Chapter 20: Acids, Bases, and pH (20.1-20.5) Flashcards
what is a bronsted-lowry acid?
a H+/proton donor
what is a bronsted-lowry base?
a H+/proton acceptor
the general equation for acid dissociation
HA < – > H+ + A-
A- is the conjugate base
Define a conjugate acid-base pair
a conjugate acid-base pair contains 2 species that can be interconverted by transfer of a proton
(eg HNO3 and NO3- can be interconverted by transfer of a proton/H+ ion)
define monobasic acid
has 1 proton (H+) that can be replaced in an acid-base reaction
define dibasic acid
has 2 protons (H+) that can be replaced in an acid-base reaction
define tribasic acid
has 3 protons (H+) that can be replaced in an acid-base reaction
hydronium ion
H3O+
what is a salt
a salt is the product of a neutralisation reaction with an acid and a base where the H+ ion has been replaced with either a metal or ammonium (NH4+)
a strong acid
releases all of its H+ ions in aqueous solution
HA –> H+ + A-
example: HCl, H2SO4
weak acid
does not release all of its H+ ions in aqueous solution
HA < – > H+ + A-
example: all organic acids eg CH3COOH
equation for pH
pH = -log[H+]
use base 10
acid dissociation constant or Ka
Ka = [H+][A-] / [HA] or = [H+]^2 / [HA]
HA < – > H+ + A- (Ka is only for weak acids)
pKa =
-log(Ka)
base 10
dilute acid + metal –>
redox
salt + H2
eg 2H+(aq) + Zn(s) –> Zn 2+(aq) + H2(g)
common bases
carbonates, metal oxides, metal hydroxides, alkalis
acid + carbonate –>
neutralisation
salt + water + CO2
eg 2H+(aq) + CaCO3(s) –> Ca 2+(aq) + H2O(l) + CO2(g)
acid + metal oxide (s) –>
acid + metal hydroxide (s) –>
(neutralisation, they are the same)
salt + water
acid + alkali –>
neutralisation
H2O
eg H+(aq) + OH-(aq) –> H2O (l)
[H+] =
[H+] = 10^(-pH)
larger acid dissociation (Ka) means
stronger acid
a larger pKa means
a weaker acid (a smaller Ka)
neutralisation is
a chemical reaction in which an acid and a base reacts to form water
in strong acids: [H+] =
[H+] = [HA]
weak acid:
[HA] (at equilibrium) =
And why does this not necessarily work for ‘stronger’ weak acids?
[HA(aq)] (start)»_space; [H+(aq)] (equilibrium)
HA(aq) = [HA(aq)] (start) - [H+(aq)] (eqm)
[HA] (at eqm) ~ [HA] (at start)
Breaks down when [H+] becomes significant and there is a real difference between
[HA(aq)] (equilibrium) and [HA(aq)] (start) - [H+(aq)] (equilibrium)
what is Kw (definition)
the ionic product of water
H2O dissosiates slightly, acting as a weak acid
H2O < – > H+ + OH-
what is Kw (equation)
Kw = [H+] X [OH-]
you are given the value of Kw
what is Kw (number value)
Kw (at rtp) = 1 X 10^-14
if [H+] = [OH-]
solution is neutral
if [H+] > [OH-]
solution is acidic
if [H+] < [OH-]
solution is alkali
strong bases
completely dissociate in aqueous solution
KOH –> K+ + OH-
how to find the pH of a strong base
[KOH] = [OH-]
then use Kw = [H+][OH-] to find [H+]
use pH = -log[H+]
Why do we use pH rather than [H+]?
pH makes the numbers more manageable, as it takes a log of the really small [H+] values
Weak acid approximation for [H+] at equilibrium
[H+] (equilibrium) ~ [A-] (equilibrium]
i.e there is a negligible dissociation of water (weak solutions would cause there to be more H+ due to its dissociation in water)
