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Transition Metal Properties

Key Concepts

Transition metals are located in the middle of the Periodic Table and have an electron configuration filling the d-subshell .

s subshell d subshell p subshell
H He  
Li Be   B C N O F Ne
Na Mg   Al Si P S Cl Ar
K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr
Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe
Cs Ba La Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn
Fr Ra Transition Metals

In comparison with main group metals, transition metals generally show:

Hardness, Density, Melting and Boiling Point

  Transition Metals Non-Transition Metals
Symbol V Cr Mn Fe Co Cu Na Mg Al K Ca Ba
Atomic Number
(Z)
23 24 25 26 27 29 11 12 13 19 20 56
Valence Shell
Electron Configuration
3d34s2 3d54s1 3d54s2 3d64s2 3d74s2 3d104s1 3s1 3s2 3s23p1 4s1 4s2 6s2
Density
(g cm-3)
6.1 7.2 7.4 7.9 8.9 8.9 0.97 1.7 2.7 0.86 1.6 3.5
Melting Point
(oC)
1900 1900 1250 1540 1490 1083 98 650 660 64 838 714
Boiling Point
(oC)
3450 2642 2100 3000 2900 2600 892 1110 2450 770 1490 1640
  • Transition metals have smaller atomic volumes than Group 1 and 2 metals because additional electrons are being progressively added to the inner atomic orbitals resulting in stronger attraction to the nucleus.

  • These atoms of smaller volume can pack together more closely resulting in higher densities and hardness.

  • Closer packing results in stronger bonding so more energy is required to melt or boil transition metals.

Ionisation Energy and Chemical Reactivity

  Transition Metals Non-Transition Metals
Symbol V Cr Mn Fe Co Cu Na Mg Al K Ca Ba
Atomic Number
(Z)
23 24 25 26 27 29 11 12 13 19 20 56
Valence Shell
Electron Configuration
3d34s2 3d54s1 3d54s2 3d64s2 3d74s2 3d104s1 3s1 3s2 3s23p1 4s1 4s2 6s2
First
Ionisation
Energy
(kJ mol-1)
656 659 724 766 764 752 502 744 584 425 596 509
  • The smaller atomic radii of transition metals means the valence shell (outer-shell) electrons are more strongly attracted to the nucleus and therefore less easily removed resulting in higher first ionisation energies compared to Group 1 and 2 metals.

  • Because electrons are less easily lost, the transition metals are less chemically active than Group 1 and 2 metals.

  • The lower chemical reactivity of the transition metals means they will be placed lower down in the activity series of metals compared to Group 1 and 2 metals.

  • Since oxidation relates to the loss of electrons, transition metals are less easily oxidised than Group 1 and 2 metals.

Oxidation States

  Transition Metals Non-Transition Metals
Symbol V Cr Mn Fe Co Cu Na Mg Al K Ca Ba
Atomic Number
(Z)
23 24 25 26 27 29 11 12 13 19 20 56
Valence Shell
Electron Configuration
3d34s2 3d54s1 3d54s2 3d64s2 3d74s2 3d104s1 3s1 3s2 3s23p1 4s1 4s2 6s2
Common Oxidation
States
+2
+3
+4
+5
+2
+3
+6
+2
+3
+4
+6
+7
+2
+3
+2
+3
+1
+2
+1 +2 +3 +1 +2 +2
  • The energies of the 3d and 4s orbitals are very close.

  • Often the lowest oxidation is +2 corresponding to the loss of 2 ns orbital electrons.

  • Higher oxidation states correspond to the additional loss of (n-1)d orbital electrons.

  • The decrease in maximum states after manganese in the first transition metal series (and after iridium in the second series and osmium in the third series) reflects the difficulty of breaking into a half-filled d subshell.

Coloured Compounds

  Transition Metals Non-Transition Metals
Symbol V Cr Mn Fe Co Cu Na Mg Al K Ca Ba
Atomic Number
(Z)
23 24 25 26 27 29 11 12 13 19 20 56
Valence Shell
Electron Configuration
3d34s2 3d54s1 3d54s2 3d64s2 3d74s2 3d104s1 3s1 3s2 3s23p1 4s1 4s2 6s2
Common Oxidation States +2
+3
+4
+5
+2
+3
+6
+2
+3
+4
+6
+7
+2
+3
+2
+3
+1
+2
+1 +2 +3 +1 +2 +2
Colour of Chloride Compound VCl2
green
CrCl3
red
MnCl2
pink
FeCl2
yellow
CoCl2
blue
CuCl2
yellow
NaCl white MgCl2 white AlCl3 white KCl
white
CaCl2 white BaCl2 white
Colour of Aqueous Solution (Mn+) V2+
violet
Cr2+
blue
Mn2+
pink
Fe2+
green
Co2+
pink
Cu2+
blue
Na+ colour-
less
Mg2+ colour-
less
Al3+ colour-
less
K+ colour-
less
Ca2+ colour-
less
Ba2+ colour-
less
  • A substance will appear coloured if it absorbs light from some portion of the visible spectrum.

  • Ions with d orbital electrons appear coloured because energy from visible light is absorbed and used to promote a d orbital electron to a higher energy d sublevel (referred to as d-d transitions).

  • Ions with no d orbital electrons are colourless, eg, Sc3+, Ti2+.

  • Ions with d10 electron configurations are colourless, d-d transitions are impossible because the d orbitals are all filled, eg, Zn2+

Paramagnetism

  Transition Metals Non-Transition Metals
Symbol V Cr Mn Fe Co Cu Na Mg Al K Ca Ba
Atomic Number
(Z)
23 24 25 26 27 29 11 12 13 19 20 56
Valence Shell
Electron Configuration
3d34s2 3d54s1 3d54s2 3d64s2 3d74s2 3d104s1 3s1 3s2 3s23p1 4s1 4s2 6s2
Paramagnetism of Aqueous Ions Yes
V3+
Yes
Cr2+
Cr3+
Yes
Mn2+
Mn3+
Yes
Fe2+
Fe3+
Yes
Co2+
Yes
Cu2+
No
Na+
No
Mg2+
No
Al3+
No
K+
No
Ca2+
No
Ba2+
  • Paramagnetism is a weak attraction into a magnetic field.

  • Substances with unpaired electrons can be paramagnetic.

  • Paramagnetism is caused by both the orbital and spin motions of electrons (any rotating or revolving charged object generates a magnetic field).

  • The magnetic fields of paired electrons cancel out, so only unpaired electrons contribute to paramagnetism.

Ferromagnetism

  • Ferromagnetism is a strong attraction into a magnetic field.

  • Ferromagnetism occurs when atoms with unpaired electron spins are just the right distance apart to permit the individual spins to align with each other within a relatively large region.
    The individual spins within this region act cooperatively resulting in a large magnetic effect.

  • Only solids can show ferromagnetism.

  • The only ferromagnetic elements at room temperature are iron, cobalt and nickel.

  • Ferromagnetic compounds such as CrO2 and Fe3O4 also exist.


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