Download TRANSITION ELEMENTS

TRANSITION ELEMENTS
Transition elements refer to those in the d-block of the Periodic Table. The sequence from scandium (Z = 21) to zinc (Z = 30)
form what is generally regarded as the first transition series. Virtually ALL the properties of transition elements are related to
their electronic structures, in particular the relative energy levels of the orbitals available for their electrons.
Electronic Configuration
Atomicr
Number
Element
Symbol
Electronic Configuration
Ground state
Most common oxidation state
21
22
23
24
25
26
27
28
29
30
1
As the shells of electrons get further away from the nucleus, successive shells become closer in energy. By the time the
fourth shell is reached, there is, in fact, an overlap between the third and fourth shells, making the energy of the 3d orbitals
much higher than that of the 4s orbitals.
2
Notice the unexpected electron structures for chromium and copper, as a result of the extra stability acquired from their halfor fully-filled subshells.
3
Once the 3d level is occupied by electrons, these repel the 4s electrons even further from the nucleus. The 4s electrons are
therefore pushed to a higher energy level (ie. 4s > 3d). Consequently, when transition metals form ions, they lose electrons
from the 4s level before the 3d level. This means that all transition metals will have similar chemical properties, which
are dictated by the behaviour of the 4s electrons in the outermost shell.
Characteristic Properties of Transition Elements and their Compounds
Characteristic properties
1
Variable Oxidation States
2
Formation of Coloured Compounds
3
Catalytic Property
4
Formation of Complex Ions
Typical examples
d-Block 2
Variable Oxidation States
Transition metals exhibit more than one oxidation states in their compounds because :
1
there is very little energy difference between 3d and 4s electrons; and
2
there is only a slight increase in the successive ionization energies as electrons are being removed progressively.
Consequently, they have electrons of similar energy in both the 3d and 4s levels. This means that one particular element can form
ions of roughly the same stability by losing different numbers of electrons, thus allowing the existence of more than one oxidation
state.
Sc
Ti
V
Cr
Mn
Fe
Co
Ni
Cu
Zn
7
3
6
6
6
5
5
5
5
5
4
4
4
4
4
4
4
3
3
3
3
3
3
3
3
2
2
2
2
2
2
2
2
1
1
1
1
1
1
1
1
2
The following generalizations emerge from a study of the oxidation states :
1
The common oxidation states for each element at the beginning of the series corresponds to the maximum oxidation
available, which obviously involves the loss of all 4s and 3d electrons to give a noble gas structure.
2
The common oxidation states for each element towards the end of the series corresponds is +2, which involves the loss of 4s
electrons only. As the nuclear charge increases progressively across the period, electrons are more firmly held, making the
removal of all 3d electrons practically impossible.
3
Elements rarely form simple ions with oxidation state greater than +3, since this would result in ions of extremely high
charge density. Compounds with higher oxidation states are usually found in oxides, fluorides and oxy-anions, all of which
are highly electronegative and less likely to be polarized.
4
Both scandium and zinc have only one oxidation state in their compounds. In fact, their properties differ so much from the
typical transition metal properties that they are usually not regarded as transition elements.
AN ELEMENT WITH AT LEAST ONE ION WITH A PARTIALLY FILLED d SUBSHELL IS DEFINED
AS A TRANSITION ELEMENT. IN OTHER WORDS, THE TYPICAL PROPERTIES OF TRANSITION
ELEMENTS ARE ASSOCIATED WITH PARTIALLY-FILLED d SUBSHELLS.
Physical Properties of the First Transition Series
Element
K
Ca
Sc
Ti
V
Cr
Mn
Fe
Co
Ni
Cu
Zn
Atomic radius / nm
0.24
0.20
0.16
0.15
0.14
0.13
0.14
0.13
0.13
0.13
0.13
0.13
Melting point / °C
64
850
1540
1680
1900
1890
1240
1540
1500
1450
1080
420
Boiling point / °C
770
1490
2730
3260
3400
2480
2100
3000
2900
2730
2600
910
Density / gcm-3
0.86
1.54
3.0
4.5
6.1
7.2
7.4
7.9
8.9
8.9
8.9
7.1
d-Block 3
1
The variation of atomic radius (hence I.E., electronegativity etc.) across a transition series is small, compared with
that in the main period (ie. from Na to Ar)
--
In building up elements in the main period, electrons are being added to the OUTER shell and the nuclear charge is
increasing by the addition of protons. The added electrons shield each other only weakly from the extra nuclear charge
(2° shielding effect), so atomic radii decrease sharply across a period.
--
In moving across the transition series, however, the nuclear charge is also increasing, but electrons are being added to an
INNER d sub-shell. The outer 4s electrons, which determine the atomic radii, are more effectively shielded from the
increasing nuclear charge (1° shielding effect). Consequently, the atomic radii decrease much less rapidly.
⇒ Transition metals usually form alloys by replacement of metal atoms of similar sizes with little distortion of the crystal
lattice (e.g. brass [Cu + Zn] ; stainless steel [Fe + Cr + Ni + C] )
2
Transition metals have higher melting points, higher boiling points and higher densities than s-block metals
--
Most of the transition metals have a close-packed structure in which each atom has 12 nearest neighbours. This,
together with their small atomic size, results in strong metallic bonds between atoms.
⇒ The strong interatomic bonding in transition metals is also reflected in their high tensile strengths and good mechanical
properties (e.g. use of titanium as spacecraft bodies and steel in construction purposes)
Chemical Properties
1
The small, highly charged cations cause greater polarization of associated anions, resulting in the following chemical
properties of transition metal compounds compared with those of the s-block metals :
their oxides and hydroxides in oxidation states +2 and +3 are less basic and less soluble;
their salts are less ionic and less thermally stable;
their salts and aqueous ions are more hydrated and more readily hydrolyzed forming acidic solutions;
their ions are more easily reduced than main group metal ions.
2
Most transition metals react slowly with dilute acids, as compared with s-block metals, because :
they are less electropositive than the s-block metals;
the presence of a thin, unreactive layer of oxide protects the metal from chemical attack (similar to Al )
Formation of Complex Ions
Complex ions are composed of a central metal ion surrounded by anions or molecules, called ligands. Lone-pair electrons on
the ligand form co-ordinate bonds (dative covalent bonds) to the central metal ion, which provides vacant orbitals to
accommodate these electron pairs. The number of co-ordinate bonds from ligands to the central ion is known as the coordination number of the central ion.
Ligand
Metal Ion
MLx
Co-ordination number
Most ligands can form only one co-ordinate bond with a central ion, and are
said to be monodentate ligands. In some cases, each ligand molecule or
anion can form more than one link with the metal ion. These are said to be
polydentate.
d-Block 4
H - N - CH2 - CH2 - N - H
H
H
ethane-1,2-diamine (en)
EDTA (ethylenediaminetetraacetate ion)
Bonding in Complex Ions
co-ordination
number
example
2*
[ Cu (NH3)2 ]+
hybridization of
central metal ion
shape of
complexes
[ Co Cl4 ]2-
4
[ Ni(CN)4]2-
6
*
[Fe(H2O)6]3+
Most transition metal ions with full-filled d10 configuration (e.g. Cu+, Ag+) co-ordinate with only 2 ligands to form linear
complex ions.
Nomenclature of Complex Ions
1
Name the ligands before the central metal ion.
2
Name ligands with negative charges first, and then neutral ones.
-ve charge ligands :
ClFneutral ligands
:
H2O
OH-
CN-
NH3
3
State the number of each type of ligands around the central ion using prefixes : di-, tri-, tetra-, penta-, hexa- etc.
4
Name the central ion using its English name directly in the case of a cationic complex; but ends with the suffix -ate for
anionic complex.
5
Al
Cr
Fe
Ni
Mn
Cu
Indicate the oxidation number of the central metal ion using Roman numerals inside brackets.
d-Block 5
Class Work
Name the following complexed compounds :
[Cu(H2O)6]SO4
Ag(NH3)2OH
[Co(NH3)6]Cl3
[Co(NH3)5Cl]Cl2
K4[Fe(CN)6]
K3[Fe(CN)6]
Stereo-structure of Complex Ions
(I) Co-ordination no. = 4
(a) Ma b
2 2
(b) Ma bc
2
(c) Mabcd
(II) Co-ordination no. = 6
(a) Ma b
2 4
(c) Ma b c
2 2 2
(b) Ma b
3 3
d-Block 6
(III) With bidentate ligands
(a) M(en) a
2 2
(b) M(en)
3
Stoichiometric Determination of Complex Ions
1
Elemental Analysis
Class Work
A cobalt complex Q, known to have water molecules and chloride ions as ligands, gives the following composition by mass:
Co, 21.5 % ; O, 35.1 % ; Cl, 39.0 %. Deduce the molecular formula of Q.
(Relative atomic mass : Co, 58.9 ; O, 16.0 ; Cl, 35.5)
2
Titration
--
to distinguish between free ions and covalently-bonded ligands (hence reveal the structural formula of the complex)
--
add excess aq. AgNO3 to known amount of complex; weigh the mass of AgCl ppt. so formed
⇒ 1 mol. of
produces
mol. of AgCl ppt.
⇒ 1 mol. of
produces
mol. of AgCl ppt.
⇒ 1 mol. of
produces
mol. of AgCl ppt.
Stability of Complex Ions
1
Role of water as a complexing agent
In aqueous solution, transition metal ions exist as hydrated complexes with water molecules. (When an ionic crystal
dissolves in water, lattice energy has to be overcome and this is compensated by the hydration of ions, ie. complexing with
water molecules)
Owing to the high charge density on the central metal ion, these hydrated complexes usually dissociate, making the aqueous
solutions of most transition metal compounds acidic (e.g. CuSO4 and FeCl3) :
[Fe(H2O)6]3+(aq)
2
[Fe(H2O)5OH]2+(aq) + H+(aq)
Stability constant
When a ligand stronger than water is added to an aqueous solution of a transition metal, water (the weaker ligand) would be
displaced and an equilibrium would take place. In general :
[ MLx]
M (H2O)x + x L
MLx + x H2O
Kst = --------------------[M(H2O)x] [L]x
where Kst is known as the stability constant, whose magnitude reflects the stability of the complex formed.
d-Block 7
3
Factors affecting stability constants
Kst values
Fe2+
Co2+
Ni2+
Cu2+
Zn2+
M (en)2+
9.5
13.8
18.1
18.6
12.1
M (EDTA)2-
14
16
18.5
19
16
(a) Generally speaking, hexadentate ligands form more stable complexes than bidentate ligands, and, in turn, more
stable than monodentate ones. This is essentially an entropy effect, which states that a larger increase in disorder is
always more favourable.
e.g.
[Cu(H2O)4]2+(aq) + 4 NH3(aq)
[Cu(NH3)4]2+(aq) + 4 H2O(l)
[Cu(H2O)4]2+(aq) + EDTA4-(aq)
[Cu(EDTA)]2-(aq) + 4 H2O(l)
(b) Metal ions of higher charge density usually form more stable complexes.
(c) The exceptionally lower tendency for zinc to form stable complex is attributed to its full-filled d10 configuration, which
acquires extra stability and has no vacant, low-lying d orbitals available to accommodate lone-pair e- from ligands.
Uses of Complex Ions
1
Detection of metal ion by the color of its complex, e.g. Fe3+ in [Fe(SCN)]2+ -- blood red , Cu2+ in [Cu(NH3)4]2+ -- deep blue
-- the complex must have an intense and characteristic colour so that the metal ion can be readily identified
2
Extraction of metals
e.g. aluminium and silver are extracted via the formation of soluble complexes AlF63- and Ag(CN)2-
3
Softener of hard water
-- when hard water, containing Ca2+ and Mg2+ ions, is mixed with soap, it reacts with the anions in these salts, producing insoluble compounds which we see
as 'scum' on the surface of water
-- various ligands (e.g. polyphosphate) can form complex with Ca2+ and Mg2+ ions, thus preventing them to react with the anions in soap and the water is said
to be 'softened'
4
Formation of essential biological marcomolecules, e.g. chlorophyll (complex of Mg2+) and haemoglobin (complex of Fe2+)
Notice that the chemistry of a complex ion is, in most cases, entirely different from that of the free metal ion.
Formation of Coloured Compounds
Most of the compounds and complexes of transition elements are coloured. The colour of these compounds can often be related
to incompletely filled d-orbitals in the transition metal ion.
The outer electronic orbitals of transition metal ions have only small energy difference and many of these ions have
unpaired electrons in the outermost shell. These unpaired electrons can absorb electromagnetic radiation and undergo
transition in the visible region and hence the ions are coloured.
Important points to note :
1
Compounds formed by ions with empty (d0 ) and full-filled (d10 ) configuration are colorless, because electronic transition
within 3d and 4s orbitals is impossible.
2
Any colour change accompanying a chemical reaction involving transition metal ions could be taken as the first sign that
either one of the following has taken place : (i) a change in oxidation state; or (ii) ligand displacement.
d-Block 8
(i) Common oxidizing and reducing agents :
(ii) Common complexing agents :
Oxidizing agents
Reducing agents
MnO4- → Mn2+
2 I- → I2
Mn+ → [ M(H2O)6 ]n+
Cr2O72- → 2Cr3+
SO2 → SO42-
Mn+ → [ M(OH)4 ](n-4)
Fe3+ → Fe2+
C2O42- → 2 CO2
Mn+ → [ M(NH3)4 ]n+
H2O2 + 2H+ + 2e- → 2H2O
H2O2 → O2 + 2H+ + 2e-
Mn+ → [ M Cl4 ](n-4)
F2 → 2 F Class Work
Explain, using equations, the following observations :
(a) When aq. NaOH is added to aq. iron(II) sulphate, green precipitate is obtained instantaneously, which rapidly turns brown on
exposure to air.
(b) On shaking solid copper(I) iodide with aq. ammonia, it dissolves to give a colorless solution, which gradually turns deep blue
on standing. Addition of conc. HCl to the resulting solution changes its colour to green.
Catalytic Property
Transition metals and their compounds are important catalysts in industry and in biological systems.
1
Heterogenous catalysis
N2(g) + 3 H2(g)
2 NH3(g)
Transition metals which possess vacant, low-lying 3d orbitals may act as electron acceptors to form temporary bonds
between catalyst and reactants, thus bringing reactant particles closer enough to result in successful reaction.
2
Homogenous catalysis
S2O82-(aq) + 2 I-(aq)
2 SO42-(aq) + I2(aq)
Variable oxidation states of transition elements allow electron transfer between reactants and products by means of the
catalyst interchanging between two oxidation states, thus facilitating an oxidation-reduction cycle.