Pre- & Post- Transition Metals Flashcards
transition metal
an element from the d-block that normally forms complexes with an incomplete d-shell
pre-transition metals
metals before central d-block
A group metals
- in group 1, group 2, group 3 + Lanthanides
-also includes Al
group A elements: largely hard (but with exceptions)
post-transition metals
B group transition metals
- elements in group 11-12
- metals in group 13-16
Group B elements: largely soft
hard and soft acids and bases
hard prefers hard
soft prefers soft
in this definition, acid and base refer to lewis acid and lewis base
hard and soft refers to the polarisability of a metal ion or species
- hard ion: does not deform much when brought up to ion of opposite charge
- soft ion: readily deforms when brought up to ion of opposite charge
lewis acid
electron pair acceptor (Na+)
lewis base
electron pair donor (F-, NH3)
hard (for lewis acids)
~ pre-transition metals are hard
~ transition metals in high oxidation states are hard
a high oxidation state would be a charge of +3 or higher
- however, for 2nd and 3rd row transition metals, sometimes higher oxidation states are needed
- Pt4+ is soft
soft (for lewis acids)
~ post-transition metals are soft
~ transition metals in low oxidation states are soft
hard (for lewis base)
F, N, O donors are typically all hard bases
F-, -OH, NH3, RNH2, ROH
soft (for lewis base)
all other donors are soft bases
S, Se, Te, As, P donors are typically all soft bases
H2S, PR3, H3R3, TeR2
determine if transition metal ion is hard/soft
look at the stability of the corresponding halide
if we have: F- > Cl- > Br- > I-
metal cation is hard
if we have: F- < Cl- < Br- < I-
metal cation is soft
stability of complex determined by looking at the stability constant (Kn, Bn)
Kn - stability constant
for a metal ion (cation) combining with a ligand (lewis base) the equilibrium constant is called the stability constant or formation constant
addition of ligand occurs stepwise
K1 - first ligand addition
K2 - second ligand addition
βn - overall stability constant
considers adding all the ligands in one go (rather than stepwise)
βn = K1 x K2 x K3 … Kn
trends in stability constants
if stability constant for a reaction is high, then that complex will be readily formed and is encompassed by our definition of a strong complex via HSAB principle
it is commonly observed that the stepwise stability constants lie in the order:
K1 > K2 > K3 > K4 … Kn
stepwise reduction in Kn values
consider reactions in water and the number of sites available for a displacement reaction
K1 - 6 sites
K2 - 5 sites
K3 - 4 sites
K4 - 3 sites
K5 - 2 sites
K6 - 1 sites
less sites available for L - therefore, statistically less likely to happen
exceptions to stepwise reduction in Kn values
occur when there is a structural change at the metal (i.e. 6 to 4 coordination)
occur due to variation in high spin - low spin
occur due to changes associated wth the John-Teller distortion (Cu2+ ~ John-Teller ion)
changes to stability constant when structural change occurs at metal
stability constant for Cd with Br-
-logK1 = 1.58
-logK2 = 0.54
-logK3 = 0.06
-logK4 = 0.37
increase in 4th stability constant - this anomaly suggests a structural change at the metal
[Cd(H2O)6]2+ is a 6-coordinate complex but Br- is a larger ligand than H2O
stability constant K4 relates to :
[CdBr3(OH)3]- + Br- ⇌ [CdBr4] + 3H2O
OCTAHEDRAL ⇌ TETRAHEDRAL
addition of 4th Br- ligand is favourable (increase in entropy ~ 2 to 4 molecules) - as Br- is so large, there is not enough space for H2O molecules when 4th ligand is added
the chelate effect
K1 for a bidentate ligand (en) compared to β2 of corresponding di-ligand complex (diamine)
logK1 > logβ2
essentially, the same two Cu-N bonds are formed ~ but the chelated complex is more favoured
ΔS is large and positive for the en addition
ΔG more negative
reattachment of chelate ring
loss of one arm of the chelate ring - can easily reattach due to high local concentration of the end of the molecule
reattachment best for smaller ring sizes
chelate effect maximised for smaller ring sizes
chelate effect maxmised for ring sizes of 5 and 6 and it virtually non existant above a ring size of 9 atoms
macrocylic ligand
ligand that has a ring or a caged structure
macrocyclic effect
the enhanced stability of complexes with macrocyclic ligands compared to their open chain analogues
arises from changes in configurational entropy
- the open chain analogue loses a lot of configurational entropy (rotation around bonds) on complexation
- the macrocyclic ligand loses less configurational entropy as it is already in ring form
there is a cavity size effect for macrocyclic ligands
K+ fits best in the cavity for 18 crown 6
Rb+ fits best in cavity for valimaycin
the size selection properties of macrocyclic ligands can be used in the separation of various metal ions
macrocylic ligands that are three dimensional are known as cryptands
solubilities of ionic substances
solubility of ionic salts related to four main factors
relating to change in ΔH
1. lattice energy lose on dissolution
2. solvation energy gain on dissolving
relating to change in ΔS
3. the entropy gain by the ions moving in solution
4. the entropy lost by the solvent molecules
solubility of NaCl
Na+ and Cl- massively gain entropy
- there is very little comparative entropy in the rigid NaCl lattice
- on dissolvingm the anion and cation are free to move in solution
H2O molecules lose entropy
- attracted to ions
solvation energy is given by:
the solvation energies and solubilities increase with the dielectric constant of the solvents (dielectric constant - measure of how easy it is to sepatate ions)
the solubilities of the salts generally increase with size of cation and anion
-entropy change accompanying dissolution is less favourable for small ions
solvation energy inverserly proportional to ionic radius
lattice energy is given by:
lattice energy inversely proportional to the ionic radius
there is a charged effect - the lattice energy increases faster for multiply charged ions
lattice energy is higher for ions of similar size as they pack well in the lattice - salts are likely to be more soluble if the cation and anion differ in size
lattice energy higher for ions of similar size - less soluble salt
NaF 1.0 mol dm-3 (less soluble)
NaCl 6.2 mol dm-3
NaBr 8.8 mol dm-3
NaI 11.9 mol dm-3 (more soluble)
Na+ 0.95 Å
F- 1.36 Å
Cl- 1.81 Å
Br- 1.95 Å
I- 2.21 Å
- Na+ is closer in size to F-
- NaF has good packing in a lattice and high lattice energy
- much more difficult to dissolve
- in NaI, there is a big size difference between the anion and cation
- low lattice energy, less energy is needed to dissolve NaI
multiply charged ions are less souble than singly charged ions
require more solvent ions to go around them
NaI (singly charged) more soluble than MgSO4 (muliply charged)
oxidation and reduction
oxidised: species lost an electron
A → A+ + e-
reduced: species gained an electron
A + e- → A-
oxidising agent: species that removed electrons for oxidation
reducing agent: species that supplies electrons for reduction
REDOX REACTION:
Zn + 2H+ → Zn+ + H2
can be viewed as the sum of two half reactions - each with its own ΔG
2H+ 2e- → H2
Zn → Zn2+ + 2e-
ΔG from half reactions
ΔG = -nFE
ΔG - change in Gibbs Free energy
n - number of electrons transferred in reaction
F - Fariday constant
E - standard reduction potential
if E > 0, ΔG will be negative
for reaction to be spontaneous, E must be positive
Latimer diagram
applies thermodynamics to solution inorganic chemistry. it is in the form:
Ox → Red
enables us to calculate a non-adjacent couple ~ for example the E value for:
Cu2+ → Cu
Latimer diagrams are normally shown for pH = 0, but they are also tabulated for pH = 14
Latimer diagram (example)
(number of electrons) x (potential E for stepwise reduction) = (number of electrons transferred) x (potential of overall reaction)
Latimer diagram - strongest oxidising agent
look at the highest potential for any couple ~ oxidising agent is on the LHS of the couple
Latimer diagram - strongest reducing agent
look at the lowest potential for any couple
~ reducing agent is on the RHS of the couple
Latimer diagram - species likely to undergo disproportionation
M2+ → M+ + M3+
species goes from one oxidation state to a higher and lower oxidation state
a species is unstable to disproportionation if the potential on the left is lower than the potential on the right
Latimer diagram - species likely to undergo comproportionation
M+ + M3+ → M2+
a higher and lower oxidation state converge into a single species (opposite of disproportionation)
Frost diagram
pictorial representation of a Latimer diagram (constructed from Latimer diagram)
plot of electropotential against oxidation state
tells us about:
-comproportionation/disproportionation
-most thermodynamically stable species
-strongest oxidising/reducing agent
Frost diagram - comproportionation/disproportionation
Frost diagram - most thermodynamically stable species
species which is lowest on the diagram
Frost diagram - strongest oxidising/reducing agent
strongest oxidising agent: species which lies at the top of the most positive slope
- [O] agent is on the RHS of the most positive slope
strongest reducing agent: species which lies at the top of the most negative slope
- [R] agent is on the LHS of the most negative slope
Frost and Latimer diagrams
tell us about solution thermodynamics and not about kinetics
but an overpotential of 0.6V normally means the reaction occurs kinetically fast
overpotential for oxidation: 0.6V greater than that required for the minimum thermodynamic amount
overpotential for reduction: 0.6V less than than required for the minimum thermodynamic amount
stability field of water
region 1: any oxidising agent can oxidise water rapidly
- will kinetically and thermodynamically oxidise water
region 2: any species can thermodynamically oxidise water, but the reaction is kinetically slow
region 3: water is stable to oxidation and reduction
region 4: species can thermodynamically reduce water, but the reaction is kinetically slow
region 5: any reducing agent can reduce water rapidly
- will kinetically and thermodynamically reduce water
oxidation line relates to oxidising water. lines slope downwards to higher pH
- easier to oxidise species at high pH
- easier to reduce species at low pH
thermodynamic and kinetic oxidation of water
thermodynamic oxidation
- at pH=0, a species with a potential greater than 1.23V required
- at pH=10, a species with a potential greater than 0.63V required
kinetic oxidation
an overpotential of 0.6V is required for reaction to proceed quickly
- at pH=0, a species with a potential greater than 1.83V
- at pH=10, a species with a potential greater than 1.23V required
thermodynamic and kinetic reduction of water
thermodynamic reduction
- at pH=0, species with a potential less than 0V required
- at pH=10 species with a potential less than -0.6V required
kinetic reduction
an overpotential of 0.6V is required for reaction to proceed quickly
- at pH=0, species with a potential of less than -0.6V required
- at pH=10, species with a potential of less than -1.2V required
Pourboix diagram
phase diagrams which plot electropotential against pH
produce a form of phase diagram that enable us to predict the form of a species at a particular potential and pH value
Pourboix diagram example
strongest oxidising agent: at the top of the diagram ~ FeO4 2-
strongest reducing agent: at the bottom of the diagram ~ Fe
areas in the Pourbaix diagram
mark regions where a single species is stable
- more stable species tend to occupy larger regions
lines in the Pourbaix diagram
mark places where two species exist in equilibrium
horizontal lines: pure REDOX reactions
- reactions are not pH-dependent
- no acid-base equilibria
verticle lines: pure acid-base reactions
- reactions do not depend on potential
downward slope of -0.0592 V/pH: reactions that are both acid-base and redox
- acid/base exchange is an important part of process
acidity of cations in solution
dissolution of ionic salts in water can cause changes to solution pH
- due to metal ion polarising the water molecules and producing H+ in solution
smaller, multiply charged ions tend to cause more distortion of the water molecules (related to the charge to radius ratio of the ion)
pH
-log[H+]
lower pKa value more acidic
pH ~ Henderson Hasselbalch equation
pH = pKa + log10 ([base]/[acid])
metal is able to polarise the M-O bond towards itself
the bond gets elongated and easier to break dependent on the q/r ratio of the cation
ions of high charge/small radius are more polarising - more chance of breaking the O-H bond and forming H+ in solution
lower pKa values indicates stronger acids
metal aquo ions act as Bronsted acids ~ H+ donors
metal aqua ions - acidity
- smaller ions have higher acidity
- higher charged metal ions have higher acidity
- electronegativity of ion also greatly affects pKa and can override other effects
most electronegative metal: gold
