topic 14

Acid: a substance that donates a proton (H+)
Base: A substance that accepts a proton (H+)

-A reaction where a proton is transferred from an acid to a base

HCl (acid) + H₂O (base) → H₃O⁺ + Cl⁻

-conjugate acid: when a base accepts a proton the species formed is the conjugate acid
-conjugate base: when a acid donates a proton, the species formed is the conjugate base

Cl⁻ is HCl's conjugate base because it is what remains after HCl loses a proton.

H₃O⁺ is H2O's conjugate acid because it forms when H₂O gains a proton.

monoprotic (monobasic/monoacidic): Donating or accepting on proton
Diprotic (dibasic/diacidic): donating or accepting two proton
Triprotic (tribasic/triacidic): donating or accepting three protons

a substance that can act as both an acid and a base

- in the reaction with HCl, H2O acts as a base
- in the reaction with NH3, H2O acts as an acid

strong acids: almost completely dissociated in aqueous solution (HCl → H⁺ + Cl⁻ )
(all strong acids go to completion)

weak acids: partially dissociated in aqueous solution (CH₃COOH ⇌ H⁺ + CH₃COO⁻)
(all weak acids form equilibrium)

pH = -log[H⁺] ([H⁺] is measured in mol dm-3)

pH < 7 → acidic
pH = 7 → neutral
pH > 7 → basic

For a weak acid HA ⇌ H⁺ + A⁻,

Ka​ = [H+][A-]/[HA]

- the intital concentration of the acid is equal to the concentration of the acid at equilibrium (since the acid slightly dissociates slightly[Ka is a very small value])

-[H+] comes only from the dissociation of the weak acid not of the water in the aqeuous solution

pKa = −log(Ka)
pKw = −log(Kw)

The acid is stronger, because it dissociates more in solution.

The acid is stronger, because it has a larger Ka (more dissociation).

The ionic product of water(Kw) is the equilibrium constant for the dissociation of water:
Kw = [H⁺][OH⁻]

1- Find [H⁺] = Kw / [OH⁻]
2- Then calculate pH = −log[H⁺]

A: A solution where the concentration of the acid or base is the same in moles per dm³.

pH increases by 1 unit each time because [H⁺] decreases (more dilute → less acidic).

pH increase by 0.5 units each time because weak acids do not fully dissociate.

strong acid + strong base → neutral pH (~7);
strong acid + weak base → acidic pH (<7);
weak acid + strong base → basic pH (>7)

[HA] <--> [A-][H+]

1. Calculate number of moles (m/mr)
2. workout the concentration (no. moles/ vol. of acid)
3. workout the [H+] = 10^(-pH)
4. Ka = [H+]^2/ (concentration - [H+])

- strong acid + strong base (pH = 7)
-Strong acid + weak base (pH < 7)
-Strong base + weak acid (pH > 7)
-Weak acid + Weak base ( hard to detect due to absence of a sharp decrease in graph)

Pick an indicator whose pH transition range falls within the steep vertical section of the curve (around the equivalence point).

The indicator partially dissociates and forms an equilibrium:

HIn <--> (H+) + (In-)

HIn = acidic form (color A)

In⁻ = basic form (color B)

A: The equilibrium shifts to HIn → acidic color dominates.

A: The equilibrium shifts to In⁻ → alkaline (basic) color dominates.

It shows a mixture color between the acidic and basic forms.

I.e. methyl orange = orange colour
acid: red
basic: yellow

When [In-] is 10x [HIn]

when [In-] is 10x [HIn]

second equivalent point = 1st equivalnce point x2 (if the graph is going up)

second equivalent point = 1st equivalnce point / 2 (if the graph is going down)

A buffer solution is one that resists change in pH when a small amount of acid or base is added

- weak acid + conjugate base
- weak base + conjugate acid

- the intial concentration of acid is the same as the concentration at equilibrium

- the concentration of [CH3COO-] is made of the dissacoiation of the salt

- the concentrations of the acid and salt are half of the orginal due to the volume doubling (n/v) (no. moles/ volume )

[H+] = Ka * [acid]/[salt]

1- calculate the no. Moles of both acid and base
2- workout total volume (total volume = vol. Of base+ vol. Of acid )
3- workout the new concentration of acid & base (no. Of moles/ total vol)
4-[H+] = Ka * ([acid]/[base])
5. pH= -log(H+)

1- find the concentration of [H+] = (10^-pH)
2- rearrange [H+] = Ka * ([acid]/[base])
3- find the value of ([acid]/[base])

-the pH of blood is controlled by the buffer mixture:
H2CO3<---> (HCO3 -) + (H+)

The buffer mixture maintains pH of blood around 7.35 - 7.45

1- there is a large resevoir of the weak acid and conjugate base
2- weak acid (H+) reacts with added alkali (OH-) & forms water
3- conjugate base reacts with added acid
4- the concentration of both acid & base remains almost constant so little change to pH