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 Transition
Elements
and their Properties
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IONISATION ENERGIES
OXIDATION STATES
REACTION WITH WATER
REACTION WITH ACIDS
OXIDES
HYDROXIDES
COMPLEX FORMATION
COLOR OF COMPLEXES
CATALYTIC PROPERTIES
Trends in Ionisation Energies
The first, second and third ionization enthalpies of dblock metals are given below :
Sc
Ti
V
Cr
Mn
Fe
Co
Ni
Cu
Zn
760 760
740
750
910
First IE kJmol-1
630 660 650 650
Second I.E. kJmol-1
720
1240 1310 1410 1590 1510 1560 1640 1750 1960 1700
Third I.E. kJmol-1
2390 2650 2870 2990 3260 2960 3230 3390 3560 3800
（a The first ionization enthalpy of zinc is
exceptionally greater than that of copper. The
electronic configuration of zinc is probably more
stable in having full 3d and 4s subshells.
b）The second ionization enthalpy of Chromium is
slightly greater than that of its preceding and
succeeding neighbours. The second ionization
enthalpy involves the removal of an electron from
a half-filled subshell (3d5), which has extra stability :
The case is similar for copper which possesses a full
3d subshell :
（c）The third ionization enthalpy of manganese is
greater than its preceding and subsequent
neighbours.
The third ionization enthalpy involves the removal
of an electron from a half-filled subshell (3d5),
which has extra stability .
Variable oxidation states
d-block elements have electrons of similar energy
in both 3d and 4s subshells. Thus one
particular element can form ions of roughly the
same stability by losing different number of
electrons.
Transition elements from titanium to copper
exhibit two or more oxidation states in their
compounds.
Sc Ti
V
3
4- 1 5-1
3 4,3 5,3
Cr
Mn
6-1 7-1
6,3 7,4,2
Fe Co
6-1 5-1
3,2 3,2
Ni
4-1
2
Cu
3-1
2,1
Generalizations :
（a）The common oxidation states for each element include
+2 or +3 or both. +3 states are relatively more common at
the beginning of the series, whereas +2 states are more
common towards the end.
（b）The highest oxidation states up to manganese
correspond to the involvement of all the electron outside the
argon core : 4 for Ti, 5 for V, 6 for Cr and 7 for Mn. After
this, the increasing nuclear charge binds the d electrons
more strongly and so one of the common oxidation states is
that which involves the weakly held in the outer 4s shell
only : 2 for Fe, 2 for Co, 2 for Ni and 1 for Cu.
Zn
2
2
（c）Ti,
V, Cr and Mn never form simple ions in their
highest oxidation state since this would result in ions of
extremely high charge density. Hence the compounds of
these elements in which they exhibit their highest
oxidation state are either covalently bonded or contain
complex ions, e.g. VO3-,CrO3, Cr2O72-, Mn2O7, MnO4（d）The stability of Mn(II) and Fe(III) can be explained
by the fact that both of their electronic structure is 3d5.
This half-filled subshell in these ions has a particular
stability.
The Cu(I) ion with its full 3d10 electronic structure, is
less stable than the Cu(II) ion (3d9) in aqueous solution.
•A rule of thumb is that oxidation states are stable when
a d orbital is empty (Cr(VI)), half full 3d5 Fe3+, or full 3d10
Cu1+. The reason is that double occupation of a sub
orbital causes electron repulsion increasing the energy
of that configuration
•In all oxidation states above 3+ there is no doubt that
the bonding in compounds and complexes is
predominantly covalent. This arises because small 3+ ions
would be very polarising; rather than metals donating
electrons to a non-metal in an ionic bond electrons are
shared in a polar covalent bond.
• For example Fe(III) chloride FeCl3 has a melting point of
only 306oC - it sublimes when it forms by heating iron with
chlorine. By comparison NaCl has a melting point of 806oC,
and iron (II) chloride 672oC: here the bonds are highly ionic,
some covalent character in FeCl2, but the bond in FeCl3(s) is
polar but mostly covalent.
In all reactions the metals may lose 2 outer electrons to
form M2+ ions with non-Noble gas electron arrangements.
Transition metals are not as powerful reducing
agents as Group II - their Eo values are much more
positive. The metals are thus much more resistant
to oxidation by water and oxygen, which makes
them good materials for engineering purposes.
• Other oxidation states occur because the loss of
or gain of an electron allows compounds to form
with ionic or covalent bonds that are stable relative
to the elements.
In the complex anions MnO4-(VII) and
Cr2O72- (VI) the bonds are covalent. Here the
oxidation numbers show the number of
electrons the metal is using in bonding 6 and
7 respectively. As the d block are in the 4th
period the number of electrons allowed in the
outer shell is 18. The p shell is unoccupied
and can accommodate electrons from
oxygen. The bonding about Mn and Cr is
tetrahedral, but in dichromate the central
oxygen links two tetrahedra.
•Reaction of the d block with water - a redox reaction - to
form metal hydroxides
The d block elements, unlike group I elements and Group II
elements, react very slowly if at all with water to form hydrogen and
metal hydroxides. Even those with negative Eo values , do so only
slowly due to the oxide coat protecting the bare metal.
•Copper does not react with water or steam, because it is too
unreactive - positive Eo.
•Cr when red hot will react with steam to form Cr(III)oxide; an oxide
coat renders it passive to reaction in cold water. It is thus alloyed
with Fe to make stainless steel.
•Managanese and Iron in the presence of oxygen and water rusts
slowly to form hydrated oxide. This is a major problem with the use
of iron and steel in engineering projects e.g. car bodies. The reaction
is very slow at RTP and takes days to even start.
Fe (s) + O2(g) + (x +1)H2O (l)

Fe2O3.xH2O(s) + H2 (g)
•Reaction of the metals with acids - a redox
reaction
•Copper does not react with dilute acids, a positive Eo
means it is below hydrogen in the reactivity series.
• Some of the metals react slowly with dilute acids.
•Zinc is the only one that does so quickly
•Zn (s)+
2HCl(aq)

H2 (g) + ZnCl2 (aq)
•0
2(+1)
2 0
2+
•Iron and manganese react relatively slowly to give iron
(II) and Mn(II) salts
•e.g. Fe (s) +2HNO3(aq)

H2 (g) + Fe(NO3)2 (aq)
•All the metals go into solution in concentrated nitric
acid, without evolution of hydrogen - this is a redox
reaction.
•d blockOxides
•d block metals react exothermically with oxygen to form simple ionic
oxides. The formulae vary. These have giant structures and high melting
points because the ionic bond is very strong.
•Because of their high lattice energy the oxides are insoluble in water and
alkalis, but will dissolve in acids - i.e. they are insoluble bases.
Chromium
Managanese
Iron
Cr(III) - Cr2O3
Mn(II) - MnO, Mn(III) - Mn2O3; Mn(IV) - MnO2 black
Fe(II) - FeO; Fe(III) - Fe2O3 - rust coloured most stable.
Fe3O4 is formed by heating in air, this a mixture of FeO.Fe2O3
Copper
Cu(I)-Cu2O red brown & most stable; Cu(II) - CuO black
Zinc (3d10)
ZnO white when cold yellow when hot (loses some oxygen)
e.g.
CuO(s) + H2SO4(aq)------ CuSO4(aq) + H2O(l)
Fe2O3 (s) + 6HNO3(aq)---- 2Fe(NO3)3(aq) + 3H2O(l)
d block hydroxides are insoluble and can all be precipitated with
sodium hydroxide
Some of the metal hydroxides are amphoteric and will dissolve in excess
alkali to give a complex anion. All metal hydroxides will of course dissolve
in acid to give a salt solution .
a) Copper Cu2+(aq) in CuCl2(aq) Add dilute sodium hydroxide
solution:; a pale blue precipitate forms which does not dissolve in
excess.
b)
Cu2+(aq) +2OH-(aq)
blue soln.
Cu(OH)2(s)
pale blue ppt.
b) add ammonia solution drop by drop to CuCl2(aq); a pale blue precipitate of
copper hydroxide forms initially; the pale blue ppt of Cu(II) hydroxide forms as
above; then this dissolves to form the tetrammine complex which is deep
blue/purple. Note ammonia solution is both an alkali ( contains OH- ions) and a
ligand ammonia (Lewis base).
Cu(OH)2(s) + 4 NH3(aq)
---------
[Cu(NH3)4]2+(aq) + 2 OH-(aq)
deep blue/purple
. Iron Iron(II) Fe 2+e.g. in FeCl2 aq - pale green in solution
•a) Add dilute sodium hydroxide solution; a dark green gelatinous
precipitate that does not dissolve in excess NaOH. This is used as a
diagonostic test for Fe(II). This green precipitate goes rusty brown in
about 10 minutes; it is oxidised by oxygen in the air to Fe3+.
•
FeCl2(aq ) + 2NaOH(aq)--------- Fe(OH)2 (s) + 2NaCl(aq)
Fe itself and all iron compounds end up as Fe3+ in moist air, other
oxidising agents like chlorine, conc. HNO3 and H2O2 also carry out this
oxidation.This is why there are many more Bronze Age artifacts than the
younger Iron Age artifacts
3. Iron Iron(III) Fe 3+ e.g. in FeCl3 aq or s, often pale yellow in solution
and red brown in solids
Add dilute sodium hydroxide solution; a red brown gelatinous precipitate
forms that does not dissolve in excess NaOH. This is used as a
diagonostic test for Fe(III).
FeCl3(aq ) + 3NaOH(aq) ------------
Fe(OH)3 (s) + 3NaCl(aq)
yellow
red brown ppt.
4. Chromium Chromium(III) Cr3+ e.g. in CrCl3(aq)
Add dilute sodium hydroxide solution; a green precipitate formed
that does dissolve in excess NaOH. This is th esme equation as for
Al3+ ions.
[Cr(H2O)6]3+ (aq) + 3 OH- (aq) < ------ [Cr(H2O)3(OH)3](s) + 3 H2O(l)
H+(aq)
Chromium hydroxide is amphoteric - dissolves in alkali and acid
Formation of coloured complex ions.
•a LIGAND is a neutral molecule or negative ion which can
datively covalently bond with a metal or metal cation. Ligands are
Lewis bases and nucleophiles e.g. :NH3 and :OH•Ligands may be monodentate, bidentate, tetradentate and
hexadentate and more. This means that a ligand has 1,2, 4 or
6 atoms in the molecule or ions which can attach via a dative
covalent bond to a metal
A COMPLEX AND COMPLEX ION. A complex results when 1 or
more ligands form a larger molecule or ion with a metal or metal
cation. The oxidation numbers are conserved. The resulting complex
may be positive and is called cationic complex ion, netural - a neutral
complex, or negatively charged - an anionic complex ion. Examples
are given below.
a cation is attracted to the cathode(-) and is therefore postively
charged e.g [Cr(NH3)6]3+
an anion is attracted to the anode(+) and is therefore negatively
charged e.g. [CuCl4]2-
a neutral complex occurs when the charge on the metal is
exactly balanced by the charges on the ligand. e.g .i)
Fe(OH)2(s) is really [Fe(H2O)4(OH)2]o 0 because the 2+
charge on Fe is balanced by the 2 - charges on each OH; it
is a complex because 2 water ligands have been displaced
by 2 stronger hydroxide ligands.ii) the Ph(OH)CO2- is a
bidentate uninegative ligand; it complexes with Cu2+ to
give a neutral complex, since the 2 ligands bring in a
charge of 2- [Cu(Ph(OH)CO2)2]o this square planar and is
called bis2-hydroxybenzoatecopper(II) - the bis merely
means 2 ligands.
Colour of complex ions
Colour in an octahedral complex
•When 6 ligands surround a transition metal the ligand lone pair of
electrons interact differently with the different d sub-orbitals.
•The 5 d sub-orbitals split into 2 sets: an upper set of 2 orbitals
and a lower set of 3 sub-orbitals. The available number of d
electrons fill these orbitals starting witht the lower set.
•If there is a gap in the upper set of orbitals visible light will excite
an electron from a lower to upper orbital. This light will be removed
from the incident light and colour will be seen. If for example red
light is absorbed the complex will be blue.
complexes with 3d10arrangements are colourless because
their is no gap in the upper set of sub-orbitals e.g. Cu(I) ,
Zn(II) and Ag(I) compounds are colourless.
Catalytic properties of transition metals and
their compounds
Transition metals and their compounds are important
catalysts in industry and in biological systems.
Some of the transition metals including copper,
manganese, iron, cobalt, nickel and chromium are
essential for the effective catalytic activity of various
enzymes. One of the most important enzymes
containing copper is cytochrome oxidase. This enzyme
is involved in the process energy is obtained from the
oxidation of food. In the absence of copper, cytochrome
oxidase is completely inhibited and
the animal or plant is unable to metabolize food
effectively.
Numerous transition metals and their
compounds are important industrial catalysts.
Transition metal and their compounds can
catalyse reactions because they are able to
introduce an entirely new reaction mechanism
with a lower activation energy than the
un catalysed reaction. Since the activation
energy of the catalysed reaction is lower, the
reaction rate is faster.
Chemist believe that the catalytic activity of
transition metals and their compounds depends
on their ability to exist in various oxidation
states.
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