12. Introduction to the Chemistry of Transition Elements A2 Flashcards
Transition elements
d-block elements that form one or more stable ions with incomplete d orbitals
5 physical properties of transition elements compared to calcium (a typical s-block element)
- Higher melting points
- more delocalised electrons, stronger metallic bonds - Higher densities
- larger Ar & smaller atomic radius
- closer packing (face-centred cubic structure) *body-centred cubic structure for s-block elements - Smaller atomic radii & smaller ionic radii
- more protons in nucleus, stronger nuclear attraction - Higher 1st ionisation energies
- more protons in nucleus, stronger nuclear attraction - Higher electrical conductivity
- more delocalised electrons (both 3d and 4s electrons can be delocalised, due to similar energy levels)
Tendency of transition elements to have variable oxidation states
- TE can form ions of about the same stability by losing diff. no. of electrons, due to the similarity in energy of 4s and 3d electrons
- the likely oxidation states of a transition element can be predicted from a given electronic configuration
- only 4s and 3d electrons are lost in forming ions
Ligand
including monodentate, bidentate, polydentate ligands {+ examples}
A species that contains a lone pair of electrons that forms a dative covalent bond to a central metal atom/ion
Monodentate eg. H2O, NH3, Cl-, OH-, CN-
- donates 1 lone pair to central metal atom/ion
Bidentate eg. 1,2-diaminoethane (en), ethanedioate ion (ox)
- contains 2 lone pairs that each form a dative bond
Polydentate eg. EDTA^4- (hexadentate)
Complex
A molecule or ion formed by a central metal atom/ion surrounded by one or more ligands
Co-ordination number
The number of bonds formed to an atom/ion/species/metal
Reaction of copper(II) ions with ligands to form complexes
With:
- water molecules
- hydroxide ions
- ammonia molecules
- chloride ions (add conc. HCl)
Anhydrous CuSO4 solid = white
Hydrated CuSO4.5H2O solid = blue
- Cu(II) salts dissolve in water –> [Cu(H2O)6]2+ blue solution
* pale blue - [Cu(H2O)4(OH)2] pale blue ppt. forms, insoluble in excess NaOH(aq)
- acid-base reaction - [Cu(H2O)4(OH)2] pale blue ppt. forms, soluble in excess NH3(aq) to give a dark blue solution [Cu(NH3)4(H2O)2]2+
- ligand exchange reaction - [Cu(H2O)6]2+ blue solution changes to green then yellow solution [CuCl4]2-
- ligand exchange reaction, reversible
Reaction of cobalt(II) ions with ligands to form complexes
With:
- water molecules
- hydroxide ions
- ammonia molecules
- chloride ions (add conc. HCl)
Anhydrous CoCl2 solid = blue
Hydrated CoCl2∙6H2O = pink
- Cobalt(II) salts dissolve in water –> [Co(H2O)6]2+ pink solution
- [Co(H2O)4(OH)2] blue ppt. forms, turns pink/red on standing, insoluble in excess NaOH(aq)
- acid-base reaction - [Co(H2O)4(OH)2] blue ppt. forms, soluble in excess NH3(aq) to give a pale brown solution [Co(NH3)6]2+, turns darker brown on standing in air
- ligand exchange reaction - [Co(H2O)6]2+ pink solution changes to blue solution [CoCl4]2-
- ligand exchange reaction, reversible
Why is there little change in atomic/ionic radii and 1st ionisation energies across the period for transition elements?
Proton number and no. of electrons increase by 1 across the period.
Each additional electron enters the inner 3d subshell, providing more shielding effect between nucleus and outer 4s electrons.
Hence, effective nuclear charge remains fairly constant across period of TE.
4 chemical properties of transition elements
- Form coloured compounds
- Have variable oxidation states
- Form complexes
- Have catalytic activity
Splitting of degenerate d orbitals into two energy levels in octahedral and tetrahedral complexes
- In an isolated atom, the 5 d-orbitals have same energy (degenerate)
- In complexes, the presence of ligands splits the 5 orbitals into a group of 3 and a group of 2 (non-degenerate). Both groups have different energy levels & have an energy gap ΔE between them.
- Since ligands have LONE PAIRS of electrons,
- the (electrons in) d orbitals pointing towards the ligands have higher energy, due to greater repulsion
❖Octahedral complexes: 2 orbitals higher, 3 lower
❖Tetrahedral complexes: 3 orbitals higher, 2 lower
Origin of colour in transition element complexes
- Visible light range = 380nm to 700nm
❖ΔE = hf = hc/λ f = frequency of light absorbed h = Planck's constant c = speed of light λ = wavelength of light
- d-orbitals split into LOWER & UPPER orbitals (2 different groups) by ligands, with an energy gap between them and VACANCIES in the upper d-orbitals
- When light shines on complex, electrons ABSORB wavelength/frequency of VISIBLE light & are promoted from a lower to a higher energy level; known as d –> d transition
- Colour of the complex observed is the COMPLEMENT of the colour absorbed
- Colour depends on energy difference, ΔE, between the split d-orbitals: ΔE = hf
Effects of different ligands on absorption, and hence colour
- Energy gap, ΔE is different for the ligands
- Different frequency/wavelength of light ABSORBED
- Stronger ligands cause larger splitting of d-orbitals, more energy needed for d–>d transition
❖ Larger energy gap corresponds to shorter wavelengths
eg. [Cu(H2O)6]2+ pale blue, [Cu(NH3)4(H2O)2]2+ dark blue
- NH3 stronger ligand than H2O, larger splitting
- ΔE with NH3 ligands different from that with H2O ligands; with NH3 larger absorbance peak
hence colour change
Why do some compounds lack colour?
d –> d transition cannot occur!
Sc3+ (3d0): empty d subshell, no d electrons
Cu+, Zn2+ (3d10): full d subshell, no electrons can be promoted
Stability constant, Kstab
The equilibrium constant for the formation of a complex ion in a solvent from its constituent ions or molecules
