transition metals

coloured compounds, d orbital splitting, catalysis equations, autocatalysis, ligand substitution, vanadium, redox, oxidation states, pH, everything !!

defintions !!

defintions !!

characteristics of transition metals

• variable oxidation states

• catalytic activity

• forms coloured compounds

• forms complexes

how do these characteristics arise?

incomplete d subshell

co-ordinate bond (dative covalent)

where one atom donates both the electrons in the covalent bond

ligand

a molecule or ion which forms a co-ordinate bond with a transition metal by donating a pair of electrons

catalysis!!

adsorption

solid catalysts adsorbs molecules onto active site on surface of the catalyst

increases proxmitiy of molecules and weakens the covalent bonds

what can happen to heterogeneous catalysts? industrial/economic impact of this?

become poisoned by impurities which adsorb to the active site, preventing reactants from binding.

needs to be replaced, increasing costs as production may have to stop.

example of catalytic poisoning

sulfur impurities in methane during the haber process will adsorb to iron catalyst

what is the contact process

making SO₃ from SO₂ (involved in manufacturing sulphuric acid)

type of catalyst used in contact process

heterogeneous

catalyst used in contact process

Vanadium V - V₂O₅

contact process equations with catalyst

V₂O₅ + SO₂ → SO₃ + V₂O₄

V₂O₄ + 0.5O₂ → V₂O₅

number of activation energy profiles for homogeneous catalysts

2

Catalyst used in oxidation of iodide ions with persulfate ions and type of catalyst

Fe²⁺ in solution, homogeneous

Equations for oxidation of iodide ions with persulfate ions

S₂O₈²⁻ + 2Fe²⁺ → 2SO₄²⁻ + 2Fe³⁺

2Fe³⁺ + 2I⁻ → 2Fe²⁺ + I₂

autocatalysis example

Mn²⁺ formed from reaction of MnO₄⁻ with C₂O₄²⁻

MnO₄⁻ half equation

MnO₄⁻ + 8H⁺ + 5e⁻ → 4H₂O + Mn²⁺

C₂O₄²⁻ ethanedioate ion half equation

C₂O₄²⁻ → 2CO₂ + 2e⁻

Autocatalysis equations for C₂O₄²⁻ with MnO₄⁻

MnO₄⁻ + 4Mn²⁺ + 8H⁺ → 4H₂O + 5Mn³⁺

C₂O₄²⁻ + 2Mn³⁺ → 2Mn²⁺ + 2CO₂

Colour !!

How does colour arise with transition metals?

Some wavelengths are absorbed and the rest of the wavelengths are reflected or transmitted.

What happens when light is absorbed?

d electrons move from ground state to excited state when light is absorbed

Calculation for energy difference between ground state and excited state

∆E = hv = hc/λ

Frequency of light absorbed depends on…

∆E

Factors which affect ∆E

• Change in oxidation state

• Change in type of ligand

• Change in co-ordination number

How does type of ligand affect ∆E

different ligands will split the d orbital by a different amount of energy

ligands will cause different levels of repulsion to d orbital

how does co-ordination number influence ∆E

affects the strength of the metal ion-ligand interactions

technique used to find concentration of compounds using the colour of the ions

spectroscopy

principles of spectroscopy, colorimeter

• shine white light through a coloured filter ( complementary to colour of solution) aimed at the sample

• more light it absorbs, higher the concentration of the solution

• compare the amount of light absorbed to a calibration curve

ligand sub, shapes and complexes !!

example of incomplete substitution of ligands, and colour of this compound

[Cu(NH₃)₄(H₂O)₂]²⁺ deep blue

Size of Cl⁻ ligand compared to NH₃ and H₂O

Larger, forms compounds with co-ordination number 4 and tetrahedral shape

Bidentate ligands examples

C₂O₄²⁻ ethanedioate ion and H₂NCH₂CH₂NH₂ 1,2-diaminoethane

Shape of complexes bidentate ligands form and isomerism they exhibit

Octahedral and optical

Isomerism shown by octahedral complexes like [Cu(H₂O)₄Cl₂] and how to identify each one

cis-trans

trans isomer has the ligands opposite eachother

cis isomer has the ligands on the same same

Multidentate ligands examples

EDTA⁴⁻ and haem with iron(II)

How is oxygen transported in the blood

Oxygen forms a co-ordinate bond to Fe(II) in haemoglobin

How is CO toxic

forms a stronger co-ordinate bond with Fe(II) in haemoglobin, replacing oxygen

Chelate effect

Replacing monodentate ligands with bidentate or multidentate ligands, increases number of moles of products, negative ∆G means reaction is spontaneous and favourable

Example of square planar complex

Cisplatin

cisplatin is used in cancer treatment, what does it do?

• stops DNA replication

• ligand replacement reaction with DNA

• bond forms between nitrogen atom on guanine and platinum in cisplatin

Linear complex example and common use

[NH₃→Ag←NH₃]⁺ Tollens’ reagent for testing aldehydes

Vanadium chem !!

Vanadium species. different oxidation states and colours

Yellow VO₂⁺ +5

Blue VO²⁺ +4

Green V³⁺ +3

Violet V²⁺ +2

How are the IV, III, II oxidation states formed?

Reduction of VO₂⁺ using Zinc in acidic solution

Vanadium(V) —> Vanadium (IV) and colour change

2VO₂⁺_(aq) + Zn_(s) + 4H⁺_(aq) → Zn²⁺_(aq) + 2VO²⁺_(aq) + 2H₂O_(l)

_{Yellow to blue}

Vanadium IV to Vanadium III and colour change

2VO²⁺(aq) + Zn(s) + 4H⁺ → Zn²⁺(aq) + 2V³⁺(aq) + 2H₂O(l)

Blue to green

Vanadium III to Vanadium II and colour change

2V³⁺(aq) + Zn(s) → Zn²⁺(aq) + 2V²⁺(aq)

Green to violet

Redox potential

Show how easily a metal can be reduced (same as electrode potential)

Higher values mean ion is less stable and more easily reduced

Redox potentials depend on…

• Ligands

• pH (higher H⁺ concentration means a higher redox potential)

Colours of hexaaqua ion complexes

[Cu(H₂O)₆]²⁺

blue

[Fe(H₂O)₆]²⁺

green

[Fe(H₂O)₆]³⁺

yellow (purple)

[Al(H₂O)₆]³⁺

colourless solution

hexaaqua ions with OH⁻/NH₃ and colour of complex formed !!

[Cu(H₂O)₆]²⁺

[Cu(H₂O)₆]²⁺(aq) + 2OH⁻ (aq)→ [Cu(OH)₂(H₂O)₄](s)

blue ppt

[Fe(H₂O)₆]²⁺

[Fe(H₂O)₆]²⁺(aq)+ 2OH⁻(aq) → [Fe(OH)₂(H₂O)₄](s)

green ppt, brown ppt on standing (as 3+ is formed as it reacts with O₂ in the air)

[Fe(H₂O)₆]³⁺

[Fe(H₂O)₆]³⁺(aq)+ 3OH⁻(aq) → [Fe(OH)₃(H₂O)₃](s)

brown ppt

[Al(H₂O)₆]³⁺

[Al(H₂O)₆]³⁺(aq)+ 3OH⁻(aq) → [Al(OH)₃(H₂O)₃](s)

white ppt

with excess NH₃

[Cu(H₂O)₆](s) + 4NH₃ → [Cu(NH₃)₄(H₂O)₂](aq) + 4H₂O

deep blue SOLUTION

with excess OH⁻

[Al(OH)₃(H₂O)₃](s) + OH⁻ → [Al(OH)₄]⁻ + 3H₂O

colourless solution reformed

[Al(OH)₃(H₂O)₃](s) with H⁺

[Al(OH)₃(H₂O)₃](s) + 3H⁺ → [Al(H₂O)₆]³⁺

this shows [Al(OH)₃(H₂O)₃](s) is…

amphoteric - can act as both an acid and a base

reactions of hexaaqua ion complexes with CO₃²⁻…

2⁺ ions form…

XCO₃(s) ppt

3⁺ ions form…

same as OH⁻/NH₃ ppt, with CO₂ gas

colour of complex formed when [Cu(H₂O)₆]²⁺ reacts with CO₃²⁻

CuCO₃ is green/blue ppt

colour of complex formed when [Fe(H₂O)₆]²⁺ reacts with CO₃²⁻

FeCO₃ is green ppt

pH !!

Explain why an aqueous solution containing [Fe(H₂O)₆]³⁺ ions has a lower pH than an aqueous solution containing [Fe(H₂O)₆]²⁺ ions.

• [Fe(H₂O)₆]³⁺ has higher charge to size ratio/ higher charge density

• [Fe(H₂O)₆]³⁺ polarises the water molecules more

• More O-H bonds break, releasing more H⁺ ions.

Explain, with the use of an equation, why a solution containing [Al(H2O)6] 3+ has a pH ˂7

• [Al(H₂O)₆]³⁺ ⇌ [Al(H₂O)₅(OH)]²⁺ + H⁺

• [Al(H₂O)₆]³⁺ has a high charge density

• weakens the O-H bond, releasing H⁺ ions