as its pH at end point

• colour change takes place over range of pH = pKa ± 1

choosing the right indicator

• indicator is appropriate if it has an endpoint range that coincides with pH at equivalence point

buffer solutions

• solution which resists changes in pH when small amounts of acid/base are added

• used to keep the pH almost constant

• consists of weak acid- conjugate base or weak base conjugate acid

acidic buffers

• e.g aqueous mixture of sodium ethanoate and ethanoic acid

• ethanoic acid is weak, ionises in solution to form relatively low conc of ethanoate ions

• sodium ethanaote is salt, fully ionises in solution

• contains reserves supplies of the acid and its conjugate base

  • high conc of ethanoic acid due to partial ionisation

  • high conc of ethanoate ions due to full ionisation of sodium ethanaote

• in buffer solution, ethanoic acid in equilibrium with hydrogen and ethanoate ions

adding H+ to acidic buffer

• equilibrium shifts left as H+ ions react with CH3COO- to form more CH3COOH until equilibrium reestablished

• large reserve supplies of both CH3COO- and CH3COOH so pH remains relatively constant

adding OH- to acidic buffer

• reacts with H+ to form water

• equilibrium position shifts right, more CH3COOh molecules ionise to form more H+ and ethanoate ions until equilibrium reestablished

• pH remains constant

basic buffers

• made by mixing a solution of a weak base with its salt

  • e.g NH3 and NH4Cl

• ammonia partially dissociate in water forming NH4+ and OH-

• NH4Cl fully dissociates forming NH4+ and Cl- ions

• adding acid/base will result in no change in pH

  • additional H+ will combine with NH3 to make NH4+

  • additional OH- will combine with NH4+ to form NH3 and H2O

buffer calculations

• to determine pH of buffer:

  • [H+] needed, found using:

    • [H+]= Ka [acid]/[salt]

    • pH = pKa + log[salt] /[ acid]

• to determine pOH

  • [OH-] needed, found using

    • [OH-] =Kb [base] / [salt]

    • pOH = pKb + log[salt] / [base]

factors affecting buffers

• dilution

  • Ka and Kb are equilibrium constants so not changed by dilution

  • doesn’t change ratio of acid/base conc to salt conc

• temperature

  • constant temp must be maintained when using buffers

  • temp affects values of Ka and Kb