5.1.3 Acids, Bases and Buffers Flashcards
Bronsted lowry acid
Substance that can donate a proton
Bronsted lowry base
Substance that can accept a proton
How to calculate ph
pH = -log [H+]
Conjugate base pairing which acts as an acid
Substance with higher Ka
Role of H+ in acid metal equations
2H+ + Mg -> Mg2+ +H2
Role of H+ in acid alkali equations
H+ + OH- -> H2O
Role of H+ in acid carbonate equations
2H+ +CO3 2- -> H2O + CO2
Calculating pH of strong acids
Completely dissociate
H+ is equal to conc of acid
What do you give pH values to in exams
2 dp
how to calculate [H+]
[H+] = 10^-ph
Ionic product for water describe
In all aqueous solutions and pure water the neutralisation of water equilibrium occurs
Uses Kw
Kw
Kw = [H+] [OH-]
Value of Kw at 25 degrees
1 x 10^-14 mol2dm-6
What can Kw be used to find
[H+]
[OH-]
Finding the pH of pure water
Neutral because H+ = OH-
Kw = [H+]^2
So H+ = square root of Kw
Different temps pH of pure water
Changes
Dissociation of water is endothermic as bonds are broken
Increasing temp pushes equilibrium to right
Bigger conc of H+ ions and lower pH
Calculating ph of strong bases
Normally get given conc of hydroxide ion
To work out pH work out H+ using Kw
Strong bases completely dissociate into their ions
Weak acids
Only slightly dissociate when dissolved in water
Uses Ka
Whats the weak acids dissociation expression
Ka= [H+][A-] divided by [HA]
Larger Ka stronger the acid
Products divided by reactants
pKa
pKa = -logKa Ka= 10^-pka
Calculating pH of a weak acid
Ka = [H+]^2 divided by [HA]
Assumptions made to simplify Ka
- [H+] = [A-] because they have dissociated according to a 1:1 ratio
- As the amount of dissociation is small we assume the initial conc of the undissociated acid has remained constant
pH of diluted strong acid
H+ = H+ old x old volume divided by new volume
Then pH = -log [H+]
pH of diluted base
OH- = OH- old x old volume divided by new volume H+ = Kw divided by OH- pH = - logH+
What is a buffer solution
One where the pH doesnt change significantly if small amounts of acid or alkali are added to it
Basic buffer solution
Made from a weak base and a salt of that weak base (react weak base with a strong acid)
Ammonia and ammonium chloride
Acidic buffer solution
Made from a weak acid and a salt of that weak acid (reacting weak acid with strong base)
Ethanoic acid and sodium ethanoate
How could salt content be added
Salt solution added to acid or some solid salt added
Buffer made by partially neutralising a weak acid with alkali and produces a mixture of salt and acid
Ethanoic acid buffer
CH3CO2H (reversible) CH3CO2- + H+
Much higher conc of salt ion than in pure acid
If small amounts of acid are added to buffer
Equilibrium shift to left
Removing H+ added
As theres a large conc of the salt ion in the buffer the conc ratio CH3CO2H/CH3CO2- stays almost constant so pH stays constant
If small amounts of alkali are added to buffer
OH- ions will react with H+ ions to form water
Equilibrium shifts to right to produce more H+ ions
Conc of H+ and pH remains constant but some ethanoic acid molecules are changed to ethanoate ions
Calculating pH of buffer solutions
Ka= H+ x A- divided by HA
Assume that A- conc is due to added salt only
Assume intial conc of acid has remained constant because amount that has dissociated or reacted is small
If small amount of alkali is added to a buffer what happens (moles)
Moles of buffer acid would reduce by number of moles of alkali added and the moles of salt would increase by same amount
CH3CO2H + OH- -> CH3CO2- + H2O
If small amount of acid is added to a buffer what happens (moles)
Moles of buffer salt would reduce by the number of moles of acid added and the moles of buffer acid would increase by the same amount
CH3CO2- + H+ -> CH3CO2H
Buffering action in blood
Carbonic acid-hydrogencarbonate equilibrium acts as buffer in control of blood pH
H2CO3/HCO3- buffer present in blood plasma mainting pH 7.35-7.45
Equilibrium in blood
H2CO3 (reversible) H+ + HCO3-
Adding alkali reacts with H+ so shifts to right forming new H+ and more HCO3-
Constructing a pH curve
- Acid into conical flask
- Measure intial pH with pH meter
- Add alkali in small amounts noting the volume
- Stir mixture to equalise pH
- Measure and record pH
- Repeat but when approaching end point add smaller volumes of alkali
- Add until alkali in excess
Need to calibrate meter first
Calibrating meter in pH curve
Measure known pH of buffer solution
pH meters can lose accuracy on storage
Put probe in a set buffer and press calibrate
Maintain constant temp also to improve accuracy
4 main types of titration curves
- Strong acid and strong base
- Weak acid and strong base
- Strong acid and weak base
- Weak acid and weak base
Strong acid strong base titration curve
HCl and NaOH
pH at equivalence point 7
Long steep part 3 to 9
work out neutralisation volume from titration data given in question, standard titration calculations
Key points to sketching titration curve
Intial and final pH
Volume at neutralisation
General shape (pH at neutralisation)
Where does the equivalence point lie
Lies at mid point of extrapolated vertical point of curve
Weak acid strong base
CH3CO2H and NaOH at start the pH rises quickly and then levels off, flattened part is buffer region and is formed because buffer solution is made Starts near 3 Steep part >7 (7 to 9) Equivalence point >7
Half neutralisation volume
Use Ka at 1/2 HA= A
Ka = H+ and pKa = pH
If you know Ka can work out pH at half
pH of weak acid at half neutralisation will equal pKa
Strong acid weak base
HCl and NH3
Equivalence <7
Steep <7 around 4 to 7
Weak acid weak base
CH3CO2H and NH3
No vertical part
What can indicators be considered as
Weak acids
Acid must have different colour to conjugate base
End point of a titration
[Hln]=[ln-]
Choose indicator whose end-point coincides with equivalence point for titration
Applying Le Chateliers to give colour
Hln (reversible) ln- + H+
Colour A. Colour B
In acid solution H+ ions push equilibrium towards reactants, colour A is acidic colour
In alkaline solution the OH- ions will react and remove H+ ions causing equilibrium to shift to products, colour B is alkaline colour
When will an indicator work
If pH range of indicator lies on stteo part of titration curve
Indicator will change colour rapidly and the colour change will correspond to neutralisation point
When do you use phenolphthalein
In titrations with strong bases but not weak bases
Colour change: colourless acid -> pink alkali
When do you use methyl orange
Titrations with strong acids but not weak acids
Colour change: red acid -> yellow alkali (orange end point)
