acids and bases Flashcards
What is an acid base equilibira?
A reaction involving the transfer of protons
Define a Bronsted Lowry acid
A proton donor
Eg NH4+
Define a Bronsted Lowry base
A proton acceptor
Eg OH-
Define concentration
how many moles are in a given volume of solution
Define acid strength
The strength is a measure of how ionised the acid is in solution
What’s the difference between a strong acid and a weak acid?
A strong acid will fully ionise in solution whereas a weak acid will partially ionise in solution.
What is pH
a measure of the concentration of hydrogen ions in solution.
As the concentration of hydrogen ions for a strong acid is very high, yet the concentration of hydrogen ions for a strong alkali is very low a logarithmic scale must be used.
So,
pH = -log10[H+].
equation to determine [H+] with ph
[H+] = 10^–pH
Write an equation for the dissossiation of sulfuric acid in water
H2SO4(l) → 2H+(aq) + SO42–(aq)
Diprotic acid
What causes a solution to be acidic?
H+ (hydrogen ion), or more accurately H3O+ (hydroxonium ion), protons react with H2O to form it.
Explain the dissociation of water
Water very slightly dissociates into its ions according to the following equation:
2H2O(l) ⇋ H3O+(aq) + OH-(aq)
Write an equation for the ionisation of water
2H2O ⇌ H3O+ + OH
//
H2O ⇌ H++ OH-
Explain the nature of the formation of hydroxonium and OH- ions
hydroxonium ion is a very strong acid, and the hydroxide ion is a very strong base. As fast as they are formed, they react to produce water again.
The net effect is that an equilibrium is set up.
2H2O(l) ⇋ H3O+(aq) + OH-(aq)
At any one time, there are incredibly small numbers of hydroxonium ions and hydroxide ions present. (The H2O is in large excess)
Water very slightly dissociates
Derive Kw using the equation for the ionisation of water
H2O ⇌ H++ OH-
Keq = [H+] [OH-]/ [H2O]
Keq [H2O] = [H+][OH-]
[H2O] is so large compared to [H+] and [OH-] that Keq[H2O] can be considered to be constant
Keq[H2O] = Kw
Therefore, Kw = [H+][OH-]
What is the difference between ionising and dissassociating?
Just an NB
Ionisation is the process that involves the formation of ions whereas dissociation is the process of breaking up of a moiety into its constituent atoms, molecules and ions
What ius the value of Kw at room temperature (25 *C)
1 x 10 ^-14
What is the forward reaction of the ionisation of water
endothermic
Is the forward or backward reaction favoured when temperature of water is increased?
Forward, because its endothermic.. so more [H+] is produced so becomes more acidic, pH decreases
What physical factors affect the value of Kw? How do they affect it?
Temperature only- if temperature is increased, the equibilibium moves to the right so Kw increases and the pH of opure water decreases.
Why is pure water still neutral, even if pH does not equal 7?
[H+]=[OH-]
What is the realtionship between pH and conc of H+?
Lower pH= higher concentration of H+
If two solutiuons have a pH difference of 1, what is the difference in [H+]?
A factor of 10
Why does pure water have a pH of 7?
That question is actually misleading! In fact, pure water only has a pH of 7 at a particular temperature - the temperature at which the Kw value is 1.00 x 10-14 mol2 dm-6.
In pure water at room temperature the Kw value tells you that:
[H+] [OH-] = 1.00 x 10-14
But in pure water, the hydrogen ion (hydroxonium ion) concentration must be equal to the hydroxide ion concentration. For every hydrogen ion formed, there is a hydroxide ion formed as well.
That means that you can replace the [OH-] term in the Kw expression by another [H+].
[H+]2 = 1.00 x 10-14
Taking the square root of each side gives:
[H+] = 1.00 x 10-7 mol dm-3
Converting that into pH:
pH = - log10 [H+]
pH = 7
How would you calculate pH of 0.0500 mol dm-3 Ba(OH)2
[OH-]= 0.05 x 2
[H+] = kw/ [OH-]
[H+] = 10^-14/ 0.1
pH = -log [H+]
pH= -log [1.00 x 10^-13]
pH = 13.00
How would you calculate [H+] of [Ba(OH)2] with pH 13.30?
[H+] = 10 ^-pH
[H+] = 10 ^-13.3 = 5.01x10^-14
[OH-] = Kw/[H+]
[OH-]=10^-14/ 5.01x10^-14 = 0.2
[Ba(OH)2] = 0.2 x 1/2 = 0.100 mold dm^-3
Give examples of strong monoprotic acids
HCl, HNO3
Give examples of strong diprotic acids
H2SO4
Give examples of strong triprotic acids
H3PO4
Give examples of strong monobasic bases
NaOH KOH
Give examples of strong diprotic bases
Ca(OH)2
Ba(OH)2
Give examples of weak monobasic bases
ammonia
methylamine
Give examples of strong dirpotic bases
1,2-diaminoethane
Give examples of weak acids
mono - ethanoic
di - ethandioic
tri - citiric
What is Ka
acid dissociation constant
Weak acids and bases only dissocaite partially in water therefore the reaction has an equilibirum dissassociation constant Ka
Explain ka
We define a new constant as Ka =
Ka = [H+][A-]/[HA]
HA refers to the un-ionised acid molecule, and A- refers to the anion remainder after loss of a proton.
- A large Ka value means that a lot of the acid ionises in solution. (stronger acid)
- A small Ka value means that very little acid ionises in solution. (weaker acid)
pKa=
pKa = – log Ka
the stronger the acid:
the bigger Ka
the smaller pKa
Basic titration method
Standard solution of acid/alkali prepared in volumetric flask
25 cm3 samples taken out with pipette and put into conical flask
Indicator + acid/alkali in conical flask
acid/alkali added from burette until permanent colour change
How do you make a standard solution RP1
- mass a weighing bottle/boat
- add the correct mass of solid
- pour the contents of the solid into a beaker
- mass the weighing bottle/boat again (the difference in two recorded masses will provide the exact
mass of the solid added to the beaker) - dissolve the solid in a little distilled water using a stirring rod
- pour the solution into a volumetric flask using a funnel and stirring rod
- wash the beaker several times, pouring the washings into the flask
- top up to the mark carefully (drop-wise at the end)
- swirl, invert and repeatedly shake to ensure it is homogenous (thoroughly mixed)
What is a back titration?
Used where reaction between acid and base is slow (e.g. low solubility in water)
React an excess of an acid/base with the acid/base – some is leftover
Titrate the leftover acid/base to see how much leftover
What is an indicator?
Weak acids which have a different colour to their conjugate bases.
HIn ⇌ H+ + In-
colour 1 ⇌ colour 2
How do indicators work?
HIn ⇌ H+ + In-
low pH: equilibrium pushed left = colour 1
high pH: equilibrium pushed right = colour 2
methyl orange
colour at low ph
range
colour at high pH
red
3.2-4.4
organge/yellow
phenolphthalein
colour at low ph
range
colour at high pH
colourless
8.2-10.0
purple
What is a titration curve
shows how pH of a solution changes in an acid-base reaction
How is the neutralisation point identified
a large vertical section therough the equivalence point
Define equivalence point
The point at which the exact volue of base has been added to just neutralise the acid or vice versa.
strong acid-strong base
pH at equivalence:
suitable indicator:
7
phenolphthalien/methyl orange
strong acid-weak base
pH at equivalence:
suitable indicator:
< 7 (4 )
methyl orange
weak acid-strong base
pH at equivalence:
suitable indicator:
> 7 (9 )
phenolphthalein
weak acid-weak base
pH at equivalence:
suitable indicator:
depends on the relative strength of th acid and base
What is a buffer solution
solution that resists changes in pH when small amounts of acid pr alkali are added
NB- pH does change, just not by much
Acidic buffer
buffers that have a pH lower than 7
Basic buffer
buffers that have a solution higher than 7
What is an acidic buffer made of?
mixture of a weak acid and one of its salts (ie HA and A-)
What is key in an acidic buffer?
the [acid] and [salt] are much higher than the [H+]
What are the two routes to make an acidic buffer?
+ how to calculate the pH of the buffer in each route
Route 1: mixture of weak acid and one of its salts (use [A-] and [HA] with Ka equation to calculate H+)
Route 2: Mixture oif an excess of weak acid and a strong base
(HA+OH—>A- +H2O to calulate left over HA and formed A- then use Ka)
What is a basic buffer made of?
a mixture of a weak alkali and one of its salts
What is the key in a basic buffer
the [base] and the [salt] are much higher than [OH-]
What are the two routes to make an basic buffer?
Route 1: mixture of weak base and one of its salts
Route 2: mixture of an excess weak base and a strong acid
How do buffer solutions work?
buffers pH depends on the ratio of [acid]:[salt] or [base]:[salt] so when small amounts of acid (H+) or base (OH-) are added the ratio remains roughly constant and so the pH hardly changes.
What happens when you add a little acid to an acidic buffer?
HA⇌ H+ + A-
The added H+ is removed by reacting with A- to form more HA.
The [A- ]falls slightly and the [HA] rise slighty but as [HA] & [A-]»_space; [H+] the ratio of [HA]:[A-] remains roughly constant.
What happens when you add a little base to an acidic buffer?
HA⇌ H+ + A-
The added OH- reacts with H+ and so some HA breaks down to replace that H+
The [A- ]rises slightly and the [HA] falls slighty but as [HA] & [A-]»_space; [H+] the ratio of [HA]:[A-] remains roughly constant.
What happens when you add water to an acidic buffer?
The ratio of [HA]:[A-] remains constant so pH remains constant
How do you calculate when acid is added to an acidic buffer
- calc mol acid added
- A-+H+–>HA before and after
- plug into Ka
- to get new H+
How do you calculate when base is added to an acidic buffer
-calc mol OH- added
- HA+OH—>A- +H2O before and after
- calc [HA] and [A-] after
- plug into Ka to find H+
same pocess as finding pH of acidic buffer formed via routte 2
