ions + solutions - acids/bases

Question 1

arrhenius acid definition

Answer 1

a substance that increases [H+] in solution

Question 2

arrhenius base definition

Answer 2

a substance that increases [OH-] in solution

Question 3

give 3 limitations of arrhenius acids/bases

Answer 3

• doesn’t consider non-aqeous solutions
• lacks knowledge of H+ = proton
• ammonia is unexplained, it behaves basically but has no OH-

Question 4

what model of acids/bases solved the limitations of arrhenius’ model?

Answer 4

bronsted-lowry model solves these issues
bronsted-lowry acid = proton donor
bronsted-lowry base = proton acceptor
acid HA has conjugate base A-
base B has conjugate acid HB+

Question 5

compare the ionisation of water from the arrhenius and brosnted perspectives

Answer 5

arrhenius proposed that water has non-zero conductivity as it self-ionises/auto-dissociates
H2O <–> OH- + H+ in aq solution

whereas bronsted and lowry suggested that this involved at least 2 water molecules
2H2O <–> H3O+ + OH- in aq solution

Question 6

amphoteric definition

Answer 6

when substances can behave as both acids + bases
e.g. water as an acid has conjugate base OH- and as a base has conjugate acid H3O+

Question 7

what is the idea behind the grotthuss mechanism?

Answer 7

protons are much more mobile in water than other ions as it doesn’t have to be the same proton - proton swapping/hopping can occur rapidly along a chain of water molecules
OH is highly mobile for the same reason
- remember than mobile ions contribute a conductivity proportional to their conentration, OH-/H+ ions have a much higher conductivity due to the grotthuss mechanism

Question 8

show how pH scale is based on equilibirum of water + kw

Answer 8

2H2O <–> H3O+ + OH- gives
k = [H3O+][OH] / [H2O]^2
however typically [H3O+] and [OH-] are so dilute than [H2O] is constant, they are absorbed into expression
assuming ideal solution behaviour ([] instead of {}):
kw = [H3O+][OH-] and at 25C = 1x10^-14
so solving for neutral water at 25C has [H+] ~ 1x10^-7 and [H+] = [OH-]
- pH is based on this

Question 9

what environmental conditions affect kw + how?

Answer 9

kw is temperature + pressure dependent, so water at other temps can be neutral but [H+] will not = 1x10^-7

Question 10

neutral definition

Answer 10

when [H+ = [OH-], regardless of pH

Question 11

how is pH affected by ideality?

Answer 11

if assuming ideal solution, pH is based on [H+]
if not ideal, pH is based on {H+}

Question 12

strong acid definition

Answer 12

an acid that fully dissociated in aq solution into solvated protons + its conjugate base - therefore pH can be estimated from molarity of solution

Question 13

solvent levelling definition

Answer 13

when a solvent limits the strength of acids/bases within it, with by levelling it eith their own acidic/basic properties
- why very large amounts of acid are needed to move pH to <0, as the pH scale is logarithmic

Question 14

what is the difference between pH and pOH?

Answer 14

pH is based on [H+]
pOH is the -ve base 10 logarithm of [OH-] - this is useful for discussing strong bases

Question 15

how can pH/pOH be related from the kw equilibrium?

Answer 15

kw = [H3O+][OH-] and at 25C = 1x10^-14
-logkw = pOH + pH = 14 at 25 C
therefore, pH = approx 14-pOH

Question 16

how can metal-aqua complexes display acidic behaviour?

Answer 16

proton exchange can convert a water molecule in an aqua complex into an OH-, liberating H+ into the bulk water, therefore causing solutions of metal salts to become acidic

Question 17

weak electrolyte definition

Answer 17

substances that do not dissociate completely in solution

Question 18

give the equilibrium equation for a dissociation reaction

Answer 18

• it is important to remember that there is an equilibrium between dissociated and associated forms in solution
  Ka = [X][Y] / [XY] - if assuming ideality
  Ka = association constant

Question 19

what environmental factors does Ka depend on?

Answer 19

only temperature - boltzmann

Question 20

how can Ka be used to describe acidity?

Answer 20

using a similar base 10 -ve logarithmic scale:
pKa = -logKa
so strong acids have very small pKas

Question 21

show how the degree of dissociation calculation is derived from Ka

Answer 21

Ka = [H][A] / [HA]
= (xCo)(xCo) / (1-x)(Co)
= ( x^2 / 1-x ) * Co

where Co = concentration of weak acid
and x = fraction that dissociates
- this is simplified by assuming there is enough acid present to ignore kw

Question 22

why does the pH of water change if it is allowed to stand open to the air?

Answer 22

CO2 in atmosphere is very soluble, dissolving in it to form carbonic acid then dissociating into bicarbonate

Question 23

what is Kb?

Answer 23

the basicity constant = Kb
this is found from the equilibrium of bases with water:
B + H2O <–> BH+ + OH-
H2O is absorbed therefore assumed constant
similarly to pKa, can -ve log Kb to get pKb, which can be used to find [OH], pOH, and pH

Question 24

relate pKa and pKb using kw equation

Answer 24

pKa + pKb = pKw = 14 at 25 C

Question 25

buffer solution defintion

Answer 25

weak acid + conjugate base mixture, able to resist small changes in pH

Question 26

how does dissociation of buffers compare to acids?

Answer 26

degree of dissociation will be less than the acid alone (used to create the buffer solution)

Question 27

derive the equation for the pH of a buffer

Answer 27

Ka = [H][A]/[HA] => -pKa = log[H]+log[A] - log[HA] => -pKa = -pH + log([A]/[HA]) this is the hendersson-hasselbach equation, and it works as long as both [A] and [HA] are present in large quantity

Question 28

considering the hendersson-hasselbach equation, what does it mean if acid + base are equimolar?

Answer 28

acid being [HA] and conjugate base being [A] when these are equal, pH = pKa