Chapter 20 - Acids, Bases + pH

Brønsted-lowry acid

a species that donates a proton

e.g. HCl - equation = HCl → H+ + Cl-

Brønsted-lowry base

a species that accepts a proton

e.g. NH3 - equation = NH3 + H+ → NH4+

conjugate-base pairs

two species that can be interconverted by the transfer of a proton

-can have more than 1 pair in an equation

example of conjugate-base pairs: HNO2 ⇌ H+ + NO2-

HNO2 and NO2- are the conjugate base pair

-HNO2 is the acid as it donates a proton to form the conjugate base NO2-

-NO2- is the conjugate base as it accepts a proton to form the acid HNO2

monobasic acid, example: HNO3 (DO EQUATION)

an acid that can only donate 1 proton to a base

example: HNO3 - equation = HNO3 → H+ + NO3-

dibasic acid, example: H2SO4 (DO EQUATION)

an acid that can donate 2 protons to a base

example: H2SO4 - equations = H2SO4 → H+ + HSO₄^-
HSO₄^- ⇌ H+ + SO₄^2-

equilibrium arrow used as weaker acid is produced (HSO₄^-)

tribasic acid, example: H3PO4 (DO EQUATION)

an acid that can donate 3 protons to a base

example: H3PO4 - equations = H₃PO₄ → H+ + H₂PO₄^-
H₂PO₄^- ⇌ H+ + HPO₄^2-
HPO₄^2- ⇌ H+ + PO₄^3-

equilibrium arrow used as weaker acid is produced

reaction of metal with acid - full + ionic equation → HCl + Mg

2HCl_(aq) + Mg_(s) → MgCl_2(aq) + H_2(g)

2H⁺ + Mg → Mg²⁺ + H₂

reaction of metal carbonate with acid - full + ionic equation → HCl + MgCO₃

2HCl_(aq) + MgCO_3(s) → MgCl_2(aq) + H₂O_(l) + CO_2(g)

2H⁺ + MgCO₃ → Mg²⁺ + H₂O + CO₂

-the carbonate is a solid so cannot dissociate into ions so remains undissolved

reaction of metal oxide with acid - full + ionic equation → HCl + MgO

2HCl_(aq) + MgO_(s) → MgCl_2(aq) + H₂O_(l)

2H⁺ + MgO → Mg²⁺ + H₂O

reaction of alkali with acid - full + ionic equation → HCl + NaOH

HCl + NaOH → NaCl + H₂O

H⁺ + OH^- → H₂O

strong acid

completely dissociates in aqueous solution

e.g. HCl

pH scale + strong acids

-a negative logarithmic scale of conc of H⁺

strong acids = [H⁺_(aq)] = [HA_(aq)]

pH and H⁺ equations

pH = -log₁₀[H⁺_(aq)]

[H⁺_(aq)] = 10^-pH

weak acid

partially dissociates in aqueous solution

e.g. CH3COOH

acid dissociation constant

-for weak acids

for HA ⇌ H⁺ + A^-

-expression: Ka = [H⁺] [A^-] / [HA] → units = mol dm^-3

Ka value

-larger Ka value means greater extent of dissociation so stronger weak acid - reverse for smaller

pKa and Ka equations

pKa = -log₁₀Ka

Ka = 10^-pKa

pKa value

-smaller pKa value means large Ka value so greater acid dissociation so stronger weak acid - reverse for larger

pH of weak acids

[H⁺_(aq)] = √Ka x [HA_(aq)]

limitations of scientific approximations

-assume that the acid is so weak that the conc at equilibrium = the same as the original conc

-but is actually [HA]_equilibrium > [HA]_original

-only works for weaker weak acids

water as amphoteric

-acts as an acid by donating a proton

-acts as a base by accepting a proton

deriving towards K_w

H2O ⇌ H+ + OH-

Ka = [H+][OH-] / [H2O] → Ka x [H2O] = [H+][OH-]

Kc is really small as water is a very weak acid so small amount of dissociation so [H2O] is constant

Kw equation + value

Kw = [H+][OH-] = 1 × 10^-14 mol² dm^-6

for strong bases [A-] = [H+] or [OH-]

pOH and [OH-] equations

pOH = -log[OH^-]

[OH^-] = 10^-pOH