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A Level Chemistry > 5.1.3 Rates of Reactions- acids and bases > Flashcards

5.1.3 Rates of Reactions- acids and bases Flashcards

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1
Q

bronsted-lowry acids

A
  • acids dissociate and release H+ ions in aqueous solution

- proton donor

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2
Q

bronsted-lowry bases

A

proton acceptor

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3
Q

alkalis

A

disscoate and release OH- ions in aqueous solution

a soluble base

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4
Q

conjugate acid-base pair

A

a conjugate acid is the species formed when a base accepts a proton. a conjugate base is the species formed when an acid donates a proton.
they form conjugate acid-base pairs
e.g. HCl and Cl-
-in forward direction, HCl releases a proton to form its conjugate base, Cl-
-in the reverse direction, Cl- accepts a proton to form its conjugate acid, HCl

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5
Q

hydronium ion, H3O+

A

HCl + H2O <=> H3O+ + Cl-

  • H2O has accepted a proton to form its conjugate acid, the hydronium io,
  • H3O+ is very important, as it is the active acid ingredient in any aqueous acid
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6
Q

monobasic, dibasic and tribasic acids

A

refer to the total number of hydrogen ions in the acid that can be replaced per molecule in an acid-base reaction
monobasic- HCl and CH3COOH
dibasic- H2CO3
tribasic- H3BO3

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7
Q

redox reactions between acids and metals

A

dilute acids undergo redox reactions with some metals to produce salts and hydrogen gas
acid + metal –> salt + hydrogen

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8
Q

neutralisation of acids with carbonates

A

carbonates are bases that neutralises acids to form a salt, water and CO2

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9
Q

neutralisation of acids with metal oxides

A

an acid is neutralised by a solid metal oxide or hydroxide to form a salt and water only

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10
Q

neutralisation of acids with alkalis

A

with alkalis, the acid and base are in solution. the overall reaction is the same as for metal oxides
acid + alkali –> salt + water

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11
Q

pH scale

A

a logarithmic scale

  • a low value of [H+(aq)] matches a high value of pH
  • a high value of [H+(aq)] matches a low value of pH
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12
Q

mathematical relationship between pH and [H+(aq)]

A

pH = -log[H+(aq)]
reverse is [H+(aq)] = 10^-pH
-a change in one pH number is equal to a 10 times difference in [H+(aq)]
-to dilute a solution from a pH 1 to pH 4 would require dilution by 10x10x10= 1000 times

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13
Q

calculating pH of strong acids

A

in aqueous solution, a strong monobasic acid, HA, completely dissociates:
HA –> H+ + A-
1 mol 1 mol
so, for a strong acid, [H+(aq)] is equal to the conc of the acid, [HA(aq)]:
[H+(aq)]=[HA(aq)]

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14
Q

strong acid

A

HCl

completely dissociates in aqueous solution

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15
Q

weak acid

A

CH3COOH

partially dissociates in aqueous solution

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16
Q

acid dissociation constant Ka

A

one of several equilibrium constants used for acid-base equilibria
Ka= [H+(aq)][A(aq)]/[HA(aq)]
Ka changes with temp and recorded Ka values are usually standardised at 25oC
the larger the numerical value of Ka, the further the equilibrium is to the right
the larger the Ka value, the greater the dissociation and the greater the acid strength

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17
Q

Ka and pKa

A

The pKa value is one method used to indicate the strength of an acid. pKa is the negative log of the acid dissociation constant or Ka value. A lower pKa value indicates a stronger acid. That is, the lower value indicates the acid more fully dissociates in water.
pKa = -logKa
Ka = 10^-pKa

18
Q

pH of weak acids

A

monobasic acid HA:
a weak acid HA partially dissociates in aq solution
HA <==> H+ + A-
[H+(aq)] depends upon:
-con cof the acid [HA(aq)]
-acid dissociation constant Ka
Ka= [H+(aq)]eqm [A-(aq)]eqm/[HA(aq)]eqm = [H+(aq)]eqm [A-(aq)]eqm / [HA(aq)]start - [H+(aq)]eqm

[H+(aq)] = square root of Ka x [HA(aq)]

19
Q

approximations for pH of weak acids equation

A

-HA dissociates to produce equilibrium concs of H+ and A- that are equal
- [HA(aq)]start&raquo_space; [H+(aq)]eqm
[HA(aq)]eqm = [HA(aq)]start - [H+(aq)]eqm approximates to [HA]eqm=[HA]start

20
Q

determination of Ka for a weak acid method

A

-prepare a standard solution of the weak acid of known conc
-measuring the pH of the standard solution using a pH meter
Ka= -log(pH)^2 / conc

21
Q

approximations in calculations involving weak acids

A
  • assumes that the dissociation of water is negligible

- assumes that the conc of acid is much greater than the H+ conc at eqm

22
Q

ionisation of water

A

water ionises very slightly, acting as both an acid and a base
H2O + H2O <=> H3O+ + OH-

23
Q

Kw

A

ionic product of water - the ions in water (H+ and OH-) multiplied together
Kw= [H+(aq)][OH-(aq)]

24
Q

importance of Kw

A

significance of Kw having value of 1.00 x 10-14 mol2dm-6 at 25oC is huge
the value sets up the neutral point in the pH scale
Kw controls the conc of H+ and OH- ions in aq solutions

25
Q

pH of strong bases

A

calculated from conc of base and ionic product of water Kw

[H+(aq)] = Kw/[OH-(aq)]

26
Q

buffer solution

A

a system that minimises pH changes when small amounts of an acid or a base are added
contain two components to remove added acid or alkali: a weak acid and its conjugate base

27
Q

how does a buffer solution work

A

when alkalis and acids are added to a buffer, the two components in the buffer solution react and will eventually be used up
as soon as one component has all reacted, the solution loses its buffering ability towards added acids or alkalis
as the buffer works, the pH does change but only by a small amount

28
Q

formation of a buffer solution from a weak acid and its salt

A

a buffer solution can be prepared by mixing a solution of ethanoic acid, CH3COOH, with a solution of one of its salts, e.g. sodium ethanoate CH3COONa
when ethanoic acid is added to water, the acid partially dissociates and the amount of ethanoate ions in solution is very small. ethanoic acid is the source of the weak acid component of the buffer solution
CH3COOH <=> H+ + CH3COO-
salts of weak acids are ionic compounds and provide a convenient source of the conjugate base. when added to water, the salt completely dissolves. dissociation into ions is complete and so the salt is the source of the conjugate base component of the buffer solution
CH3COONa + aq –> CH3COO- + Na+

29
Q

formation of a buffer solution from the partial neutralisation of the weak acid

A

a buffer solution can be prepared by adding an aqueous solution of an alkali, such as NaOH, to an excess of the weak acid. the weak acid is partially neutralised by the alkali, forming the conjugate base. some of the weak acid is left over unreacted
the resulting solution contains a mix of the salt of the weak acid and any unreacted weak acid

30
Q

two reservoirs to remove added acid and alkali

A

in the ethanoic equilibrium, the equilibrium position lies well towards ethanoic acid
when CH3COO- ions are added to CH3COOH, the equilibrium position shifts even further to the left, reducing the already small conc of H+ ions, and leaving a solution containing mainly the two components, CH3COOH and CH3COO-
CH3COOH <=> H+ + CH3COO-
CH3COOH and CH3COO- act as two reservoirs that are able to act independently to remove added acid and alkali
this is achieved by shifting the buffer’s equilibrium system either to the right or left

31
Q

action of the buffer solution- conjugate base

A

the conjugate base removes added acid
on addition of an acid, H+:
1 [H+(aq)] increases
2 H+ ions react with the conjugate base, A-
3 the equilibrium position shifts to the left, removing most of the H+ ions

32
Q

action of the buffer solution- weak acid

A

the weak acid removes added alkali
on addition of an alkali, OH-:
1 [OH-(aq)] increases
2 the small conc of H+ ions reacts with the OH- ions:
H+ + OH- –> H2O
3 HA dissociates, shifting the equilibrium position to the right to restore most of H+ ions

33
Q

calculating the pH of a buffer solution

A

the equilibrium and Ka expression for a weak acid is:
HA <=> H+ + A- Ka= [H+(aq)] [A-(aq)] / [HA(aq)]
when you calculate the pH for a weak acid, you make an approximation that [H+(aq)] = [A-(aq)]
for a buffer solution, this is no longer true as A- has been added as one of the components of the buffer

[H+(aq)] = Ka x [HA(aq)]/[A-(aq)]
Ka= [H+(aq)] and pKa=pH
34
Q

carbonic acid-hydrogencarbonate buffer system

A

a general buffer system based on HA/A- controls pH. the carbonic acid-hydrogencarbonate operates in a similar way:
on addition of an acid, H+:
1 [H+(aq)] increases
2 H+ ions react with the conjugate base, HCO3-
3 the equilibrium position shifts to the left, removing most of the H+ ions

on addition of an alkali, OH-:
1 [OH-(aq)] increases
2 the small conc of H+ ions reacts with the OH- ions:
H+ + OH- –> H2O
3 H2CO3 dissociates, shifting the equilibrium position to the right to restore most of H+ ions

the body produces far more acidic material than alkaline, which the conjugate base HCO3- converts to H2CO3. the body prevents H2CO3 building up by converting it to CO2, which is then exhaled by the lungs

35
Q

pH titration curves

A

when the base is first added, the acid is in great excess and the pH increases very slightly. as the vertical section is approached, the pH starts to increase more quickly as the acid is used up more quickly
eventually, the pH increases rapidly during addition of a very small volume of base, producing the vertical section. only drops of solution will be needed for the whole vertical section
after the vertical, the pH will rise very slightly as the base is now in great excess.
the equivalence point of the titration is the volume of one solution that exactly reacts with the volume of the other solution
the equivalence point is the centre of the vertical section of the pH titration curve

36
Q

pH meter

A

1 using a pipette, add a measured volume of acid to a conical flask
2 place the electrode of the pH meter in the flask
3 add the aqueous base to the burette and add to the acid in the conical flask, 1cm3 at a time
4 after each addition, swirl the contents. record the pH and the total volume of the aqueous base added
5 repeat steps 3 and 4 until the pH starts to change more rapidly. then add the aqueous base dropwise for each reading until the pH changes less rapidly
6 now add the aqueous base 1cm3 at a time again until an excess has been added and the pH has been basic, with little change, for several additions

a graph of pH against total volume of aq base added is then plotted

37
Q

acid-base indicators - the end point

A

an acid-base indicator is a weak acid, HA, that has a distinctively different colour from its conjugate base, A-, e.g. the common indicator methyl orange
-the weak acid is red
-the conjugate base is yellow
at the end point of a titration, the indicator contains equal concs of HA and A- and the colour will be in between the two extreme colours

38
Q

explain indicator colour changes

A

an indicator is a weak acid
the equilibrium position is shifted towards the weak acid in acidic conditions or towards the conjugate base in basic conditions, changing the colour as it does so
in a titration in which a strong base is added to a strong acid, methyl orange is initially red as the presence of H+ ions forces the equilibrium position well to the left
on addition of a basic solution containing OH- ions:
-OH- ions react with H+ in the indicator
-the weak acid HA dissociates, shifting the equilibrium to the right.
-the colour changes, first to orange at the end point and finally to yellow as the equilibrium position is shifted to the right.

if methyl orange is added initially to a basic solution and an acid is added:

  • H+ ions react with the conjugate base A-
  • the equilibrium position shifts to the left
  • the colour changes, first to orange at the end point and finally to red when the equilibrium position has shifted to the left
39
Q

how sensitive is the end point?

A

the pH of the end point is the same as the pKa value of HA

the sensitivity of an indicator depends upon the indicator itself and eyesight

40
Q

choosing an indicator

A

in a titration you must usa an indicator that has a colour change which coincides with the vertical section of the pH titration curve
ideally the end point and equivalence point would coincide

A Level Chemistry (16 decks)

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