unit 8#
H₃O⁺ and H⁺ are used interchangeably in here, they’re basically kinda the same thing. [H₃O⁺] = [H⁺]
acids:
substances that produce H⁺ ions in aqueous solutions
bronsted-lowry acid: proton donor
HA(aq) + H₂O(l) ⇋ H₃O⁺(aq) + A⁻(aq)
HA: acid
H₂O: base
H₃O⁺: conjugate acid (what is formed after the proton is transferred to the base)
A⁻: conjugate base (everything that remains from the acid molecule after losing the proton)
in the above equation, HA donates a proton (H⁺) to the water molecule to form the hydronium, which is why water is considered a base.
this equation represents competition between the base and conjugate base for the proton
if HA is a strong acid, A⁻ is a weak conjugate base → water pulls the H⁺ from HA and equilibrium shifts to the right (so rxn goes to completion)
if HA is a weak acid, A⁻ is a strong conjugate base → A⁻ stronger than water → A⁻ pulls the proton from the hydronium ion and equilibrium shifts to the left (rxn does not go to completion)
strong acids:
dissociate completely at equilibrium, equilibrium lies far to the right
have very weak conjugate bases
weak acids:
dissociate partially at equilibrium, equilibrium lies far to the left
have strong conjugate bases
organic acids are weak acids
as Ka increases, acid strength increases
bases:
substances that produce OH⁻ ions in aqueous solutions
bronsted-lowry base: proton acceptor
B(aq) + H₂O(l) ⇋ BH⁺(aq) + OH⁻(aq)
B: base
H₂O: acid (remember that water is an amphoteric substance)
BH⁺: conjugate acid
OH⁻: conjugate base
amphoteric substances behave as either acids or bases, depending on what they’re reacting with (eg water or ammonia)
conjugates are total opposites of their original form. for example, strong acids produce weak conjugate bases
pH scale:
<7 = acidic
>7 = basic
7 = neutral
pH indicators:
litmus paper
acid → red, base → blue
phenolphthalein
acid → colorless, base → pink
methyl orange
acid → red, base → yellow
strong acids:
HI - hydroiodic
HBr - hydrobromic
HCl - hydrochloric
HNO₃ - nitric
HClO₄ - perchloric
H₂SO₄ - sulfuric
strong bases:
group one hydroxides
group two hydroxides starting from calcium and going down
strong acid-base rxns use single arrows & stoichiometry, while weak acid-base rxns use double arrows & equilibrium
remember that strong acids and bases are always aq
pH = -log[H₃O⁺]
pOH = -log[OH⁻]
pH + pOH=14 (at 25°C)
Kᴀ = [H⁺][A⁻]/[HA]
Kᴀ: acid dissociation constant at equilibrium
Kb = [OH⁻][BH⁺]/[B]
Kb:
Kᴡ = [H⁺][OH⁻]
Kᴡ = 10⁻¹⁴ at 25°C
Kᴡ increases as temperature increases
pKa=-log(Ka)
pKb=-log(Kb)
pKw=-log(Kw)
pKw = 14 at 25°C
pKa + pKb = 14 (at 25°C)
Ka x Kb = Kw (10⁻¹⁴ at 25°C)
%ionization = 100 x ([H⁺ or OH⁻]/[weak acid/base(initially)])
as weak acid molarity increases, percent ionization decreases
%ionization increases as the solution becomes more dilute
acid-base titrations:
titrant: the thing in the burette being dropped onto the analyte (usually base)
analyte: the thing in the beaker originally (usually acid)
the strength/curve of a titration depends on the reactions of the conjugates with water
end point: the point in a titration at which the indicator undergoes its color change
titration curves:
x axis → volume of titrant added, y axis → pH
when graphing them, the half equivalence point is horizontal, while the equivalence point is vertical
strong acid + strong base
originally, the pH is very low, and only the strong acid HA is present in the beaker. base is added until we reach the equivalence point, where pH = 7 & [HA] = [OH⁻]. at the equivalence point, the conjugates of our strong acid & base do not react with water to a significant extent. therefore, pH =7
after the equivalence point, OH⁻ will be in excess and will cause the graph to continue progressing to a pH > 7
weak acid + strong base
originally, the pH is relatively low (<7). as OH⁻ is added, the pH increases (still <7) until the buffer zone is reached. in the buffer zone (buffer zone resists changes in pH as OH⁻ is added), both the weak acid and its conjugate base are present. as we continue to add OH⁻ in the buffer zone, we reach the half equivalence point, where [acid] = 1/2[base], pH = pKa, and & [weak acid] = [conjugate base].
after the half equivalence point, [conj base] > [weak acid].
we continue to add OH⁻ until pH > 7. we reach the equivalence point somewhere after pH 7.
remember that the pH at the equivalence point is basic because of the reaction of the conjugate base with water to release OH⁻, not because the titrant is a strong base
conjugate base only species present at the equivalence point
as acid strength decreases:
starting pH increases
bigger jump to buffer zone
higher pH at equivalence point (bc weaker acid = stronger conjugate base)
strong acid + weak base
this time, unlike in the past two cases, the titrant is the acid, and the analyte is the base. so, the pH starts at pH > 7
H⁺ is added until the buffer zone is reached. in the buffer zone, both the weak base and its conjugate acid are present. we continue to add H⁺ until we reach the half equivalence point (still in the buffer zone, also still pH>7), where [base] = 1/2[acid], [conjugate acid] = [weak base], pOH = pKb.
after the half equivalence point, [conj acid] > [weak base]
we continue to add H⁺ until a pH < 7 is reached.
we eventually reach the equivalence point at pH < 7
remember that the pH at the equivalence point is acidic because of the reaction of the conjugate acid with water to produce hydronium (so H⁺), not because the titrant is a strong acid
at the equivalence point, the conjugate acid is the only species present
as base strength decreases
starting pH decreases
bigger drop to buffer zone
lower pH at equivalence point (bc weaker base = stronger conjugate acid)
HA(aq) + BOH(aq) → BA + H₂O(l)
neutralization rxn
double arrow if we have weak acid and weak base
if we have one strong acid / base we use single arrow
if we have a strong acid and a strong base reacting, determine the limiting & excess to figure out the pH from the excess reactant
actually if we have any acid-base rxn, find limiting and excess and net ionic equation to see what remains after excess then use what remains as an equilibrium reactant to find protons or hydroxide ions or whatever they need
(usually pH)like after an acid or base in excess reacts with the limiting reactant it has its own dissociation thingy
[acid] = [base] is the equivalence point
pH determined by the conjugate at the equivalence point
[acid] = 1/2[base] (weak acid, strong base) or [base] = 1/2[acid] (weak base, strong acid) is the half equivalence point
weak acid, strong base → pH = pKa
weak base, strong acid → pOH = pKb
buffered solution:
a solution that resists changes in pH (upon the addition of small amounts of strong acid or base)
if base is added, either the weak acid or conjugate base reacts with the added base
if acid is added, either the weak base or conjugate base reacts with the added base
can be either:
a weak acid and its conjugate base
partial neutralization of weak acid with a strong base
weak acid + salt of its conjugate base
a weak base and its conjugate acid
partial neutralization of weak base with a strong acid
weak base + salt of its conjugate acid
in a buffer, we need an acid capable of reacting with added OH⁻ and a base that can consume added H⁺
the acid & base in a buffer cannot react with each other!!!
to solve:
first, use stoichiometry. assume that the reaction of the buffered solution with the added H⁺ or OH⁻ goes to completion
second, use equilibrium. use the final concentrations found in the previous step as the initial concentrations in the equilibrium reaction. the equilibrium rxn will be the weak acid/base’s dissociation into its conjugate
henderson-hasselbalch equation:
pH = PKa + log([A⁻]/[HA])
[A⁻]=[HA], pH=PKa
[A⁻]>[HA], pH>PKa
[A⁻]<[HA], pH<PKa
diluting a buffer solution will not change its pH!!!
increasing the concentration of the buffer components (weak acid/base + conjugate) will increase the amount of acid or base the buffer can neutralize
THE LOWER THE CONCENTRATIONS OF THE COMPONENTS, THE LOWER THE BUFFER CAPACITY. THIS IS BECAUSE A SMALLER NUMBER OF MOLES ARE AVAILABLE TO REACT WITH ANY ADDED ACID OR BASE
strong acids experience an inductive effect due to the highly electronegative oxygen atoms
polarity in the molecule draws electrons away from the hydrogen atom, making it easily ionizable
electronegative elements in an acid tend to stabilize the conjugate base, and in turn strengthen the acid
the strength of bases depends on: how easily the lone pair picks up a hydrogen ion & the stability of the ions being formed
adding water to a strong acid increases pH because the addition of water decreases H⁺ to produce H₃O⁺