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18... Flashcards

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1
Q

lewis acid and bases theory

A

general definiton of acids in terms of electron pair acceptor/donor

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2
Q

lewis acid

A

electron pair acceptor

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3
Q

lewis base

A

electron pair donor

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4
Q

what is the lewis base and lewis acid in a complex ion

A
  • water molecule - lewis base
  • metal ion - lewis acid
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5
Q

lewis acids are electrophile or nucleophiles?

A

electrophiles
- electron -deficient species that can accept a lone pair from a nucleophile

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6
Q

pH curves

A

shows the pH of a solution changes as the acid (or base) is added

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7
Q

featured of a pH curve

A
  • all pH curves show an s-shaped curve
  • pH curves yield useful information about how the acid and alkali react together
  • midpoint of the inflection is called the equivalence point
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8
Q

what can be determined from the curves

A
  • determine the pH of the acid by looking where the curve starts on the y-axis
  • find the pH at the equivalence point
  • Obtain the range of pH at the vertical section of the curve
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9
Q

Four typed of Acid-Base Titration

A

strong acid + strong base
weak acid + strong base
weak base + strong acid
weak acid + weak base

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10
Q

pH curve of a strong acid and strong base

A
  • base added to acid
  • pH intercepts on the y-axis starts at a low pH due to strength of acid
  • as base is added there is a gradual rise in pH until titration approaches the equivalence point
  • once all acid has been neutralised, curve flattens out and continues to rise gradually
  • at the end, pH will be high due to strength of base
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11
Q

pH curve of a weak acid and strong base

A
  • base added to acid
  • pH on the intercept on the y axis starts at roughly 3 due to strength of acid
  • initial rise in pH is steep as the neutralisation of the weak acid by the strong base is rapid
  • conjugate base of the acid is formed, creating a buffer
  • buffer formed will resist changes in pH so the pH rises gradually, “buffer region”
  • half equivalance point is the stage where exactly half the weak acid has been neutralised
  • at this point pKa = pH at half equivalance
  • equivalance point in this titration is above 7
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12
Q

What is a buffer

A

consists of a weak acid and its conjugate base or weak base and its conjugate acid

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13
Q

pH curve of a weak base and strong acid

A
  • acid added to base
  • pH on the intercept on the y axis at roughly 11 doe to strength of base
  • pH will fall as base is neutralised and conjugate acid is produced
  • “buffer region”, so pH falls gradually
  • “half equivalence point is the stage where half the amountof weak base has been neutralised
  • pKb = pOH at half equivalence
  • equivalence point is below 7
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14
Q

pH curve of a weak acid and weak base

A
  • acid is added base
  • pH roughly at 11 due to strength of base
  • change in pH for titration is very gradual
  • equivalance point is hard to determine
  • equicalence point is roughly 7 but difficult to determine
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15
Q

what is an acid-base indicator?

A
  • weak acid which dissociates to give an anion of a different color
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16
Q

what does the color of the indicator depend on

A
  • pH of the solution
  • ## color does not change suddenly at a certain pH but changes gradually over pH range
17
Q

when does color change occurs in indicators

A

pH=pKa +- 1

18
Q

How to choose a suitable indicator

A
  • equivalence point of titration - where the pH changes rapidly
  • indicators change color over a narrow pH range - aprox. centered arounf the pKa of indicator
  • indicator will be good for a titration if the pH range of indicator falss within the rapid change for titration
19
Q

indicator for strong acid and strong base

A

pH changes from 4 to 10 at end-point
- methyl red and phenophthalein

20
Q

indicator for weak acid - strong base

A

pH changes from 7 to 10 at end point
- Phenolphthalein

21
Q

indicator for strong acid and weak base

A

pH changes from 4 to 7 at end-point
- methyl red
- methyl orange

22
Q

indicator for weak acid and weak bases

A
  • no sudden pH cnages
  • no suitable indicator
23
Q

buffer solution

A

solution which resists changes in pH when small amounts of acid or base are added
A buffer solution is used to keep the pH almost constant
A buffer can consist of weak acid – conjugate base or weak base – conjugate acid

24
Q

common buffer solution

A

aq. mixture of ethanoic acid and sodium ethanoate
CH3COOH (aq) ⇌ H+ (aq) + CH3COO- (aq)

high conc. high conc.

25
Q

what occurs when H+ ions are added to buffer solution

A
  • equilibrium shifts to the left
  • pH remains remains constant
  • conjugate base reacts with H+ ions
26
Q

what occurs when OH- ions are added to buffer solution

A
  • OH- reacts with H+ to form water
  • H+ conc decreases
  • more acid ionise to form H+ and conjugate base ions
  • ## pH remains reasonably constant
27
Q

salt hydrolysis / neutralisation

A
  • ionic salt is formed from the neutralisation reaction of an acid and base
  • ionic salt, MA, formed will dissociate in water
  • reaction of the salt will vary depending on the strength of the acids and bases used in the neutralisation reactions
  • The use of the differing strengths of the acids and bases will directly influence the type of salt hydrolysis and the pH of the final solution
28
Q

salt hydrolysis for strong acids and strong bases

A
  • ions do not act as Bronsted-Lowry acids or bases
  • no change in pH
29
Q

Salt hydrolysis of strong acid and weak bases

A
  • salt formed is acidic solution
  • solution becomes more acidic
  • ## pH drops
30
Q

salt hydrolysis of strong base and weak acid

A
  • salt formed is alkaline solution
  • solution becomes more alkali
  • pH increases
31
Q

salt hydrolysis of weak acid and weak base

A
  • in order to determine the pH of the resulting solution of a reaction between a weak acid and weak base we must take into account the Ka and Kb values
  • Using the reaction between ammonia, NH3 (aq), and ethanoic acid, CH3COOH (aq), as an example:
    NH3 (aq) + CH3COOH (aq)→ CH3COONH4 (aq)
  • Both the cation (positive ion) and anion ion (negative) produced will have acid-base properties
    CH3COO– (aq) + H2O (l) → CH3COOH (aq) + OH- (aq)
    NH4+ (aq) + H2O (l) → H3O+ (aq) + NH3 (aq)

Ka (cation) = kw / kb(parent base)
Kb (anion) = kw / ka (parent acid)

If the Ka is larger, the solution will be acidic
If the Kb is larger the solution will be basic
If Ka = Kb, then the pH will be 7

32
Q

metal salt hydrolysis

A
  • small metal ions that have high charge will exhibit a high charge density (Al3+)
  • makes a highly charges metal ions iseal for forming complexes as they can coordinately bond with ligands
  • complexes can act as weak acids
  • high charge density of metal ion, increases polarity of the ater molecule pulling the e- towars itself, until O-H bond breaks
  • metal ion with 1+ and 2+ charge will not releace an H+ ion, dereasing the pH
33
Q

Acid dissociation constant, Ka

A
  • weak acids
    HA (aq) + H2O (l) ⇌ A- (aq) + H3O+ (aq)
    OR
    HA (aq) ⇌ A- (aq) + H+ (aq)

When writing the equilibrium expression for weak acids, we assume that the concentration of H3O+ (aq) due to the ionisation of water is negligible

34
Q

base dissociation constant, Kb

A
  • weak base
    B (aq) + H2O (l) ⇌ BH+ (aq) + OH- (aq)
35
Q

converting between pKa and Ka / pKb and Kb

A
36
Q

relative strengths of acid and bases

A

The larger the Ka value, the stronger the acid
The larger the pKa value, the weaker the acid
The larger the Kb value, the stronger the base
The larger the pKb value, the weaker the base

37
Q

Kw

A

Kw = [H3O+][OH-]

1x10^-14

38
Q

Kw in relation to Ka/pKa and Kb/pKb

A

(Ka)(Kb)=Kw
(Ka)(Kb)=10^-14

pKa + pKb = pKw
pKa + pKb = 14

[H3O+] = Kw / [OH-]

39
Q

relationship between Kw and temperature

A
  • ionisation of water is endothermic
  • inc. temp = decreases pH of water
  • dec. temp = increases pH of water

Chem (13 decks)

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