5.1.3 - Acid, Bases and Buffers Flashcards
define bronsted-lowry acid
a species that donates a proton
define brosted-lowry base
a species that accepts a proton
what happens in terms of electrons in a brosted lowry base
lone pair of electrons that form a dative covalent bond with H
what is an amopheric substance
a substance that can act as either an acid or a base depending on the condition
example of an amopheric substance
water
monobasic acids
an acid that can donate 1 proton
example of a monobasic acid
HCl and CH3COOH
dibasic acids
acids with 2 protons to donate
example of a dibasic acid
H2SO4 and H2CO3
tribasic acids
acids with 3 protons to donate
example of a tribasic acid
H3BO3
ionic equation for: acid + metal
2H+(aq) + metal(s) –> metal ion2+(aq) + H2(g)
ionic equation for: acid + carbonate
2H+(aq) + CO3 2+(aq) –> H2O(l) + CO2(g)
ionic equation for: acid + metal oxide
2H+(aq) + metal oxide(s) –> metal ion2+(aq) + H2O(aq)
ionic equation for: acid + alkali
H+(aq) + OH-(aq) –> H2O(l)
what is the acid dissociation constant
Ka
equation for Ka
Ka = [H+][A-]/[HA]
assumptions when calulation things for a weak acid using Ka
- when HA dissociates, [H] = [A]
- [HA]eq = [HA]start
problems with the assumption that [HA]eq = [HA]start
It becomes less accurate the bigger Ka becomes
what does the size of Ka tell us about an acid
The larger the Ka, the stronger the acid and the higher the H+ concentration at equilibrium
formula for pH
pH = -log[H+]
how to find [H+] from pH
[H+] = 10^-pH
formula for pKa
pKa = -log(Ka)
how to find Ka from pKa
Ka = 10^-pKa
what does ‘p’ imply
negative logarithmic scale
what is Kw
ionic product of water:
H2O(l) ⇌ H+(aq) + OH-(aq)
equation for Kw
Kw = [H+][OH-]
why can [H2O] be ignored when calculating Kw
it is constant at a given temperature
what is Kw at room temperature (give room temperature in K)
1x10^-14
room temperature: 298K
assumptions when calculating Kw
in a neutral substance [H+] = [OH]
therefore Kw = [H+]^2
calculating the pH of strong bases using Kw
use [OH-] and Kw to find [H+]
assumptions when calculating the pH of strong bases
- fully ionise in water eg. NaOH –> Na+ + OH-
- assume [NaOH] = [Na+] = [OH-]
- [base] = [OH-]
define buffer
a solution that resists pH changes when small amounts of acid or base are added, or when diluted
what are buffers formed from, give eg.
weak acid + salt of weak acid
ethanoic acid + sodium ethanoate
why do we assume all ethanoate ions come from sodium ethanoate in a buffer
CH3COOH ⇌ CH3COO- + H+
CH3COONa —> CH3COO- + Na+
equilibrium for dissociation of ethanoic acid is so far left its negligable
therefore [CH3COONa] = [CH3COO-]
what happens when you add an acid to a buffer
pH decreases
increased [H+]
H+ reacts with CH3COO-
as H+ is removed, equilibrium shifts left and pH increases
what happens when you add alkali to a buffer
pH increases
increased [OH-]
OH- reacts with H+
[H+] decreases so equilibrium shifts right to increase [H+] again
decreasing pH
calculating pH of buffers
Ka = [H+][A-]/[HA]
why cant we assume [H+] = [A-] when calculating the pH of buffers
[A-] from the base
equation for the blood buffer system
H2CO3 ⇌ H+ + HCO3-
function of the blood buffer system
to maintain the pH of blood between 7.35 and 7.45
acidocis
when blood pH falls below 7.35
- fatigue, shock and death
alkalosis
when blood pH falls above 7.45
- nausea, muscle spasms, headaches
what is the equivalence point on a pH curve
when the volume of acid reacts exactly with the volume of alkali in the amounts matching the stoichiometry
how to match an indicator to a pH curve
the verical section of the curve to match the pH range of the indicator
