prof hm marques [email protected] chemistry ii inorganic chemistry part 2...
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Prof HM MarquesProf HM [email protected]@chem.wits.ac.za
Chemistry IIChemistry IIInorganic ChemistryInorganic ChemistryPart 2Part 2
Chapters 4 and 5
Chapter 4 – Acids and Bases
Revise from Chem I
Bronsted-Lowry definition of an acid & a base
Lewis definition of an acid & a base
Co
H3N
H3N NH3
OH-
NH3
NH3
3+
2+Complex ion
Ligands
Co
H3N
H3N NH3
OH-
NH3
NH3
3+
2+
N
H
HH
Co
Lewis base Lewis acid
HCl(g) + H2O(l) H3O+ (aq) + Cl- (aq)
NH3(g) + H2O(l) OH- (aq) + NH4+(aq)
Base
Acid
Water is amphiprotic (amphoteric) – it can act as either an acid or a base
(H3O)+(H2O)3 or H9O4+
HA(aq) + H2O (l) H3O+(aq) + A-(aq)
+ -3
a
[H O ][A ]=
[HA]K
Acid dissociation
constantMeasure of the strength of the
acid
log ( )a apK K
Perchloric acid (HClO4) Ka = 1010 pKa = -10
Sulphuric acid (H2SO4) 102 -2
Phosphoric (H3PO4) 7.5 x 10-3 1.92
Hydrocyanic acid (HCN) 4.9 x 10-10 9.31
Acid strength decreases
The higher the pKa, the weaker the acid
+ -3
a
[H O ][A ]=
[HA]K
+ a3 -
[HA][H O ] =
[A ]
K
+ a3 -
[HA]log [H O ] = log
[A ]
K
+3 a -
[HA]log [H O ] = log log
[A ]K
-+
3 a
[A ]log [H O ] = log log
[HA]K
-
a
[A ]pH = p log
[HA]K
-
a
[A ]pH = p log
[HA]K
a
[Base]pH = p log
[Acid]K
Henderson-Hasselbalch equation
•When [Base] = [Acid], pH = pKa
•At any pH, it can be shown (see Tut) that
ap pH
100%[Base] =
1 10 K
apH p
100%[Acid] =
1 10 K
B(aq) + H2O (l) OH-(aq) + HB+(aq)
- +
b
[OH ][HB ]=
[B]K
Basicity constant Measure of the
strength of the base
b blog ( )pK K
The higher the pKb, the weaker the base
H2O (l) + H2O (l) OH-(aq) + H3O+(aq)
- +w 3= [OH ][H O ]K
Autoprotolysis constant of
water= 1.00 x 10-14 at
25.0 oC
Oxalic acid is a polyprotic acid:
H2Ox + H2O HOx- + H3O+ Ka1 = 5.9 x 10-2
HOx- + H2O Ox2- + H3O+ Ka2 = 6.4 X 10-5
OH
O
HO
O
We will calculate the species distribution as a function of pH
(H2Ox)
- + 2- +
a1 a2 -2
[HOx ][H ] [Ox ][H ]= =
[H Ox] [HOx ]K K
- 2-2(H Ox) + (HOx ) + (Ox ) = 1x x x
-a1 2 a2
2 + +
[H Ox] [HOx ][H Ox] + + = 1
[H ] [H ]
K Kx x x
a1 2 a1 a2 22 + + 2
[H Ox] [H Ox][H Ox] + + = 1
[H ] [H ]
K K Kx x x
a1 a1 a22 + + 2
[H Ox] 1 + = 1[H ] [H ]
K K Kx
Fractional abundance, so 0 x
1
a1 a1 a22 + + 2
[H Ox] 1 + = 1[H ] [H ]
K K Kx
+ 2 +a1 a1 a2
2 + 2
[H ] [H ][H Ox] = 1
[H ]
K K Kx
+ 2 + 2
2 + 2 +a1 a1 a2
[H ] [H ][H Ox] =
[H ] [H ] βx
K K K
Similarly,
+- a1
2- a1 a2
[H ][HOx ] =
β
[Ox ] =β
Kx
K Kx
+ 2
2
+- a1
2- a1 a2
[H ][H Ox] =
β
[H ][HOx ] =
β
[Ox ] =β
x
Kx
K Kx
0.0
0.1
0.2
0.3
0.4
0.5
0.6
0.7
0.8
0.9
1.0
0.0 2.0 4.0 6.0 8.0 10.0 12.0 14.0
pH
H2OxHOx-
Ox2-
Fra
ctio
nal a
bund
ance
0.0
0.1
0.2
0.3
0.4
0.5
0.6
0.7
0.8
0.9
1.0
0.0 0.5 1.0 1.5 2.0 2.5 3.0 3.5 4.0 4.5 5.0
pH
H2Ox HOx-
Ox2-
0.020.05
0.93
Species distribution at pH 3
93% HOx– ; 5% Ox2- ; 2% H2Ox
For a proton to be acidic it must be attached to an electronegative element (O, F, Cl, Br, I; to a lesser extent N, S)
R––X––H
If X is electronegative…
R––X: + H+
…and delivering H+ to a Lewis base
…then the X-H bond can split heterolytically
with X retaining the electron pair…
The acidity of a proton
Organic acids have acidic and non-acidic protons
H3C
H2C
CH2
C
O
O
HNon-acidic protons because C is not
electronegative enough
Acidic proton because H bonded to
electronegative O
AQUA ACIDS
Acidic proton on a water molecule coordinated to a metal ion
M
L
L L
O
L
L H
H
If metal is able to polarise the M-O bond towards it…
…that will cause the H-O bond to be polarised towards O…
…releasing H+ to be accepted by a Lewis base.
AQUA ACIDS
Acidic proton on a water molecule coordinated to a metal ion
M
L
L L
O
L
L H
HM
L
L L
O
L
L H
- H++
Aqua acids are in principle polyprotic acids
[Fe(OH2)6]3+ [Fe(OH2)4(OH)2]+ + H+[Fe(OH2)5(OH)]2+ + H+
etc.
HYDROXOACIDS
Acidic proton on a hydroxyl group bonded or coordinated to a central atom
OH
Si
HOOH
OH
OXOACIDS
Acidic proton on a hydroxyl group bonded or coordinated to a central atom on which there is an oxo (=O) group
Ru
L
L L
L
OH2
OH2
Ru
L
L L
L
OH
OH
- 2 H+
Ru
L
L L
L
O
OH
- H+
Aqua acids, hydroxoacids and oxoacids may be successive stages of the deprotonation of an aqua acid
Aqua acids
M
L
L L
L
OH2
OH2
Central atom in lower oxidation states
s block, d block metals in lower (+1, +2, +3) oxidation states, metals on the left of the p block
OH
H
r+ d
Strengths of aqua acids
This can be rationalised using an electrostatic (ionic) model
radius of the metal ion of charge n+
diameter of coordinated water
molecule
H H+
Work done = [Potential at r = ] – [Potential at r = (r+ + d)]
1 2
0
1
4πε
q qE
r
1 2
0
H
0
10
4πε ( )
1
4πε ( )
cation
q q
r d
z z
r d
OH
H
r+ d
H H+ H+
For this process,
G = -RT ln K = -nFE
RT ln K = nF(E)
-RT ln K = -nFE(r +d)+
RT
-2.303 nFE(r +d)+-log K =
RT
-2.303 nFE(r +d)+pKa =
RT
-2.303 nFE(r +d)+pKa =
H
0( )
0H
14πε ( )Work 1
= = charge 4πε ( )
cation
cationr d
z z
r d zE
z r d
OH
H
r+ d
H H+ H+
H
0
1
4πε ( )cationz z
Er d
RT
-2.303 nFE(r +d)+pKa =
H
0( )
0H
14πε ( )Work 1
= = charge 4πε ( )
cation
cationr d
z z
r d zE
z r d
oa
12.303
4 ( )p
cationznF
r dK
RT
a-p( )
cationzK
rpKa should become smaller, and acidity should increase with
an increase in the charge on the ion
a decrease of the size of the ion
oa
o
12.303
4 ( )p
2.303
4 ( )
cation
cation
znF
r dK
RTnFz
r d RT
a-p( )
z
Kr
The term (z/r+) is also known as the ionic potential
Alternatively we could say2
a-p( )
z
Kr d
by adding the charge on the proton and the diameter of water.
The term (z2/(r+ + d)) is called the electrostatic parameter
How good is this electrostatic model (for gas phase) in solution?
See Fig. 4.3
as electrostatic parameter increases...
pKa
gets
sm
alle
r an
d ac
idity
incr
ease
s
How good is this electrostatic model (for gas phase) in solution?
Model quite good for s block ions, some d block ions, and the lanthanides
Model quite poor for many of the d block ions; their acidity is often much higher than predicted by the model
Major reason for failure of the model: bonding between the metal and its ligands often not purely ionic, and there is some covalency in metal-ligand bonds.
Model is worst for metals like Sn2+ and Hg2+ that form very covalent complexes.
M
L
L L
O
L
L H
H
The more covalent the M-O bond…
…the more the O-H bond is polarised
towards O…
…the more readily H+ is lost, and the more acidic
the compound
Oxoacids
Formed by
electronegative elements top right of periodic table (e.g., N, P, S, Cl)
elements in high oxidation state (e.g., Te, I, As, Se)
See Table 4.2
General formula: OpE(OH)q
H3PO3, phosphorus acid, isO1(PH)(OH)2
Pauling’s rules
Empirical rules that allow one to estimate the pKa of oxoacids
General formula: OpE(OH)q
1. pKa 8 – 5p
2. For p > 1, each successive pKa increases by about 5 units
Example
Estimate the pKas of H3AsO4
O
As
HOOH
OH
Actual values: 2.3, 6.9, 11.5
*Estimates are usually good to within 1-2 units
The strengths of oxoacids can be varied by substitution:
O
S
OOH
OH
O
S
OCF3
OH
O
S
OF
OH
CF3 and F are more electron withdrawing than OH; these acids are stronger acids than H2SO4
The strengths of oxoacids can be varied by substitution:
O
S
OOH
OH
O
S
OCF3
OH
O
S
OF
OH
Electron withdrawl...
...polarises the OH bond, making the proton more acidic
The strengths of oxoacids can be varied by substitution:
O
S
OOH
OH
O
S
OCF3
OH
O
S
OF
OH
CF3 and F are more electron withdrawing than OH; these acids are stronger acids than H2SO4
NH2 and CH3 are electron donating – hence these acids are weaker acid than H2SO4
OxidesOxides of non-metals are acidic.
When dissolved in water, they bind water and release a proton
SO3(g) + H2O(l) → H2SO4(aq) → H+(aq) + HSO4–(aq)
S
O
O O
+ H2O
O
S
OOH
OH
SO3 is the anhydride of H2SO4
Acidic oxides are neutralised by bases
SO2 + NaOH → Na+HSO3–
Oxides of metals are basic.
When dissolved in water, they accept a proton from water, producing an alkaline solution
MgO(s) + H2O(l) → Mg(OH)2(s) Mg2+(aq) + 2OH–(aq)
Acidic and basic oxides neutralised each other
CaO + SO2 → CaSO3
Oxides of the elements in the boundary region between metals and non-metals often show amphoteric behaviour.
Fig. 4.4
In the d block, metals in low oxidation states tend to be basic; amphoteric in their intermediate oxidation states; and acidic in high oxidation states
Fig. 4.5
Ga2O3 + 6H3O+ + 3H2O → 2[Ga(H2O)6]3+
Ga2O3 + 2OH– + 3H2O → 2[Ga(OH)4]–
Amphoteric axides will react with acids and bases…
Polymerisation of aqua ionsAqua ions of metals that have basic or amphoteric oxides polymerise and precipitate as pH is increased.
[Fe(OH2)6]3+ - exists in strongly acidic solution
Fe
H2O
H2O OH2
OH2
OH2
OH2
3+
increase pHFe
H2O
H2O OH2
OH
OH2
OH2
2+
[Fe(H2O)6]3+(aq) + (3+n)H2O(l) → Fe(OH)3•nH2O(s) + 3H3O+(aq)
polymer of Fe(OH)3
[Al(H2O)6]3+(aq) + (3+n)H2O(l) → Al(OH)3•nH2O(s) + 3H3O+(aq)
Similarly:
As pH is increased further, the species redissolve because both Al(OH)3 and Fe(OH)3 are amphoteric.
Al(OH)3 + OH– → [Al(OH)4]–
Polyoxyanions
With early 3d elements, or oxides of elements in high oxidation states
eg V2O5
eg PO43-
As base is added, condensation reactions occur, and polyoxyanions are formed
VO42- V2O7
2- V3O93- [H2V10O28]4-
Phosphates polymerise by condensation:
H+
H2O
See pp. 123-125
etc
Lewis Acids and Bases
We have seen that compounds such as [Co(NH3)6]3+ are complexes between ligands (Lewis bases) and a metal (Lewis acid)
Electron deficient compounds can act as Lewis acids
H3C B
CH3
CH3
H
N
HH
B
H3C
H3CH3C
N
H
HH
Self-study:Write brief notes on the Lewis acid properties of compounds of the elements of the s block, of Group 13, 14, 15, 16 and the halogens
(p. 126-128)
Hard and Soft Acids and Bases (HSAB)
R.G. Pearson
Let A be a Lewis acid, and B a base
Measure log K for the reaction
A + B AB
If for B = halide, the order of log K is
I– > Br– > Cl– > F–
then A is called a soft acid
If for B = halide, the order of log K is
I– < Br– < Cl– < F–
then A is called a hard acid
Fig. 4.10
Al3+ is a hard Lewis acid
Log K increases with the ionic potential = z/r
Al3+ Anion–
r
2o
1
4πεmetal anionz z
Fr
Strength of complex 1/r2
bonding is largely ionic
ionic radius increases
Hg2+ is a soft Lewis acid
Log K increases as the radius of the anion increases
Hg2+ Anion–
Strength of complex increases as overlap between orbitals of the anion sand orbitals of the metal increases with increasing size of anion
bonding is largely covalent
ionic radius increases
Hg2+ Anion–
Similarly, if we take a hard metal ion (like Al3+), then bases which bind strongly to it are hard Lewis bases; bases which do not bind strongly are soft Lewis bases.
Pearson’s Principle:
Hard Lewis acids prefer to bind to hard Lewis bases; soft Lewis acids prefer to bind to soft Lewis bases
See Table 5.3 for a listing of hard and soft Lewis acids and bases
In summary:
Hard metal ions are small, usually highly charged. Their electron cloud is not readily polarisable.
• Alkali metals (Li+, Na+, … )
• Alkali earth metal (Mg2+, Ca2+, … )
• H+
• Lighter transition metals in their higher oxidation states: Ti(IV), Ti(III), Cr(III), Fe(III), Co(III) …
Hard Bases: contain the smaller electronegative atoms, especially O, N, F and Cl. These donor atoms also have rather unpolarisable electron clouds
Soft metal ions are larger metal ions, often in their lower oxidation states. Their electron cloud is readily polarisable.
• Heavier transition metals (Pt, Rh, Ir)
• Transition metals in their lower oxidation stateCu(I), Ag(I), Hg(I), Hg(II), Pd(II), Pt(I), Pt(II)
• As we move across the d block the +2 oxidation state is stabilised –i.e., get Sc(III), not Sc(II); but Zn(II) not Zn(III). Hence
softness tends to increase across the d block, and down each group
Common soft bases include:
• H- (hydride)
• C donors
• S, P, As donors
• I-
Soft Bases: contain the larger, more polarisable and less electronegative atoms, especially S, Se, P, C and As.
Example
Arrange the following ligands in the order of increasing log K for binding to Fe(III) and to Pb(II) (donor atoms underlined)
(a) CH3SCH3
(b) CH3OCH3
(c) CH3S-
(d) CH3O-
Example
Explain why Cu(I) and Cu(II) are found in nature as the sulphide (CuS, Cu2S) but Ti(IV) and Fe(III) are found as their oxides (TiO2, Fe2O3)