3.1.12 Acids & Bases Flashcards
what is meant by a ‘Bronsted-Lowry acid’?
proton donor
what is meant by a ‘Bronsted-Lowry base’?
proton acceptor
what makes an acid strong? what equation can be used to show this?
- strong acid fully dissociates/ionises in solution to release a high [H+]
- HA + H2O -> H3O+ + A-
(H3O+ = H+; can be used interchangeably)
what makes an acid weak and what equation can be used to show this?
- weak acids partially dissociate in solution to release low [H+]
- HA + H2O ⇌ H3O+ + A-
how to name conjugate acids and bases?
HA + H2O -> H3O- + A-
- HA = acid
- H2O = base
- H3O- = conjugate acid
- A- = conjugate base
how to calculate pH or [H+]?
pH = -log[H+]
[H+] = 10 to the power of -pH
when calculating the pH of strong acids, what do you need to remember?
[H+] = HA because there is full dissociation
when calculating the pH of weak acids, what assumptions are made?
- initial conc. of acid = equilibrium conc. of acid
- @ equilibrium, [H+] = [A-]
how do you calculate ka (acid dissociation constant)?
[H+][A-]/[HA]
- when it is a weak acid [H+] = [A-]
what is ‘pka’ and how do you calculate it?
pka determines the strength of an acid
- pka = -log(ka)
- ka = 10 to the power of -pka
how do you use ka and pka to find strengths of weak acids?
- the higher the ka value, the stronger the weak acid because the equilibrium lies more to the right so there is a higher [H+]
- pka is the inverse of ka, so the higher the pka value the weaker the acid
what is the formula for kw (ionic product of water)?
kw = [H+][OH-]
what is the kw of water at standard temp. & pressure?
1x10 to the power of -14
units: mol²dm^-6
why does the pH of water change with temperature?
H2O ⇌ H+ + OH-
- when temperature increases, equilibrium position shifts to the endothermic side (forward reaction is favoured)
- [H+] increases, so pH decreases
why is water neutral at all pHs?
because [H+] = [OH-]
how is an acid buffer made?
- react excess acid with strong alkali e.g. NaOH
- all the NaOH is neutralised so the remaining acid and salt is left in the buffer
what is a buffer solution?
a solution that resists changes in pH when small amounts of acid and alkali are added
what does a buffer consist of?
a mixture of a weak acid and the salt of its conjugate base
how do buffer solutions work? use the symbol equation as an example:
CH3COOH ⇌ CH3COO- + H+
- **added OH- **-> pH will increase
- the added OH- will react with the H+ (H+ and OH- -> H2O)
- the equilibrium position will shift to the right so [H+] increases so pH decreases back down.
- added H+ -> pH will decrease
- equilibrium position will shift to the left so [H+] decreases, so pH increases back up.
how do you obtain a pH curve? (explain the calibration process)
- rinse pH probe w/ distilled water
- dip pH probe into a buffer solution
- wash pH probe in distilled water AGAIN
- dip pH probe in buffer solution w/ a different pH
- calibrate a pH probe and use to measure the initial pH of the alkali in the conical flask
how do you obtain a pH curve? (explain the steps post-calibration)
- add acid from the burette in 2cm³ increments and swirl the mixture
- record pH after every addition; reduce size of the portions of the increments close to the endpoint because the pH would change suddenly
- repeat until the acid is in excess
- plot a graph of pH vs. volume (of acid/alkali)
what does a strong acid + strong base pH curve look like? (e.g. HCl and NaOH)
- very large vertical section
- the pH at the equivalence point is 7 because there is full dissociation for both acid + alkali so [H+] = [OH-]
what does a strong acid + weak base pH curve look like? e.g. HCl and NH3
- relatively large vertical section
- pH at equivalence point is LESS than 7. acid has full dissociation and base has partial dissociation, so [H+] > [OH-]
what does a strong base + weak acid pH curve look like? e.g. NaOH + HCOOH
- equivalence point pH >7
- [OH-] = full dissociation; [H+] = partial dissociation
- vertical section is relatively small
