OCR A 5.1.3 Acids, Bases and Buffers

Arrhenius acid

releases H+ ion in aqueous solution

Arrhenius base

releases OH- ions in aqueous solution

Bronsted-Lowry acid

proton doner

Bronsted-Lowry base

proton acceptor

alkali

water soluble bases

-NaOH

-KOH

-Ca(OH)2

-NH3

-(aq) state symbol

bases

metal oxides/carbonates

(s) state symbol

chloric acid(VII)

HClO₄

conjugate acid base pair

NH₃ + H₂O —> NH₄⁺ + OH^-

B2 A1 A2 B1

NH₃ / NH₄⁺

H₂O / OH^-

-linked by transfer of proton

conjugate acid

any species that has gained a proton

conjugate base

any species that has lost a proton

polyprotic acids

donate more than 1 proton

monoprotic

1 mole of monoprotic acid will produce 1 mole of H+ ions

eg: HNO3

diprotic

1 mole of diprotic acid will produce 2 moles of H+ ions

H2SO4

triprotic

1 mole of triprotic acid will produce 3 moles of H+ ions

H3PO4

acid base reaction

-form pH neutral salts

-made from metal from base and non metal from acid

-H+ ions from acid react w/ OH- ions from alkalis to make water

ammonia reaction with acid

-makes ammonia salts

-no water

metal + acid

-salt+ hydrogen

weak acid dissociation

-eg: CH₃COOH & other carboxylic acids

-backward reaction favoured; less H+ produced

-only small amount of weak acid dissociates; assume [HA]_start = [HA]_equilibrium

-dissociation of acid is more than dissociation of water in the sol; assume [H⁺] = [A^-]

strong acid dissociation

-eg: HCl. H₂SO₄, HNO₃

-forward reaction favoured, more H+ produced

how does diprotic strong acid dissociate

-H₂SO₄ → H⁺ + HSO₄^- (first dissociation is strong acid)

-HSO₄^- → H⁺ + SO₄^2- (second dissociation is weak acid)

calculate [H+]

strong base dissociation

-eg: NaOH, KOH

-forward reaction favoured; more OH- produced

weak base dissociation

-NH₃

-backwards reaction favoured, less OH- produced

ionic product of water

-water exists in equilibria with its ions; water doesn’t just contain H2O molecules

-water dissociates weakly into H₃O + OH^-; assume that conc of water is constant

K_w

-same in sol at given temp

-value of 10^-14 mol²dm^-6

-value changes if temperature changes

K_w in pure water

-[H⁺] = [OH^-]

-K_w = [H⁺]²

ph

-logarithmic scale that measures conc of H+ ions in solution

pH of strong acid

HCl —> H⁺ + Cl^-

[strong acid] = [H+]

this only true for monobasic acid

pH of diprotic strong acid

[strong acid] = 2[H+]

pH of base

-dissociate fully; produce 1 OH- ion every base molecule

-[base] = [OH-]

pK_a

—logK_a

-used for weak acid

-higher the value, weaker the acid

K_a

-acid dissociation constant

-higher the value, less weaker

pH of weak acid

-HA_(aq) ←> H⁺_(aq) + A^-_(aq)

-we assume only small amount of weak acid dissociates:

-units; moldm^-3

how to measure pH experimentally

-using pH meter but calibrate before

-place in distilled water; shld read 7, if not adjust

-place in standard sol pH 4 and 10; if not show correct pH adjust

-rinse w/ distilled water between each pH sol

titration

-acid/base of known conc in burette

-base/acid unknown conc, known vol in conical flask w/ indicator

-add chemical in burette to conical flask until indicator changes colour (end-point)

-add drop by drop near this point

-read how much of chemical in burette added (bottom of meniscus, eye level)

-record to 2dp till u get concordant results (within 0.1)

strong acid/ strong base titration curve

strong acid/ weak base titration curve

weak acid/ strong base titration curve

weak acid/ weak base titration curve

equivalence point

-when the vol of 1 sol exactly reacts w/ vol of other sol

-sols have exactly reacted with each other using matching stoichiometry

-assume [H+] = [OH-]

-centre of vertical section of curve

how to choose suitable indicator

-must entirely change colour within vertical part of titration curve

-weak acid/ weak base titration; no indicator suitable

Buffer

-chemical that resists the change in pH when small amounts of acid/base added

-2 types; acidic, base

making buffer

mixture of:

-excess weak acid & conjugate base/ salt

-weak base & conjugate acid

-weak acid & salt

-partial neutralization of excess weak acid by strong alkali (mix both as acid is in excess it partially neutralized by strong alkali to make salt, sol left with weak acid and salt)

buffer solution equilibrium equation

what happens when you add acid to buffer solution

-[H+] increases

-H+ reacts w/ conjugate base A-

-POE shifts left

-most of H+ ions removed

-[HA] increases

what happens when you add alkali to buffer solution

-[OH-] increases

-OH- reacts with H+ to make H₂O

-HA dissociates to restore H+ ions

-[H+] increase

find pH of buffer solution

what happens when pH of blood below 7.35

acidosis:

-fatigue

-shortness of breath

-shock

-death

what happens when pH of blood above 7.45

alkalosis:

-muscle spasms

-light headedness

-nausea

buffer system in body

-carbonic acid-hydrogencarbonate buffer system

-H₂CO₃ + HCO₃^-

what happens when acid added to blood

-[H+] increases

-H+ reacts w/ HCO₃^-

-POE shifts left

-[H₂CO₃] increases

what happens when alkali added to blood

-[OH-] increases

-OH- reacts with H+ to make H₂O

-H₂CO₃ dissociates to restore H+ ions

-[H+] increase