transition metals

transition metal

element that forms at least one ion that has a partially filled d sub-level

why is zinc not a transition element

the only ion it forms is the +2 ion, since Zn has fully filled d sub-level

physical and chemical properties of transition metals

• high density, mp/bp

• form compounds with different oxidation states

• form coloured ions and compounds

• used as catalysts: able to change oxidation states easily

• form complex ions and complexes with ligands

• paramagnetic: has unpaired electrons → magnetic moment depends on no. of unpaired electrons

• form slightly acidic salt solutions

atomic and ionic radius across the transition metals

• across the period, nuclear charge increases

• additional electrons are in the inner 3d sub-level, not valence shell → 3d electrons repel valence 4s electrons, shielding effect increases

• effective nuclear charge = nuclear charge - shielding effect → Zeff relatively constant

• therefore, atomic radii decreases

note: decrease is small compared to periods 2/3.

• ionic radius of M²⁺ ion is much smaller than atomic radius of M: 4s electrons removed, nuclear charge attracts remaining electrons more strongly

why physical properties of first row transition metals similar?

relatively small difference in Zeff

density across the transition metals

• density = mass/volume

• Ar increases → mass increases

• Zeff increases, atomic radii decreases → volume decreases

  • note: even though decrease in radii is small, volume is r³ so effect is magnified

• therefore, density increases

• note: d-block metals are much denser than s-block metals because d-block higher mass and smaller atomic radius

mp and bp across the transition metals

• metallic bonding affected by charge density of cation + no. of delocalised electrons

• across the period, atomic radii decreases, stronger electrostatic attraction of delocalised electrons to nucleus → charge density increases

• transition metals have both 3d and 4s electrons → no. of electrons for delocalisation increases

  • therefore, all elements have mp>1000C, much higher than s-block metals

• therefore stronger metallic bonds, mp/bp increases across the period

  • exception: Mn half-filled 3d sub-level very stable → reduces availability of valence electrons for delocalisation → weaker metallic bonding, less energy needed to overcome → lower mp

electronic configuration of transition metal ions

• d-block elements have outer pair of 4s electrons + partially/fully-filled 3d sub-level

• exception: copper 3d54s1 and chromium 3d104s1 since half-filled/full-filled d orbitals are stable

• 3d sub-level has lower principal quantum number n

• when 3d sub-level is being filled the electrons penetrate more efficiently than 4s electrons → 3d electrons shielding effect on 4s electrons → raises energy of 4s sub-level

• when atoms are ionised (removing electrons from 4s first), the 3d orbitals become much more stable → +2 and +3 oxidation states all have 3dⁿ configurations

explain why transition metals can act as homogenous catalysts

4s and 3d suborbitals have similar energy levels, readiness to lose electrons from either sublevel to form ions of variable oxidation states

IEs in transition metals

• 1st IEs generally same across the period because outer 4s electrons that are removed are shielded by inner 3d electrons

successive IEs

• 4s and 3d sub-levels similar energy level → no visible jump in successive IEs, gradual slope

• vs main group: large jump when change in shell and small jump when change in sub-level → energy released during lattice formation is much more than energy needed to form the ions → 2nd to 3rd IE visible jump

why do d-block metals in the middle of the row have high oxidation states?

• note: oxidation state doesn’t represent charge on atom

  • IE needed to produce such a high charge is much higher than available energy

• involves the formation of polar coordinate bonds with electronegative O ligands

• usually these are oxidising agents (eg MnO4-)

coordination number

number of coordinate bonds formed between the transition metal ion and the ligand

eg if coordination no. is 6, 6 ligands surround central metal cation, octahedral shape.

complex ion formation

central metal cation forms coordinate covalent bonds to a cluster of molecules/anions called ligands

• transition metals have energetically accessible 3d orbitals/4s/4p that are empty → can accept lone pair

• transition metals cation has high charge density → can attract lone pair

all first row transition elements (except Ti) form octahedral complex ion with formula [M(H₂O)₆]ⁿ⁺ (aq) → aqueous suggests surrounded by water molecules, which act as ligands by donating lone pair (:OH₂)

ligands

molecules/anions with at least one lone pair of electrons, which it uses to form coordinate covalent bond with central metal cation

note: if multiple atoms have lone pair, less electronegative atom will donate

why does V3+ form different coloured ions with H2O and CN-

• CN is a stronger field ligand (due to negative charge)

• different energy gap

• different wavelength absorbed

• different wavelength of complementary colour observed

colour of coloured complexes

1. d-d splitting: d orbitals split into 2 different sets of energy levels due to ability of ligands

2. d-d transition: d electron absorbs light from visible spectrum, promoted from lower energy levels to higher energy level

3. colour of complex is complementary to colour of light absorbed

if no colour: fully-filled d sub-level → no d-d transition → no absorption of energy

further explanation

1. d-d splitting

   • ligands (which hold electrons) enter along axes

   • orbitals are also holding electrons → therefore the 3d electrons nearer ligand will be repelled and have higher energy

   • orbitals in the same sub-level should have similar energy level, but because of spatial arrangement of the orbitals, extent of repulsion is different → split into 2 energy levels

     • head-on overlap when 3d orbitals are along x/y/z axes, hence experience greater repulsion → energy increases more, 2 orbitals destabilised

     • other 3d orbitals are between axes, experience less repulsion → energy increases but less, 3 orbitals stabilised

2. d-d transition

   • the difference causes energy gap (delta E): energy gap coincides with visible wavelength (use E=hf to calculate)

   • lower energy electron can gain energy to transition to higher energy sub-level. this transition absorbs light energy.

   • reflected light (colour of complex) is complementary to absorbed light.

note: regardless of splitting pattern, as long as not fully filled, will have colour. (ie d0 and d10 cations are colourless)

factors influencing energy gap (and therefore colour of complex)

• metal ion (not charge!): nuclear charge increase, attraction of d electrons to nucleus increase, strength of coordinate bond increase → electrons nearer, repulsion between d electrons and ligand electrons increase → energy gap larger

• number of d electrons increase, repulsion between d electrons and ligand electrons increase → more d-d splitting → energy gap larger

• nature of ligand: ligand charge density increase → repulsion between d electrons and ligand electrons increase → more d-d splitting → larger energy gap

• geometry of complex

  • octahedral → ligands enter along axes

  • tetrahedral → ligands enter between axes

  • repel differently, different splitting pattern

paramagnetic vs diamagnetic

paramagnetic has unpaired electrons, diamagnetic has no unpaired electrons

monodentate vs polydentate ligands

monodenate ligands form coordinate bonds through one donor atom

polydentate ligands (bidentate, hexadentate) form coordinate bonds through multiple donor atoms → gives rise to chelates (multiple coordinate bonds such that there is a ring structure. eg EDTA, very stable complex)

note: SCN- can have lone pair on either S or N → still monodentate since only one coordinate bonds through in the end

structure of complexes

• 2-coordinate: linear

• 4-coordinate: tetrahedral/square-planar

• 6-coordinate: octahedral

explain formation of complex ion using acid-base theory

1. lone pair of electrons on ligand act as electron pair donor

2. energetically accessible 3d orbitals on transition metal cation act as electron pair acceptor

3. coordinate covalent bond formed

why do most transition metals form 2+ cations?

loss of 4s2 electrons

Complex ion

Central metal cation is closely surrounded by molecules/anions with lone pair (ligands) that it uses to form coordinate bond with the metal cation