Ionic Equilibrium I (Acid-Base)

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Arhenius theory of acids and bases

An Arhenius acid is a substance that increases [H+] when dissolved in water / produces H+ ions in aqueous solution.

An Arhenius base is a substance that increases [OH+] when dissolved in water / produces OH- ions in aqueous solution.

Bronsted-Lowry theory of acids and bases

Bronsted-Lowry acid is a species that donates protons to a base, i.e. a proton donor

Bronsted-Lowry base is a species that accepts protons from an acid, i.e. a proton acceptor

Conjugate acid-base pair

Every acid has a corresponding conjugate base which is formed after the acid has donated its proton or H+.

For the following acid, HX <-> H+ + X-HX and X- is a conjugate acid-base pair which differs by a proton.

Bronsted-Lowry acid (HA) loses proton => A- is conjugate base

Bronsted-Lowry base (B) accepts proton => BH+ is conjugate acid

eg. HCl/Cl-

H2O/OH-

CH3COOH/CH3COO-

NH3/NH4+

H2O/H3O+

CH3NH2/CH3NH3+

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Lewis Acids and Bases

A Lewis acid is a substance that can accept lone pair electrons

A Lewis base is a substance that can donate its lone pair electrons

Summary

Arrhenius

A: increase [H+]

B: increase [OH-]

Bronsted-Lowry

A: proton-donor

B: proton-acceptor

Lewis

A: electron pair-acceptor

B: electron pair-donor

Strength of Acids

Stong acid: completely ionises in water

HCl(aq) -> H+(aq) + Cl-(aq)

-> most mineral acids (HI, HBr, H2SO4, HNO3) {HF is weak acid}

Weak acid: partially ionises in water

CH3COOH + H20(l) <> CH3COO-(aq) + H3O+(aq)

-> most organic acids (CH3COOH, H2C2O4, C6H5COOH, C6H5OH, HCN)

General rule to determine if an oxoacid is strong or weak

No. of oxygen - No. of hydrogen atoms >/= 2 --> Strong acid

E.g. of strong oxoacids: H2SO4, HNO3, HClO4

E.g. of weak oxoacids: H2SO3, HNO2, HClO

Strength of Bases

Stong base: fully ionises in water

NaOH(s) -> Na+(aq) + OH-(aq)

(NaOH, KOH, Ba(OH)2)

Weak base: partially ionises in water

NH3 + H2O <> NH4+(aq) + OH-(aq)

(NH3, CH3NH2, Na2CO3)

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pH = -log10[H+(aq)]

or [H+] = 10^-pH

=> low pH, high conc of H+, hence more acidic

pOH = -log10[OH-(aq)]

or [OH-] = 10^-pOH

Autoionisation of water

pure water conducts electricity slightly, due to the fact that water dissociates slightly into ionsH20(l) <> H+(aq) + OH-(aq), ^H>0(endo)

Since H2O is in large amount and only a small fraction of water molecules are ionised, [H2O] remains almost constant

Ionic product of water, Kw = [H+(aq)][OH-(aq)] mol2dm-6

-value of Kw depends on temperature

-in pure water at 25C, [H+(aq)] = [OH-(aq)] = 10^-7 moldm-3

Kw = [H+(aq)][OH-(aq)] = 1.00 x 10^-14 mol2dm-6 at 25C

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[H+] = [OH-] => neutral

[H+] > [OH-] => acidic

[H+] < [OH-] => alkaline

-endothermic process

-when temp increases, the P.O.E shits to the right to absorb extra heat energy. Both [H+] and [OH-] increase.

-Since Kw = [H+][OH-], Kw increases.

pKw = pH + pOH = 14 at 25C

Steps in calculating pH of strong acid/base

write eqn

determine [H+] (or [OH-])

determine pH (or pOH) using formula: pH = -log10[H+] / pOH = -log10[OH-]

for bases, pH = pKw - pOH = 14 - pOH (at 25C)

Self-ionisation of water can be ignored as long as [H+] or [OH-] is greater than 1.00 x 10^-5 moldm-3

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Dissociation of weak acids/bases

Acid dissociation constant, Ka

HA(aq) + H2O(l) <-> H3O+(aq) + A-(aq)

Ka = [H3O+]eqm[A-]eqm/[HA]eqm

pKa = -log10Ka

-Ka is temperature dependent

-the higher the Ka value or lower the pKa, the stronger the acid

Base dissociation constant, Kb

B + H20 <-> BH+ + OH-

Kb = [BH+]eqm[OH-]eqm/[B]eqm

pKb = -log10Kb

-Kb is temperature dependent

-the higher the Kb value or the lower the pKb, the stronger the base

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Degree of dissociation, alpha

For strong acids/bases which ionise completely, alpha = 1

For weak acids/bases which ionise partially, alpha << 1

Ka = (alpha)2[HA]initial

alpha = sqroot(Ka/[HA]initial)

Kb = (alpha)2[B]initial

alpha = sqroot(Kb/[B]initial)

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Relationship of Kw, Ka and Kb

Ka x Kb = [H3O+][OH-]

Ka x Kb = Kw

Kw = 10^-14 at 25C

pKa + pKb = pKw

pKw = 14 at 25C

Since Kw is a constant, the higher the Ka, the lower the Kb

-stronger acid gives weaker conjugate base

e.g. phenol is a weaker acid than benzoic acid, hence phenoxide is a stronger conjugate base than benzoate ion

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Steps in calculating pH of weak acid/base

write eqn to illustrate dissociation

write the appropriate dissociation expression (Ka or Kb)

Determine [H+] or [OH-]

Determine pH (or pOH) using formula

For bases pH = pKw - pOH = 14 - pOH (at 25C)

for a weak monobasic acid solution

Ka = [H+]^2/[HA]initial

Hence, [H+] = sqroot(Ka[HA]initial)

for a weak monoacidic base solution

Kb = [OH-]^2/[B]initial

Hence, [OH-] = sqroot(Kb[B]initial)

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Properties of Salt

-insoluble in water or sparingly soluble in water

-soluble in water, forming neutral or acidic or alkaline solution

4 types of soluble salts

-Strong base and strong acid: neutral salt

-Strong base and a weak acid: basic salt

-Strong acid and weak base: acidic salt

-Weak base and weak acid: acidic, basic or neutral salt

Salt of S.A. and S.B.

-dissolves and dissociates completely

-ions do not undergo hydrolysis in water

-conc of H+ and OH- in water remain the same

=> neutral (eg. HCl + NaOH -> NaCl)

Salt of S.A. and W.B.

-dissolves and dissociates completely

-cation of weak base -> relatively strong conjugate acid and hydrolyses in water to give H3O+

=> acidic (eg HCl + NH3 -> NH4Cl)

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Salt of S.B. and W.A.

-dissolves and dissociate completely

-anion of a weak acid -> relatively strong conjugate base and hydrolyses in water to give OH-

=> alkaline (eg CH3COOH + NaOH -> CH3COO-Na+)

Salt of W.A. and W.B.

-dissolves and dissociate completely

-both ions hydrolyse in water=> slightly acidic, slightly alkaline or neutral depending on the conc of H3O+ and OH- ions formed

(eg CH3COOH + NH3 -> CH3COO-NH4+)

Ka > Kb => acidic

Ka < Kb => alkaline

Ka (approx =) Kb => neutral

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Steps in calculating pH of conjugate Acid/Base

write dissociation eqn

write dissociation expression (Ka or Kb)

recall Kb(of A-) = Kw/Ka(of HA) and Ka(of BH+) = Kw/Kb(of B)

determine [H+] or [OH-]

determine pH (or pOH) using formula

for bases, pH = pKw -pOH = 14-pOH at 25C

Acidic salt

Ka (approx =) [H+]^2/[salt]initial

= [H+]^2/[SCA]initial

Basic salt

Kb (approx =) [OH-]^2/[salt]initial]

=[OH-]^2/[SCB]initial

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Buffer

a buffer solution is a solution which resists changes in pH when a small amount of H+ or OH- is added to it or upon dilution

A buffer solution must have:

- large amount of acid to react with any OH- added

- large amount of base to react with any H+ added

- the acid and base components must not neutralize each other

2 types of buffer

1. acidic buffer - a soln of weak acid and its salt (contains conjugate base of acid)

[eg CH3COOH & CH3COONa]

2. basic buffer - a soln of a weak base and its salt (contains conjugate acid of base)

[eg NH3 & NH4Cl]

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Acidic buffer

--on adding a small amt of H+

CH3COO- + H+ -> CH3COOH

the small amt of H+ ions are removed by the large amt of CH3COO- in the buffer.

=> pH remains almost unchanged as [H+] does not change significantly

--on adding a small amt of OH-

OH- + CH3COOH -> CH3COO- + H2O

the small amt of OH- ions are removed by the large amt of CH3COOH in the buffer

=> pH remains almost unchanged as [OH-] does not change significantly

Basic buffer

--on adding a small amt of OH-

OH- + NH4+ -> NH3 + H20

the small amt of OH- ions are removed by the large amt of NH4+ in the buffer

=> pH remains almost unchanged as [OH-] does not change significantly

--on adding a small amt of H+

H+ + NH3 -> NH4+

the small amt of H+ ions are removed by the large amt of NH3 in the buffer

=> pH remains almost unchanged as [H+] does not change significantly

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For a buffer soln,

Ka = [H+][A-]/[HA]

approx = [H+][salt]/[acid]

-lg Ka = -lg [H+][salt]/[acid]

-lg Ka = -lg [H+] - lg [salt]/[acid]

Henderson-Hasselbalch eqn for acidic buffer:

pH = pKa + lg [salt]/[acid]

Henderson-Hasselbalch eqn for basic buffer:

pOH = pKb + lg [salt]/[base]

Steps in calculating pH/pOH of buffer solutions

identify type of soln

if acidic buffer use pH = pKa + lg [salt]/[acid] alkaline buffer use pOH = pKb + lg [salt]/[base]

determine conc of [HA] and [A-] (or [B] and [BH+]) in the buffer mixture

sub eqm conc into Henderson-Hasselbalch eqn and solve for pH

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Effective buffer - Maximum Buffer Capacity

[Salt]/[Acid] = 1

if [salt] >> [acid], the buffer is effective in removing H+ but not effective in removing OH-

[Salt]/[Base] = 1

if [salt] >> [base], the buffer is effective in removing OH- but not effective in removing H+

At MBC, the soln has equal ability in removing H+ and OH- by the same amt to ensure that the pH remains almost unchanged

Acidic buffer: pH = pKa + lg 1 = pKa

Basic buffer: pOH + lg 1 = pKb

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Uses of buffer solutions

1. in calibration of pH meters

2. many biological systems where a change in pH can have a great effect on the functioning of a cell. Therefore, eg all injections & eye drops are buffered as human blood & tears have a specific pH

3. in maintaining pH of soils for growth of certain plants eg azealea, rhododendron - pH 4

Blood buffer system

-narrow range of 7.35-7.45

-enzymes only work over a small range of pH

-consists of protein, a mixture of phosphates(HPO4^-2/H2PO4-) and a mixture of carbonates(HCO3-/H2CO3)

1. proteins are made up of amino acids which contain acidic and basic groups which act as buffers

2. CO2 produced as the by-pdt of metabolic rxns in the body, dissolves to form carbonic acid which dissociates partially to give hydrogen carbonate/bicarbonate ion.

CO(g) + H2O(l) <> H2CO3

H2CO3 <> HCO3- + H+

HCO3-/H2CO3 forms a buffer system in blood plasma

-if the conc of H+ increases (eg from lactic acid which is produced during exercise), HCO3- present reacts and removes the H+ ions

H+ + HCO3- -> H2CO3

-When [OH-] increases, H2CO3 present reacts and removes OH-

OH- + H2CO3 -> H20 + HCO3-

-hydrogencarbonate buffer is constantly replenished in the body from the dissolved CO2 produced during normal metabolism

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Indicators for acid-base titrations

End point of titration: the end point of a titration occurs when the indicator changes colour.

Choice of indicator: based on the fact that its pH range for colour change must coincide with the region of rapid pH change in the titration curve (vertical section of the titration curve)

-methyl orange: red(acid) / yellow(base)

[3.2 - 4.4]

-screened methyl orange: violet(acid) / green(base)

[3.2 - 4.4]

-phenolphthalein: colourless(acid) / pink(base)

[8.2 - 10.0]

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Acid-base titration

Strong acid - strong base titration

Before titration:

-at the beginning, pH starts low, showing the high [H+] of the strong acid (or a weak acid with high conc)

-at the initial point the solution contains only strong acid

When acid is partially neutralised (after initial, before equivalence point):

-pH gradually increases, the acid is gradually neutralised by the added base

-in this region, the mixture is caused by the remaining strong acid and salt formed. the pH of the mixture is caused by the remaining strong acid present as the salt formed is neutral

At equivalence point:

-occurs when the no. of moles of alkali added just neutralises the no. of moles of acid present

-at this point, the mixture contains only the salt and water formed. since the salt formed is neutral, the pH at this point is expected to be 7.

When excess alkali is added:

-alkali has nothing to react with, so the pH rises towards a final value of 14

-when excess alkali is added, the mixture contains excess strong alkali and the salt formed. the pH of the mixture is caused by the strong alkali since the salt formed is neutral

Suitable indicator:

-region of rapid pH change: 4 - 10

-end-point: the point in a titration where the indicator changes colour

-for strong acid-strong base titration: methyle orange / screened methyl orange & phenolphthalein

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Weak acid - strong base titration

Max buffer capacity:

-occurs when [acid] = [salt], pH = pKa

-this point occurs at 1/2 V cm3 where V cm3 is the volume of base required to neutralise the acid completely

At equivalence point:

-the pH at equivalence point is more than or equal to 7 due to the hydrolysis of CH3COO-Na+ salt

Suitable indicator:

-the region of rapid pH change: 7 - 10

-suitable indicator: phenolphthalein

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Strong acid - weak base titration

At equivalence point:

-the pH at equivalence point is less than or equal to 7 due to hydrolysis of NH4Cl salt

Max buffer capacity:

-occurs when [base] = [salt], pOH = pKb

-this point occurs at 2V cm3 where V cm3 is the volume of base required to neutralise the acid completely

Suitable indicator:

-region of rapid pH range: 4 - 7

-suitable indicator: methyl orange

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Amino acids titration curve

the lower the pKa, the stronger the acid. hence, -COOH is assigned the lowest pKa. the side chain amino group in lysine is more basic than the a-amino group. this is due to a longer alkyl grp which is more electron-donating and causing the lone pair of electrons on N on side chain amino group to be more available for protonation.

OR

the amine grp on the a-carbon is less basic due to close proximity to an electron-withdrawing -COOH grp causing the lone pair of electrons on N on the a-carbon amino grp to be less available for protonation