# Acids and bases

Revision cards for AQA unit 4 acids and bases.

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## An acid is a proton donor

• Bronsted - Lowry acids are proton donors
• release hydrogen ions when they are mixed with water.
• never get H+ ions by themselves in water though - always combined with H2O to form hydroxide ions H3O+
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## Base is a proton acceptor.

Bronsted Lowry bases do the opposite - they are proton acceptors

when they are in solution, they grab hydrogen ions from water molecules.

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## Acid - base eqm involves the transfer of protons

Acids and bases can only protons if there is a base to accept them

in this reaction the acid, HA, transfers a proton to the base, B.

It's and eqm, so if you add acid more HA or B, the position of the eqm moves to the right.

If you add more BH than or A, the eqm will move to the left.

• Follows Le Chatellier's principle

When an acid is added to water, the water acts as the base and accepts the proton.

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## pH = -log10 [H+]

pH scale is a measure of H+ concentration.

conc of H+ ions can vary enormously.

log scale makes the numbers more manageable.

0 = very acidic

14 = very alkaline.

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## calculation of pH of a strong acid

e.g. HCl and HNO3 (citric acid)

for strong acids, hydrogen conc = acid conc --> a strong acid fully dissociates.

Monoprotic acids produce one mole of H+ for every mol of acid.

Method

• if given conc then pH will have to be calculated (pH=-log10[H+])
• if given the pH the conc will have to be calculated ([H+]=10^(-pH))
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## Water is weakly dissociated

eqm lies well over to the left

conc of water considered to be a constant value

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## Kw = [H+][OH-]

Kw = ionic product of water

Kw always has the same value for an aqueous solution at a given temp.

value at 25 degrees C = 0.1 * 10^-14

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## calculating pH of a strong base from its concentra

• e.g NaOH & KOH
• fully ionise in water
• each have one hydroxide ion per molecule
• one mol of OH- per mol of base.
• conc of OH- ions is the same conc as the base.
• Method
• to work out pH of a base [H+] needs to be known.
• linked to [OH-] through Kw (Kw = [H+][OH-] = 1.0*10^(-4) at 298K.
• [H+] = Kw/[OH-]
• Then work out pH (pH = -log10 [H+])
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## Weak acids and bases only dissociate slightly in s

• only small [H+] formed.
• eqm lies well over to the left.
• e.g. NH3 = weak base, CH3COOH = weak acid.
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## Dissociation constant for a weak acid (Ka)

• [H+] isn't the same as the acid conc.
• therefore, Ka needed to find pH
• to find eqm constant you have to make a couple of assumptions:-
• only a tiny amount of HA dissociates, so you can assume that:
• [HA]start= [HA]eqm.
• so this gives:- Ka = [H+][A-]/[HA]
• All the H+ ions come from the acid, so [H+]=[A-]
• so Ka = [H+]^2 /[HA] = mol dm-3
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## pKa =-log10*Ka

• pKa is calculated from Ka in exactly the same way as pH is calculated from [H+] - and vice versa.
• Use Ka expression to convert pKa to Ka
• Method:
• convert pKa to Ka (Ka=10^-pKa)
• write down Ka expression (Ka=[H+]^2 /[HA])
• calculate [H+]
• [H+]^2 = Ka[HA]
• [H+] = sqrt [H+]^2
• calculate pH (pH = -log [H+])
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## pH curves

pH curves plot pH against volume of acid or base acid.

• Strong acid/strong base
• pH starts around 1, excess of strong acid
• finishes around pH 13, excess of strong base.
• Strong acid/weak base
• pH starts around 1, excess of strong acid
• finishes around pH 9, excess of weak base
• weak acid/ strong base
• pH starts around 5, excess of weak acid
• finishes around pH13, excess of strong base.
• weak acid/ weak base
• pH starts around 5, excess of weak acid
• finishes around pH 9, excess of weak base.
• All graphs have a bit that is almost vertical = equivalence point or end point.
• at this point a tiny amount of base cases a sudden change in pH
• all acid neutralised.
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## pH curves

• end point does not occur in weak acid/weak base titration
• pH meter would be required to find end point as pH indicator would not be sensitive enough.
• titrating a base with an acid cause a flipped curve.
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## using pH curves to select an appropriate indicator

• The indicator selected needs to change colour at end point of titration.
• methyl orange and phenolphthalein are indicators that are often used for acid-base titration.
• each change colour over a different pH range.
• Methyl orange = 3.1 (red)--> 4.4 (yellow)
• phenolphthalein = 8.3 (colourless) --> 10 (pink)
• which one to use:
• strong acid/strong alkali = use either
• strong acid/weak alkali = methyl orange
• weak acid/strong alkali = phenolphthalein.
• weak acid/weak base = indicators won't work.
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## calculations for the titrations of mono- and dipro

• results of titration can be used to calculate concentrations of acids and bases.
• Increasing accuracy of results:-
• measure neutralisation volume as closely as poss (nearest 0.05)
• repeat titration at least 3 times and get a mean value
• Don't use anomalous results (all results should be 0.2cm^3 of each other.)
• Method:
• write a balanced equation and decide what is known and what needs to be known
• now work out the moles of HA
• Look at the ratio of acid to base. (1:1 then moles of base = moles of acid) (1:2 then moles of base = moles of acid / 2) (2:1 then moles of base = 2* moles of acid.)
• Now work out what needs to be known.
• diprotic acids needs twice as many moles of base as 2 H+ needs to be removed.
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## Actions of acidic and bases buffers.

• Buffer resists changes in pH when a small amount of acid or alkali is added.
• Acidic buffers
• pH less than 7
• weak acid + one of its salts
• e.g. ethanoic acid and sodium ethanoate
• salt fully dissociates when it dissolves
• CH3COO-Na+ --> CH3COO- + Na+
• ethanoic acid is a weak acid
• CH3COOH (eqm symbol) -->H+ + CH3COO-
• La chateller's principle explains how buffers work
• addition of H+ (acid) moves eqm to left
• addition of OH- (alkali) moves eqm to right
• pH doesn't change much
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## Actions of acidic and bases buffers.

• Basic buffers
• made from weak base and one of its salts
• pH greater than 7
• e.g. ammonia and ammonium chloride
• NH3 + H2O (eqm) --> NH4+ + OH-
• eqm position moves to counteract pH
• addition of H+ (acid) moves eqm to right
• addition of OH- (alkali) moves eqm to left.
• avoids pH changing much
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## Applications of buffer solutions

• important in biological environments
• cells need a constant pH to allow the biochemical reactions to take place. The pH is controlled by a buffer based on an eqm between dihydrogen phosphate and hydrogen phosphate.
• H2PO4- (eqm)--> (H+) + HPO4(2-)
• blood needs to be kept at pH 7.4. it's buffered using carbonic acid (H2CO3) levels of H2CO3 are controlled by the body. by breathing out CO2 the level of H2CO3 is reduced as it moves the eqm to the right. Levels of HCO3- are controlled by kidneys with excess being excreted by the urine.
• Buffers are used in food products to control the pH. Changes in pH can ve caused by bacteria and fungi and cause food to deteriorate. A common buffer is sodium citrate, which sets up in eqm between citrate ions and citric acid. Phosphoric acis/phosphate ions and benzoate ions are used as buffers.
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## Calculating pH of acidic buffer solution

need to know Ka of weak acids and the conc of the weak acid and its salt.

Method:

• write expression for Ka of weak acid (Ka = [H+]*[A-]/[HA]
• rearrange to calculate [H+] = Ka * ([HCOOH]/[HCOO-])
• convert [H+] to pH (pH = -log10[H+]
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