Download Slide 1

Chemistry II
Inorganic Chemistry
Part 2
Prof HM Marques
C301
717-6737
[email protected]
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
Complex ion
2+
NH3
H3N
3+
NH3
Co
OH-
H3 N
NH3
Ligands
2+
NH3
H3N
3+
NH3
Co
OH-
H3 N
NH3
H
N
H
Co
H
Lewis base
Lewis acid
Water is amphiprotic (amphoteric) – it can act as either an acid or a
base
HCl(g) + H2O(l)  H3O+ (aq) + Cl- (aq)
Base
NH3(g) + H2O(l)  OH- (aq) + NH4+(aq)
Acid
(H3O)+(H2O)3 or H9O4+
HA(aq) + H2O (l)  H3O+(aq) + A-(aq)
Ka
Acid
dissociation
constant
[H3O+ ][A- ]
=
[HA]
Measure of the
strength of the
acid
pKa   log ( Ka )
Acid strength decreases
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
The higher the pKa, the weaker the acid
Ka
[H3O+ ][A- ]
=
[HA]
[H3O+ ] =
K a [HA]
[A- ]
 K [HA] 
log [H3O+ ] =  log  a - 
 [A ] 
 [HA] 
log [H3O+ ] =  log K a  log  - 
 [A ] 
 [A- ] 
log [H3O ] =  log Ka  log 

[HA]


+
 [A- ] 
pH = pKa  log 

[HA]


 [A- ] 
pH = pKa  log 

[HA]


 [Base] 
pH = pK a  log 

 [Acid] 
Henderson-Hasselbalch equation
•When [Base] = [Acid], pH = pKa
•At any pH, it can be shown (see Tut) that
%[Base] =
100
1  10pKa  pH
%[Acid] =
100
1  10pH  pKa
B(aq) + H2O (l)  OH-(aq) + HB+(aq)
Kb
Basicity
constant
[OH- ][HB+ ]
=
[B]
Measure of the
strength of the
base
pK b   log ( K b )
The higher the pKb, the weaker the base
H2O (l) + H2O (l)  OH-(aq) + H3O+(aq)
Kw
Autoprotolysis
constant of
water
= [OH - ][H 3O + ]
= 1.00 x 10-14 at
25.0 oC
Oxalic acid is a polyprotic acid:
HO
OH
(H2Ox)
O
O
H2Ox + H2O  HOx- + H3O+
Ka1 = 5.9 x 10-2
HOx- + H2O  Ox2- + H3O+
Ka2 = 6.4 X 10-5
We will calculate the species distribution as a function of pH
K a1
Fractional
abundance,
so 0  x  1
=
[HOx - ][H + ]
[H 2Ox]
K a2
=
x(H 2 Ox) + x(HOx - ) + x(Ox 2- )
[Ox 2- ][H + ]
[HOx - ]
= 1
 Ka2 [HOx - ] 
 Ka1[H2Ox] 
x[H 2Ox] + x 
 = 1
 + x
+
+
 [H ] 
 [H ] 
 Ka1[H 2Ox] 
 K a1 K a2 [H 2 Ox] 
x[H 2Ox] + x 
 + x

+
+ 2
[H
]
[H
]




  Ka1   Ka1Ka2  
x[H 2Ox] 1   +  + 
+ 2 
[H
]
[H
] 
 
 
= 1
= 1
  Ka1   Ka1Ka2  
x[H 2Ox] 1   +  + 
+ 2 
[H
]
[H
] 
 
 
 [H+ ]2  Ka1[H + ]  Ka1 Ka2 
x[H 2Ox] 

+ 2
[H
]


x[H 2Ox] =
[H + ]2
[H + ]2  K a1[H + ]  K a1K a2
Similarly,
K a1[H + ]
x[HOx ] =
β
K a1 K a2
2x[Ox ] =
β
-
= 1
= 1

[H + ]2
β
[H + ]2
x[H 2 Ox] =
β
K a1[H + ]
x[HOx ] =
β
K a1 K a2
2x[Ox ] =
β
-
1.0
HOx-
0.9
H2Ox
Fractional abundance
0.8
Ox2-
0.7
0.6
0.5
0.4
0.3
0.2
0.1
0.0
0.0
2.0
4.0
6.0
8.0
pH
10.0
12.0
14.0
Species distribution at pH 3
1.0
0.93
0.9
H2Ox
HOx-
Ox2-
0.8
0.7
0.6
0.5
0.4
0.3
0.2
0.1
0.05
0.02
0.0
0.0
0.5
1.0
1.5
2.0
2.5
pH
93% HOx– ; 5% Ox2- ; 2% H2Ox
3.0
3.5
4.0
4.5
5.0
The acidity of a proton
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…
…then the X-H bond
can split heterolytically
with X retaining the
electron pair…
R––X: + H+
…and delivering H+ to
a Lewis base
Organic acids have acidic and non-acidic protons
O
H2
C
H3 C
C
C
H2
O
H
Non-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
If metal is able to polarise the M-O
bond towards it…
L
L
L
M
L
O
L
H
…that will cause the H-O bond to
be polarised towards O…
H
…releasing H+ to be accepted by a
Lewis base.
AQUA ACIDS
Acidic proton on a water molecule coordinated to a metal ion
L
L
L
L
M
L
O
L
H
L
H
L
M
L
L
O
H
+
H+
Aqua acids are in principle polyprotic acids
[Fe(OH2)6]3+

[Fe(OH2)5(OH)]2+ + H+
 [Fe(OH2)4(OH)2]+
etc.
+ H+
HYDROXOACIDS
Acidic proton on a hydroxyl group bonded or coordinated to a central
atom
OH
Si
HO
OH
OH
OXOACIDS
Acidic proton on a hydroxyl group bonded or coordinated to a central
atom on which there is an oxo (=O) group
Aqua acids, hydroxoacids and oxoacids may be successive stages
of the deprotonation of an aqua acid
OH
OH2
L - 2 H+
L
L
Ru
L
L
- H+
L
Ru
L
OH2
O
L
Ru
L
OH
L
L
L
OH
Aqua acids
OH2
L
Central atom in lower
oxidation states
L
M
L
L
OH2
s block, d block
metals in lower
(+1, +2, +3)
oxidation states,
metals on the left
of the p block
Strengths of aqua acids
This can be rationalised using an electrostatic (ionic) model
O
H
H
H
r+
radius of the
metal ion of
charge n+
d
diameter of
coordinated water
molecule

H+
Work done = [Potential at r = ] – [Potential at r = (r+ + d)]
E
1 q1q2
 
4πε 0 r
q1q2
1
 0  
4πε 0 (r  d )

1 zcation zH
4πε 0 (r  d )
For this process,
G = -RT ln K = -nFE
RT ln K = nF(E)
-RT ln K = -nFE(r
-log K =
pKa =
+d)
+
-2.303 nFE(r
+d)
+
RT
-2.303 nFE(r
RT
++d)
pKa =
-2.303 nFE(r
RT
E
E( r  d )
++d)

1 zcation zH
4πε 0 (r  d )
1 zcation zH
4πε 0 (r  d )
zcation
Work
1
=
=
charge
zH 
4πε 0 (r  d )
pKa =
E( r  d )
-2.303 nFE(r
++d)
RT

1 zcation zH
4πε 0 (r  d )
zcation
Work
1
=
=
charge
zH 
4πε 0 (r  d )
 1
zcation 
2.303nF 

4 o (r  d ) 

pK a 
RT
 1 zcation 
2.303nF 

4 o (r  d ) 

pK a 
RT
2.303nFzcation

4 o (r  d ) RT
zcation
-pK a 
(r )
pKa should become smaller, and acidity should increase with

an increase in the charge on the ion

a decrease of the size of the ion
z
-pK a 
( r )
The term (z/r+) is also known as the ionic potential 
Alternatively we could say
2
z
-pK a 
(r  d )
by adding the charge on the proton and the diameter of water.
The term (z2/(r+ + d)) is called the electrostatic parameter 
pKa gets smaller and acidity increases
How good is this electrostatic model (for gas phase) in solution?
as electrostatic parameter increases...
See Fig. 4.3
How good is this electrostatic model (for gas phase) in solution?
Model quite poor for many
of the d block ions; their
acidity is often much higher
than predicted by the model
Model quite good for s block ions,
some d block ions, and the
lanthanides
Major reason for failure of the model: bonding between the metal and its
ligands often not purely ionic, and there is some covalency in metalligand bonds.
Model is worst for metals like Sn2+ and Hg2+ that form very covalent
complexes.
L
L
The more
covalent the
M-O bond…
…the more the O-H
bond is polarised
towards O…
L
M
L
O
L
H
H
…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, is
O1(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
O
Example
Estimate the pKas of H3AsO4
As
HO
Actual values: 2.3, 6.9, 11.5
*Estimates are usually good to within 1-2 units
OH
OH
The strengths of oxoacids can be varied by substitution:
O
O
O
O
S
S
S
OH
OH
O
OH
O
CF3
OH
F
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
O
O
S
S
OH
OH
O
...polarisesO
the OH
bond, making the
proton more acidic
OH
O
CF3
Electron withdrawl...
S
OH
F
The strengths of oxoacids can be varied by substitution:
O
O
O
O
S
S
S
OH
OH
O
OH
O
CF3
OH
F
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
Oxides
Oxides 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)
O
O
+ H2O
S
O
O
S
O
SO3 is the anhydride of H2SO4
OH
OH
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 nonmetals 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
Amphoteric axides will react with acids and bases…
Ga2O3 + 6H3O+ + 3H2O → 2[Ga(H2O)6]3+
Ga2O3 + 2OH– + 3H2O → 2[Ga(OH)4]–
Polymerisation of aqua ions
Aqua ions of metals that have basic or amphoteric oxides polymerise and
precipitate as pH is increased.
[Fe(OH2)6]3+
3+
OH2
H2O
- exists in strongly acidic solution
OH2
increase pH
H2O
OH2
OH2
OH2
Fe
Fe
H2O
2+
OH2
H2O
OH
OH2
[Fe(H2O)6]3+(aq) + (3+n)H2O(l) → Fe(OH)3•nH2O(s) + 3H3O+(aq)
polymer of Fe(OH)3
Similarly:
[Al(H2O)6]3+(aq) + (3+n)H2O(l) → Al(OH)3•nH2O(s) + 3H3O+(aq)
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 PO43As base is added, condensation reactions occur, and
polyoxyanions are formed
VO42-
V2O72-
V3O93-
[H2V10O28]4-
Phosphates polymerise by condensation:
H+
H2 O
etc
See pp. 123-125
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
H
H3C
B
CH3
H
H3C
H3C
B
N
CH3
H
H
H3C
H
N
H
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 hard acid
If for B = halide, the order of log
K is
I– > Br– > Cl– > F–
then A is called a soft acid
Fig. 4.10
Al3+ is a hard Lewis acid
Log K increases with the ionic potential
 = z/r
Anion–
Al3+
r
ionic radius increases
F

1 zmetal zanion
4πε o
r2
Strength of complex  1/r2
 bonding is largely ionic
Hg2+ is a soft Lewis acid
Log K increases as the radius of the
anion increases
Hg2+ Anion–
Hg2+ Anion–
ionic radius increases
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
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 state
Cu(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
Soft Bases: contain the larger, more polarisable and less
electronegative atoms, especially S, Se, P, C and As.
Common soft bases include:
• H- (hydride)
• C donors
• S, P, As donors
• I-
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)
(b)
(c)
(d)
CH3SCH3
CH3OCH3
CH3SCH3O-
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)