Equilibria in Aqueous Systems Flashcards
(34 cards)
what is a common ion
A common ion is an ion that is present in a solution from two different sources
NaCl and HCl = common ion of Cl because it is produced but from different sources
common ion effect
it is a reduction of dissociation in weak acid/base due to a common ion already in a solution
- the presence of a common ion shifts the equilibrium towards reactants because le chantelier principle
CH3COOH –> CH3COO- + H+
CH3COONa –> Na+ + CH3COO-
this H+ reacts with the common ion to reform the acid - buffer
its a weak acid and the salt it produces - this forms a buffer system
Acid-Base Buffer
A solution that resists changes in pH when a small amount of either strong acid or strong base is added
weak acid + conjugate base OR weak base + conjugate acid
buffer ratio = [HA]/[A]
explain the buffer ratio
[HA]/[A] is large, which means there is more acid and less conj base –> more hydronium, more acidity
[HA]/[A] is small, which means there is more conj base than acid –> less hydronium, less acidity, more basic
what happens when you add a strong acid to a buffer system
a strong acid reacts with a weak base (A)
A- + H3O+ –> HA + H2O (reverse of acid dissociation)
HA conc increases, A conc decreases
= pH decreases, making it more acidic if HA > A
what happens when you add a strong base to a buffer system
it will react with the acidic component of the buffer (HA)
HA + OH- –> A- + H2O
[HA] decreases, [A] increases because it is being produced
= HA>A
= pH increases, more basic
the less the difference between HA and A in the buffer concentration ratio, the better/worse the buffer is
better
you want close concs of the weak acid and the conjugate base to = more effective in resisting pH changes
what is henderson-hasselbalch used for
pH = pKa + log ([A]/[HA])
helps us solve for pH directly without H3O
when the HA = A the log part becomes log1 which = 0
this leaves you with pH=pKa
usually, pKa is given in the equation or it can be solved by -log (Ka)
*the buffer is the most effective when HA = A
explain buffer capacity
the measure of the strength of the buffer - its ability to maintain the pH fluctuations following the addition of strong acids or strong bases
- depends only on concentration
- works best when [A] and [HA] are very close in magnitude + high in concentration
- if one component dominates, the buffer becomes less effective
GREAT CAPACITY = less pH flucuations
what is the absolute concentration in buffer systems
the more concentrated the buffer components - A and HA are, the greater the capacity
- increasing the capacity, increases the effectiveness of the buffer - less fluctuations
what are the relative concentrations in buffer systems
the closer the component concentrations are to each other, the greater the capacity = less fluctuations
what is the buffer range
the pH range over which the buffer is effective and related to the relative concentrations
The more differences the acid and base components are the less effective the buffer = lower buffer capacity (further away from 1)
If the [A]/[HA] component is bigger than 10 or less than 0.1 = poor buffer action
*buffers can have a usable range within +/- 1 pH unit of the pKa
Briefly how to prepare a buffer
- choose the conjugate acid-base pair
- calculate the ratio of the buffer concentrations
- determine the buffer concentration
- mix the solution and correct the pH
the acid-base titration curve
a plot of pH VS the volume of titrant added (in the burette)
what happens when you titrate a strong acid and a strong base together
A: pH = acidic with lots of H3O, the pH starts to increase and become a bit more basic once the base is constantly added - H3O conc slowly decreases
B: the pH rises very rapidly
- when the amount of dissociated OH from the base is ALMOST equal to the amount of H3O remaining
C: equivalence point, when the moles of OH = moles of H3O originally present
- anions of strong acid and cations of a strong base
- these anions and cations - salt - do not further react
D: pH continuously rises slowly to increase the pH, making it more basic as OH > H3O
where is the equivalence point for strong acids and bases relatively near?
ph of 7
explain the weak acid + strong base titration curve
A: pH starts a bit higher, not as acidic as the strong acid, starts with less H3O
B: curve rises gradually (buffer-region, half-way point)
- HPr reacts with the strong base to produce Pr
- half of the HPr is equal to the amount of Pr made (pH = pKa)
C: equivalence point - where the solutions have the same amounts of [HPr] = [Pr]
D: the pH increases slowly beyond the equivalence point as excess OH is added
where is the equivalence point for weak acids and strong bases relatively near?
above a pH of 7
shifted upwards
starting pH is higher and EP is higher
explain weak base and strong acid titration curves
displays the same shape as the weak acid curve but inverted so it starts with a basic pH
A: starts with low amounts of OH, and a relatively high pH
B: buffer region zone - where pH = pKa. this is when the HB = B (weak base ALMOST = conj acid)
C: equivalence point, shifted
- where HB = exactly B and the concentrations of H3O and OH are equal
D: excess H3O from the addition of more strong acid, this makes the pH very low
where is the equivalence point for weak base and strong acids relatively near?
pH is shifted downwards below 7
what is an indicator and how does it work with monitoring pH
a weak organic acid whose colour (A) differs from the colour (B) of its conjugate base because the 2 species have slightly different structures
When the pH is lower = H3O is higher = [Hln]»_space; [ln]
As the pH increases = H3O decreases [Hln] «_space;[ln]
this ratio = Henderson-hasselbalch [ln]/[Hln]
When the ratio = 0.1 the acid colour is seen (1:10)
When the ratio = 1 the acid and base colours (1:1)merge to form an intermediate hue
When the ratio = 10 the base colour is seen (10:1)
how do you choose the acid-base indicator
the endpoint occurs when the indicator changes colour
- you choose an indicator with a colour change that is close to the pH of the EP - you want to see the end point happening via bright colour change
what is the Ksp
solubility product constant
- equilibrium constant for the dissolution of a slightly soluble compound
Solid Iron + H2O —> P1 + P2
Ksp expressions don’t have denominators (products/reactants) because the reactants are pure solids/water
at equilibrium = saturation
why can’t you always assume a salt completely dissociates
some compounds partly dissociate and undissociated, this increases the solubility
