chemical equilibrium - University of Lincoln

Report
Reversible Reactions
and Chemical
Equilibrium
University of Lincoln
presentation
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Outline
• Reversible reactions
• Chemical Equilibrium
• Le Chatelier’s Principle
• Equilibrium constants
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Reversible Reactions
BiCl3(aq) + H2O(l) ↔ BiOCl(s) + 2HCl(aq)
CH3CO2H + CH3CH2OH ↔ CH3CO2CH2CH3 + H2O
Cr2O72-(aq) + 2OH-(aq) ↔ 2CrO42-(aq) + H2O(l)
CH3CO2H(aq) + H2O(l) ↔ CH3CO2-(aq) + H3O+(aq)
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Chemical Equilibrium
• Reactions not 100% complete
–Products and Reactants exist together
• A dynamic equilibrium
• Position of equilibrium ???
• Can the position of equilibrium be
changed?
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Le Chatelier’s Principle
When an external change is made to a system
in equilibrium, the system will respond to
oppose the change
External Changes
• Concentration
• Pressure (gases)
• Temperature
Link to external video
Link to external video
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Concentration
1. BiCl3(aq) + H2O(l) ↔ BiOCl(s) + 2HCl(aq)
2. Cr2O72-(aq) + 2OH-(aq) ↔ 2CrO42-(aq) + H2O(l)
How does reaction 1 respond to addition of
hydrochloric acid?
How does reaction 2 respond to addition of alkali?
How does reaction 2 respond to addition of acid?
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Pressure
N2(g) + 3H2(g) ↔ 2NH3(g)
CO(g) + 2H2(g) ↔ CH3OH(g)
2NO2(g) ↔ 2NO(g) + O2(g)
PCl5(g) ↔ PCl3(g) + Cl2(g)
H2(g) + I2(g) ↔ 2HI(g)
CO(g) + H2O(g) ↔ CO2(g) + H2(g)
How do the above equilibria respond to:
An increase in pressure
A decrease in pressure
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Temperature
N2(g) + 3H2(g) ↔ 2NH3(g)
H2(g) + I2(g) ↔ 2HI(g)
CO(g) + H2O(g) ↔ CO2(g) + H2(g)
PCl5(g) ↔ PCl3(g) + Cl2(g)
rH = -92.2 kJ mol-1
rH = -9.4 kJ mol-1
rH = -41.2 kJ mol-1
rH = 87.9 kJ mol-1
How do the above respond to an
Increase in temperature
Decrease in temperature
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Equilibrium constants
a measure of equilibrium position
aA + bB ↔cC + dD
[C]c [D]d
Kc 
[A]a [B]b
BiCl3(aq) + H2O(l) ↔ BiOCl(s) + 2HCl(aq)
Kc 
[BiOCl (s) ][HCl (aq) ] 2
[BiCl 3 (s) ][H 2 O (l) ]
Write the expressions for Kc for the reactions given in previous
slides
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Calculating Equilibrium Constants
HNO2(aq) ↔ H+(aq) + NO2-(aq)
The table shows the equilibrium molar concentrations for
three solutions of nitrous acid in water at 25 oC
Solution [HNO2(aq)] [H+(aq)]
mol litre-1
mol litre-1
[NO2-(aq)]
mol litre-1
A
0.090
6.2 x 10-3
6.2 x 10-3
B
C
0.20
9.3 x 10-3
9.3 x 10-3
0.30
11.4 x 10-3
11.4 x 10-3
Calculate the equilibrium constant for this reaction at 25oC
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Solution A

[H (aq)][NO2 (aq)] 6.2  103  6.2  103
Kc 

 4.3  104
[HNO2 (aq)]
0.090
Units of Kc
(mol litre1 )(mol litre1 )
1

mol
litre
Kc 
mol litre1
Kc  4.3  104 mol litre1
Now try for solutions B and C
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Acids and Bases
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Outline
• Definitions
• Weak Acids
• Dissociation Constants
• Weak Bases
• Drugs
• pH
• Buffers
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Acids and Bases
• Several definitions available - most common is
Bronsted and Lowry
• Acid is a proton donor
– HCl is able to transfer H+
• Base is a proton acceptor
– NH3 is able to accept H+ and become NH4+
• Aqueous solutions
• Proton species is H3O+ (hydroxonium ion)
– HCl(aq) + H2O(l)  H3O+(aq) + Cl-(aq)
– HCl(aq)  H+(aq) + Cl-(aq)
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Strong Acids
• Strong acids are fully dissociated
HCl (aq) + H2O(l) H3O+ (aq) + Cl- (aq)
• all dissolved HCl molecules are ionised
• 1 mol dm-3 HCl(aq) there are:
– Approx 1 mol dm-3 H3O+ (aq)
– Approx 1 mol dm-3 Cl- (aq)
DO NOT confuse ‘strong’ and ‘concentrated’
1 x 10-4 mol dm-3 HCl (aq) is a dilute solution of a
strong acid
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Other strong acids
• HNO3 (nitric)
• H2SO4 (sulfuric)
• HClO4 (perchloric)
Write equations showing the dissociation of
the above acids
Which are monoprotic?
Are any diprotic?
Chemical equilibrium – K very large
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Weak Acids
• Acids that dissociate in a reversible reaction
(e.g. CH3COOH; ethanoic (acetic) acid)
CH3COOH (aq) + H2O(l) ↔ H3O+ (aq) + CH3COO- (aq)
• Solution of CH3COOH (aq) contains:
–CH3COOH (aq)
–H3O+ (aq)
–CH3COO- (aq)
• CH3COOH is partially dissociated
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How weak is a weak acid?
0.1 mol dm-3 HCl is dissociated 91.4%
[H3O+] = 0.091 mol dm-3
pH=1.04
0.1 mol dm-3 CH3COOH is dissociated 1.34%
[H3O+] = 0.0013 mol dm-3
pH=2.87
• Extent given by K
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Weak Acids
HA(aq) + H2O(l) ↔ H3O+(aq) + A-(aq)
HA Bronsted acid
H2O Bronsted base
H3O+ Bronsted acid
A- Bronsted base


[H3 O ][A ]
K
[HA][H2 O]
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Acid dissociation constant (Ka)


[H3 O ][A ]
K
[HA][H2 O]
[H3 O  ][A ]
K[H2 O] 
[HA]
[H3 O ][A ]
K[H2 O]  K a 
[HA]
• The higher the Ka value:
–greater degree of ionisation
–stronger the acid
–Data tables
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Ka Values
• HCO2H
1.8 x 10-4 mol dm-3
• CH3CO2H
1.7 x 10-5 mol dm-3
• Are these weak or strong acids?
• Which is the stronger acid?
pKa  logKa
HCO2H
CH3CO2H
3.75
4.77
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pKa values (data tables)
Acid
pKa
Conjugate base
H3PO4
2.12
H2PO4-
HNO2
3.34
NO2-
H2CO3
6.37
HCO3-
HCN
9.31
CN-
HCO3-
10.25
CO32-
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pKa Values
• Controlling the ionisation of weak acids
• pH = pKa then [HA] = [A-]
• pH > pKa then [A-] > [HA]
• pH < pKa then [HA] > [A-]
CH3COOH (aq) + H2O(l) ↔ H3O+ (aq) + CH3COO- (aq)
• CH3COOH: CH3COO- at pH = 4.77 ?
• CH3COOH: CH3COO- at pH = 3 ?
• CH3COOH: CH3COO- at pH = 7 ?
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Henderson-Hasselbach
For weak acids
[ionised]
pH  pKa  log
[un  ionised]
Use the equation with the example in the previous slide.
Do you come to the same conclusion regarding the ratio
of un-ionised to ionised acid molecules?
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Weak Bases
B(aq) + H2O(l)  BH+(aq) + OH-(aq)
CH3NH2(aq) +H2O(l) ↔ CH3NH3+(aq) + OH-(aq)
pKa = 10.66 (of conjugate acid) [B]=[BH+]
pH = 10.66
pH =8 what happens to CH3NH3+(aq): CH3NH2(aq)
pH =13 ?
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Henderson-Hasselbach
For weak bases
[un  ionised]
pH  pKa  log
[ionised]
Use the equation with the example in the previous slide.
Do you come to the same conclusion regarding the ratio
of un-ionised to ionised acid molecules?
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CH3
Acidic drugs
ibuprofen
O
CH3
OH
H3C
2-[4-(2-methylpropyl)phenyl]propanoic acid
How does this molecule ionise?
pKa=4.5
pH =3 (stomach pH)?
pH=6 (intestine)?
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Basic drugs
CH3
NH2
amphetamine (C6H5CH2CH(NH2)CH3)
Write an equation for the reaction of
amphetamine with water.
The pKa of the conjugate acid is 9.8. What
will happen to the ratio of ionised to
unionised amphetamine at:
pH 7
pH 12
Why might this be important?
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Water
• Can dissociate:
H2O(l) ↔ H+(aq) + OH-(aq)
2H2O(l) ↔ H3O+(aq) + OH-(aq)
H2O is amphoteric


[H3 O ][OH ]
K
2
[H2 O]
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Water
Kw = [H3O+][OH-]= 1 x 10-14 mol2 dm-6
• Kw the ionic product of water
• In pure water what is [H3O+] and [OH-] ?
• Kw is a very small constant
– water is only very partially ionised
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pH
• pH is defined as:
pH = -log10[H3O+]
• pH is a measure of the H3O+ concentration
in solution and can vary from 1 to 14
• pH=7 – neutral [H3O+] = [OH-]
= 1 x 10-7 mol dm-3 at 25 oC
• pH<7 – acidic [H3O+] >[OH-]
• pH>7 - alkaline/basic [H3O+] <[OH-]
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pH-examples
• 0.1M HNO3
• 0.1M CH3COOH
• What is the pH?
• pH is dependent on the ionisation of the
acid
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pH-examples
• What about alkaline solutions?
• E.g. 0.1M NaOH solution
• Will also depend on degree of ionisation
• use equation: [H+] x [OH-] = 10-14
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Buffers
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Buffers
• A buffer solution resists pH changes on
addition of small amounts of acid or base
(alkali) to a system.
• Very important
– e.g. blood has a pH of 7.4. If it varies by ± 0.4,
death can occur
• Buffer solutions rely upon the effects of a
weak acid or base and the salt of that
acid or base
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Buffers
• Ethanoic acid (a weak acid) and sodium
ethanoate (salt)
CH3COOH  CH3COO- + H+
CH3COONa  CH3COO- + Na+
(1)
(2)
• (1)-partially ionised
• (2)-fully ionised
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Buffers
Henderson-Hasselbach equation
Acidic buffers
[salt]
pH  pKa  log
[acid]
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Making a buffer solution
• Choose a weak acid with a pKa close to the
required pH of the buffer.
• Choose an appropriate salt of the weak
acid
• Determine [salt]/[acid] ratio needed to
give correct pH
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An acidic buffer:
Ethanoic acid and sodium ethanoate
What would be the pH of an ethanoate buffer
with equal acid and sodium ethanoate
concentrations?
[salt]
pH  pKa  log
[acid]
pH  4.77  0
What is the [salt] if the acid is 0.1 mol dm-3 to
give buffer solutions of
pH = 5
pH = 4
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An alkaline buffer:
ammonia solution and ammonium chloride
[base]
pH  pKa  log
[salt]
Note the base/salt ratio
What is the pH of a buffer with base:salt ratio = 1?
pH  9.24  log1  9.24
Calculate the base:salt ratios for pH 8.5 and pH
10.5
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Acknowledgements
•
•
•
•
•
•
•
JISC
HEA
Centre for Educational Research and Development
School of natural and applied sciences
School of Journalism
SirenFM
http://tango.freedesktop.org
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