CHEM UNIT 4: ACID BASE EQUILIBRIA

Acids releases protons - base accepts protons

Bronsted-Lowry acids are proton donors - release hydrogen ions (H+) when they're mixed with water. You never get H+ ions themselves in water though - always combined with H2O to form hydroxonium ions, H3O+

HA(aq) + H2O(l) ---> H3O+(aq) + A-(aq)

Bronsted-Lowry bases are proton acceptors. When they're in solution they grap H+ ions from water molecules

B(aq) + H2O(l) ---> BH+(aq) + OH-(aq)

Acids and bases can be strong or weak

• strong acids ionise almost completely in water - nearly all the H+ ions will be released. HCl is a strong acid. Strong bases like NaOH ionise almost completely in water too - NaOh(s) = water ---> Na+(aq) + OH-(aq)

• weak acids (ethanoic acid) ionise only very slightly in water - only a small no. of H+ ions are formed. an equilibrium is set up which lies well over to the left (CH3COOH(aq)          CH3COO-(aq) + H+(aq). weak bases (ammonia) only slightly ionise in water too (NH3(aq) + H2O(l)        NH4+(aq) + OH-(aq)

Protons are transferred when acids and bases react

acids can only get rid of their protons if there's a base to accept them

HA(aq) + B(aq)      BH+(aq) + A-(aq)

If you add more HA or B, the position of equilibrium moves to the right. When an acid is added to water, the water acts as a base and accepts the proton HA(aq) + H2O(l)       H3O+(aq) + A-

acids and bases form conjugate pairs

water can behave as an acid and a base

2H2O(l)      H3O+(aq) + OH-(aq) or H2O(l)     H+(aq) + OH-(aq)

Kc = [H+][OH-] / [H2O]

water only dissociates a tiny amount so the equilibrium lies well over to the left - so much water compared to the amounts of H+ and OH- ions that the conc of water is said to be a constant - ionic product of water

Kw = Kc x [H2O] = [H+][OH-] ---> Kw = [H+][OH-]

Kw always has the same value for an aqeous solution at a given temp. at standard temp. 25C (298K) Kw = 1.0 x 10^-14 mol^2/dm^6

A neutral solution has equal H+ and OH- concentrations

[H+] = [OH-]

if [H+] is greater than [OH-] - solution is acidic, [OH-] is greater than [H+] - solution is alkaline

pH scale is a measure of the hydrogen ion conc.

pH = -log[H+]

for strong monoprotic acids, hydrogen ion conc = acid conc

HCl and HNO3 are strong acids so they ionise fully and are also monoprotic - each mole of acid produces one mole of hydrogen ions - H+ conc is the same as the acid conc. so for 0.1 mol/dm^3 HCl, [H+] = 0.1 mol/dm^3 pH = -log[H+] = -log(0.1) = 1

[H+] = 10 ^-pH

Use Kw to find the pH of a base

NaOH and KOH are strong bases that ionise fully in water. The have 1 hydroxide ion - donate 1 mole of OH- ions per mole of base = conc of OH- ions is the same as the conc of the base. So for 0.02 mol/dm^3 NaOH, [OH-] = 0.02mol/dm^3

Kw = [H+][OH-] = 1.0 X 10^-14 at 298K

pH of a weak acid you use Ka (acid dissociation constant)

weak acids don't ionise fully in solution so [H+] isn't the same as acid conc. 

For a weak aqueous acid, HA, you get: HA(aq)       H+(aq) + A-(aq)

• only a tiny amount of HA dissociates so you can assume that [HA]start = [HA]equi
• Ka = [H+][A-] / [HA]
• can also assume that all the H+ ions come from the acid, so [H+] = [A-] so Ka = [H+]^2 / [HA] 

pKa = -logKa and Ka= 10^-pKa 

pH of equimolar solutions can give you info about the substances

• 1 mol/dm^3 HCl has pH 0 - [H+] =10^-pH = 1mol/dm^3 which is also its conc. so it must be completely ionised - it's a strong acid
• 1mol/dm^3 NaCl has pH 7 = [H+] = 1 x 10^-7, using Kw = [H+][OH-] = 1 x 10^-14, [OH-] = 1 x 10^-7 too, [H+] = [OH-], so substance is neutral - true for salts of strong acids with strong bases 

when acids are diltuted their pH changes 

diluting an acid reduces the conc. of H+ in the solution. This increases the pH 

strong acid - HCl 

• diluting a strong acid by a factor of 10 increases the pH by 1 [H+] = [acid] 

weak acid - propanoic acid 

• diluting a weak acid by a factor of 10 increases the pH by 0.5 

Use titration to find the conc of an acid or alkali 

titrations - find out exactly how much alkali is needed to neutralise a quantity of acid

• measure out some acid of known conc. using a pipette and putting it in a flask with an appropriate indicator 
• first do a rough titration - add the alkali to acid fairly quickly using a burette to get a n approximate idea of where the solution changes colour (the end point) give flask a swirl 
• now do an accurate titration - run alkali in to within 2cm^3 of the end point, then add it drop by drop. If you don't exactly notice when the solution changes colour you've overshot and the result won't be accurate 
• record the amount of alkali needed to neutralise the acid. best to repeat process several times

pH curves plot pH against volume of acid or alkali added 

pH curves can help you decide which indicator to use 

methy orange - red to yellow from low pH to high pH - colour change at 3.1 - 4.4 

phenolphthalein - colourless to pink from low pH to high pH - colour change at 8.3 - 10

• for strong acid/strong alkali - can use either - rapid pH change over the range for both indicators 
• for strong acid/weak alkali - only methy orange will do, the pH changes rapidly across the range for methyl orange but not for phenolphthalein 
• for weak acid/strong alkali - use phenolphthalein 
• for weak acid/weak alkali - no sharp pH change - so neither will work 

find pKa of weak acid 

• using the pH curve for a weak acid/strong base titration, involves finding the pH at the half equivalence point 
• half equivalence point - stage of titration when half of the acid has been neutralised, when half of the equivalence volume of strong base has been added to weak acid 

At half equivalence point [HA] = [A-] 

so for the weak acid HA Ka = [H+][A-] / [HA]    Ka = [H+]   pKa = pH

so pH at half equivalence is actually the pKa value for the weak acid and if you know the pKa value then you work out the Ka 

Buffers - solutions that resist change in pH when small amounts of acid or alkali are added 

buffer doesn't stop the pH from changing completely - it does make the changes very slight though, only work for small amounts of acid or alkali 

acidic buffers are made from a weak acid and one of its salts 

acidic buffers have a pH of less than 7 - made by mixing a weak acid with one of its salts eg ethanoic acid + sodium ethanoate - salt fulli dissociates into its ions when it dissolves the ethanoic acid is a weak acid and only slightly dissociates = so you have heaps of ethanoate ions from the salt and heaps of undissociated ethanoic acid molecules 

• add small amount of acid, H+ conc. increases - most of the extra H+ ions combine with the ethanoate ions to form ethanoic acid - shifts equilibrium to left, reducing the H+ conc to near its original value. so pH doesn't change much 
• small amount of alkali added, OH- conc increases, most of extra OH- ions reacts with H+ ions to form water - removing H+ ions from solution - causes more acid to dissociate to form H+ ions - shifts equilibrium to right 

Alkaline buffers are made from a weak base and one of its salts eg mixture of ammonia solution and ammonium chloride acts as an alkaline buffer 

- acts in same way as acidic buffers 

buffer action can be seen on titration curve 

important in biological environments 

• cells need a constant pH to allow the biochemical reactions to take place. pH is controlled by a buffer based on the equilibrium between dihydrogen phosphate and hydrogen phosphate 

• blood needs to be kept at pH 7.4, buffered using carbonic acid (H2CO3) which the levels are controlled by the body. by breathing out CO2 the level of it is reduced as it moves the equilibrium to the right. Levels of HCO3- is controlled by the kidneys w/ excess being excreted in the urine 
• buffers are used in food products to control the pH. changes in pH can be caused by bacteria and fungi and cause food to deteriorate, common buffer = sodium citrate which sets up an equilibrium between citrate ions and citric acid. Phosphoric acid/phosphate ions and benzoic acid/benzoate ions are also used as buffers 

how to calculate the pH of a buffer solution 

a buffer solution contains 0.40 mol/dm3 methanpoic acid, HCOOH and 0.6mol/dm3 sodium methanoate, HCOO-Na+. For methanoic acid, Ka = 1.6 x 10^-4 mol/dm3. what is the pH of this buffer?