Chapter 20 - Acids, Bases and pH Flashcards
define acids and alkalis using the Arrhenius model
- acids dissociate and release H+ ions in aqueous solution
- alkalis dissociate and release OH- ions in aqueous solution
what is the Bronsted-Lowry definition of a base
a base is a proton acceptor
what is the Bronsted-Lowry definition of an acid
an acid is a proton donor
define a conjugate acid-base pair and give an example of one
“a conjugate acid-base pair contains two species that can be interconverted by the transfer of a proton”
HCl acid with Cl- base
give an example of a reaction with two acid base pairs, identify them
HCl + OH- —> H2O + Cl-
HCl and Cl- are an acid-base pair
H2O and OH- are an acid-base pair
define a monobasic acid
“monobasic acids are acids where one hydrogen atom can be replaced per molecule”
E.g. HCl
define a dibasic acid
“dibasic acids are acids where two hydrogen atoms can be replaced per molecule”
E.g. H2SO4
define a tribasic acid
“Tribasic acids are acids where three hydrogen atoms can be replaced per molecule”
E.g. H3PO4
define pH, give the equation
-log([H+])
the negative log (base 10) of the hydrogen ion concentration in a solution
what are the things to note about the pH scale e.g. high/low values etc.
high value = low H+ conc.
low value = high H+ conc.
its logarithmic so pH 3 is 10x higher H+ conc. than pH 2
what can we say about the H+ ion concentration of a strong acid
it is the same as the concentration of the acid because the acid fully dissociates
what happens to the pH of a solution if it is diluted by 10 times
- the pH increases by 1 because the H+ ion concentration decreases by 10 times
what is the difference between a strong acid and a weak acid
- a strong acid completely dissociates/ionises in aqueous solution
- a weak acid partially dissociates/ionises in aqueous solution
what is Ka
a version of Kc to represent the strength of a weak acid/how much it dissociates/where equilibrium lies
how can we calculate Ka for a weak acid HA
HA –Reversible–> A- + H+
Ka = [A-][H+] / [HA]
what is the only thing that changes Ka
- it acts in the same way as Kc, the only thing that changes it is temperature
what is pKa
like pH,
pKa = -log(Ka)
(base 10)
what can we say about the strength of a weak acid compared to its Ka and pKa value
- the stronger the acid, the higher the Ka value, the lower the pKa value
- the weaker the acid, the lower the Ka value, the higher the pKa value
what two factors does the H+ ion concentration depend on in a weak acid
- Ka
- [HA]
how can we layout an equilibrium for a weak acid in order to find its Ka
HA –reversible–> A- + H+
[HA]start, 0,0 = initial concs
[HA]eqm, [A-], [H+] = eqm concs
[HA} eqm = [HA]start - [H+]eqm
so
Ka = [H+][A-] / [HA]eqm
what is the simplified equation for Ka of a weak acid that allows us to calculate [H+] more easily
Ka = [H+]eqm^2 / [HA]start
what are the two assumptions we make that allow us to form the simplified equation for Ka of a weak acid
- we can assume that the dissociation of water is negligible and therefore there are no additional H+ ions so [H+] = [A-], allowing us to have x^2 on the top of the equation
- we can assume that the change in concentration of HA is so small that it is negligible, thus [HA]eqm = [HA]start
how can we find [H+] and therefore pH from a Ka and initial conc. of acid
Ka = [H+]eqm^2 / [HA]start as in the simplified equation rearrange to [H+] = sqrt(Ka x [HA]start) then pH as normal
where do the two assumptions we make break down
- we can no longer assume dissociation of water is negligible if, pH > 6, so in very weak or very dilute acids
- we can no longer assume [HA]start = [HA]eqm if Ka is larger ( Ka > 10^-2) or in very dilute solutions
what is a special property of water in terms of acids/bases
- it can act as an acid or a base making it amphotic
what is the equation for the dissociation of water
H2O –reversible–> OH- + H+
what is the Ka equation for water and what can we simplify it to
Ka = [OH-][H+] / [H2O]
this can be simplified to
Kw = [OH-][H+]
why can this simplification from Ka to Kw be made
we can assume that because there is such a large excess of water molecules when it dissociates that the concentration of water is a constant, (1000/18).
this means we can do
Kw = Ka[H2O] = [H+][OH-]
what is the value of Kw at 298K and how does this link to pH
Kw = 1 x 10^-14
at 298K,
[H+] = [OH-] = sqrt(Kw) = 1 x 10^-7
therefore pH of water is 7
when does Kw change
when temperature changes, like any other equilibrium constant
what is a neutral pH
depends on the temperature,
as temperature increases, the pH of neutral decreases, because there are more H+ ions per unit volume because equilibrium moves right
However, there are still equal numbers of H+ and OH- ions so it is still neutral
is the dissociation of an acid endothermic or exothermic
endothermic, hence why Kw increases when temperature increases
what is the equation for pKw
pKw = -log(Kw)
base 10
why is Kw very important for calculations
it acts as an equilibrium constant, if we know [H+] or [OH-] then we can calculate the other
explain how to calculate the pH of a strong base, and what to remember
1) use concentration of base to find [OH-], watch out for molar ratios!!
2) substitute into Kw to find [H+]
3) use pH formula
NOTE: make sure you use concentration and not moles
explain how to find the pH of an acid-base mixture
1) calculate number of moles of OH- and H+ from conc’s and vols
2) work out which is in excess, find the amount by which it is in excess
3) using this amount and the total volume, calculate the concentration of H+ or OH-
4) use Kw and [OH-] to find [H+] if its not in excess or just use H+ directly in the pH formula
even if the dissociation of the acid does make a difference to the equilibrium concentration of HA should we account for it
NO, mark schemes don’t like it
what is the name/term for Kw
the Ionic product of water
what is the ionic equation for a neutralisation, even if it’s a weak acid
OH- (aq) + H+ (aq) —> H2O (l)
if we have the conc of a weak acid and the conc of the H+ ions, how can we tell if it’s a strong or weak acid
- in a strong acid they would be equal
- in a weak acid they are different
what can we say about the concentrations of H+ and OH- ions in pure water at any temperature
they will be equal
it is only Kw = 10^-14 at 25 degrees celcius
what do we need to watch out for when using conc’s of strong bases and Kw to calc pH
some strong bases may release more than one OH- ion per molecule e.g. Ca(OH)2
you would need to double the conc. for use in Kw
