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Chemistry Module 5 > 5.1.3 Acids, Bases and Buffers > Flashcards

5.1.3 Acids, Bases and Buffers Flashcards

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1
Q

Bronsted lowry acid

A

Substance that can donate a proton

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2
Q

Bronsted lowry base

A

Substance that can accept a proton

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3
Q

How to calculate ph

A

pH = -log [H+]

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4
Q

Conjugate base pairing which acts as an acid

A

Substance with higher Ka

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5
Q

Role of H+ in acid metal equations

A

2H+ + Mg -> Mg2+ +H2

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6
Q

Role of H+ in acid alkali equations

A

H+ + OH- -> H2O

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7
Q

Role of H+ in acid carbonate equations

A

2H+ +CO3 2- -> H2O + CO2

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8
Q

Calculating pH of strong acids

A

Completely dissociate

H+ is equal to conc of acid

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9
Q

What do you give pH values to in exams

A

2 dp

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10
Q

how to calculate [H+]

A

[H+] = 10^-ph

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11
Q

Ionic product for water describe

A

In all aqueous solutions and pure water the neutralisation of water equilibrium occurs
Uses Kw

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12
Q

Kw

A

Kw = [H+] [OH-]

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13
Q

Value of Kw at 25 degrees

A

1 x 10^-14 mol2dm-6

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14
Q

What can Kw be used to find

A

[H+]

[OH-]

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15
Q

Finding the pH of pure water

A

Neutral because H+ = OH-
Kw = [H+]^2
So H+ = square root of Kw

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16
Q

Different temps pH of pure water

A

Changes
Dissociation of water is endothermic as bonds are broken
Increasing temp pushes equilibrium to right
Bigger conc of H+ ions and lower pH

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17
Q

Calculating ph of strong bases

A

Normally get given conc of hydroxide ion
To work out pH work out H+ using Kw
Strong bases completely dissociate into their ions

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18
Q

Weak acids

A

Only slightly dissociate when dissolved in water

Uses Ka

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19
Q

Whats the weak acids dissociation expression

A

Ka= [H+][A-] divided by [HA]
Larger Ka stronger the acid
Products divided by reactants

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20
Q

pKa

A 
pKa = -logKa
Ka= 10^-pka
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21
Q

Calculating pH of a weak acid

A

Ka = [H+]^2 divided by [HA]

22
Q

Assumptions made to simplify Ka

A
  1. [H+] = [A-] because they have dissociated according to a 1:1 ratio
  2. As the amount of dissociation is small we assume the initial conc of the undissociated acid has remained constant
23
Q

pH of diluted strong acid

A

H+ = H+ old x old volume divided by new volume

Then pH = -log [H+]

24
Q

pH of diluted base

A 
OH- = OH- old x old volume divided by new volume
H+ = Kw divided by OH-
pH = - logH+
25
Q

What is a buffer solution

A

One where the pH doesnt change significantly if small amounts of acid or alkali are added to it

26
Q

Basic buffer solution

A

Made from a weak base and a salt of that weak base (react weak base with a strong acid)
Ammonia and ammonium chloride

27
Q

Acidic buffer solution

A

Made from a weak acid and a salt of that weak acid (reacting weak acid with strong base)
Ethanoic acid and sodium ethanoate

28
Q

How could salt content be added

A

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

29
Q

Ethanoic acid buffer

A

CH3CO2H (reversible) CH3CO2- + H+

Much higher conc of salt ion than in pure acid

30
Q

If small amounts of acid are added to buffer

A

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

31
Q

If small amounts of alkali are added to buffer

A

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

32
Q

Calculating pH of buffer solutions

A

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

33
Q

If small amount of alkali is added to a buffer what happens (moles)

A

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

34
Q

If small amount of acid is added to a buffer what happens (moles)

A

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

35
Q

Buffering action in blood

A

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

36
Q

Equilibrium in blood

A

H2CO3 (reversible) H+ + HCO3-

Adding alkali reacts with H+ so shifts to right forming new H+ and more HCO3-

37
Q

Constructing a pH curve

A
  1. Acid into conical flask
  2. Measure intial pH with pH meter
  3. Add alkali in small amounts noting the volume
  4. Stir mixture to equalise pH
  5. Measure and record pH
  6. Repeat but when approaching end point add smaller volumes of alkali
  7. Add until alkali in excess

Need to calibrate meter first

38
Q

Calibrating meter in pH curve

A

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

39
Q

4 main types of titration curves

A
  1. Strong acid and strong base
  2. Weak acid and strong base
  3. Strong acid and weak base
  4. Weak acid and weak base
40
Q

Strong acid strong base titration curve

A

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

41
Q

Key points to sketching titration curve

A

Intial and final pH
Volume at neutralisation
General shape (pH at neutralisation)

42
Q

Where does the equivalence point lie

A

Lies at mid point of extrapolated vertical point of curve

43
Q

Weak acid strong base

A 
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
44
Q

Half neutralisation volume

A

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

45
Q

Strong acid weak base

A

HCl and NH3
Equivalence <7
Steep <7 around 4 to 7

46
Q

Weak acid weak base

A

CH3CO2H and NH3

No vertical part

47
Q

What can indicators be considered as

A

Weak acids

Acid must have different colour to conjugate base

48
Q

End point of a titration

A

[Hln]=[ln-]

Choose indicator whose end-point coincides with equivalence point for titration

49
Q

Applying Le Chateliers to give colour

A

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

50
Q

When will an indicator work

A

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

51
Q

When do you use phenolphthalein

A

In titrations with strong bases but not weak bases

Colour change: colourless acid -> pink alkali

52
Q

When do you use methyl orange

A

Titrations with strong acids but not weak acids

Colour change: red acid -> yellow alkali (orange end point)

Chemistry Module 5 (7 decks)

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