Transition metals

"The transition elements are those elements having a partially filled d or f subshell in any common oxidation state."

The term "transition elements" most commonly refers to the d-block transition elements. The 2B elements zinc, cadmium and mercury do not strictly meet the defining properties, but are usually included with the transition elements because of their similar properties. The f-block transition elements are sometimes known as "inner transition elements". The first row of them is called the lanthanides or rare earths. The second row consists of the actinides. All of the actinides are radioactive and those above Z=92 are manmade in nuclear reactors or accelerators.

The general properties of the transition elements are

1. They are usually high melting point metals.
2. They have several oxidation states.
3. They usually form colored compounds.
4. They are often paramagnetic.

The transition elements include the important metals iron, copper and silver. Iron and titanium are the most abundant transition elements. Many catalysts for industrial reactions involve transition elements

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                                  2 to 9 are typical transition metal ion/compound coloured solutions

There are four transition series:

=>The first transition series: Scandium (Sc) through Copper (Cu):

=>The second transition series: Yttrium (Y) through Silver (Ag):

=>The third transition series: Lanthanum (La) to Hafnium (Hf) through Gold (Au))

=>The forth transition series which is incomplete: Actinium (Ac) to element 104 through element 109: 6d      

     subshell is filling, If elements 110 and 111 are found then this will complete this series.

Electronic Configuration of first transition series:

Group (column) number	3	4	5	6	7	8	9	10	11	12
atomic numberelement symbol	21Sc	22 Ti	23 V	24Cr	25Mn	26Fe	27Co	28Ni	29Cu	30Zn
electron configuration	3d14s2	3d24s2	3d34s2	3d54s1	3d54s2	3d64s2	3d74s2	3d84s2	3d104s1	3d104s2

The valence configuration for first series transition metals (Groups 3 - 12) is usually 3dⁿ 4s².

Exceptions:  The electron configurations for chromium (3d⁵ 4s¹) and copper (3d¹⁰ 4s¹) are exceptions.
This is because the 3d and 4s orbitals are very close in energy, and the energy of the 3d orbital drops going across the row.
For both chromium and copper the configuration which has more electrons in in the 3d subshell is of lower energy.  

At chromium  the difference in 3d and 4s orbital energies is of the order of the pairing energy (remember electron pairing requires energy).
The  3d⁵ 4s¹ configuration is of lower energy because this configuration has the maximum number of unpaired electrons for a d-subshell.

At copper (near the end of the transition series)  the 3d orbital energy has dropped so that 3d orbitals are actually lower in energy than 4s orbitals
This means the 3d¹⁰ 4s¹ configuration is of lower energy because it has more in 3d orbitals.

For the transition metal atoms, the total number of valence electrons equals the number of the column (group) in the periodic table (counting from the left).

For transition metal ions having charge ≥ +2, the number of d electrons equals the total number of valence electrons minus the charge on the ion.
This is because

Orbitals in the 3d and 4s subshells are of similar energy.
In transition metal atoms  the 4s subshell is of lower energy than the 3d subshell.
In transition metal ions of charge ≥ +2, 3d is of lower energy than 4s.
In transition metal ions of charge ≥ +2, all valence electrons in the d-subshell.
Therefore
Ni (Group 10) has 10 valence electrons and Ni²⁺ is d⁸
Fe (Group 8) has 8 valence electrons and Fe³⁺ is d⁵
Ti (Group 4) has 4 valence electrons and Ti³⁺ is d¹

*d and f – blocks elements show exceptional properties.

*d – block; the covalent radii of the elements decrease from left to half midway (to Cr) and the radii of the elements, from Cr to Cu are very close to one another until near the end where the size increase slightly, across a period; (Why?) because on passing from left to right, extra protons are placed in the nucleus and extra orbital electrons are added. The orbital electrons shield the nuclear charge incompletely (because shielding effect order s > p > d > f). Because of this poor shielding effect by d - electrons, the nuclear charge attracts all of the electrons more strongly: hence a contraction in size occurs.¹

            The radii of the elements, from Cr to Cu are very close to one another; this is due to the fact that the successive addition of d – electrons screens the outer electron (4s) from the inward pull of the nucleus. As a result of this, the size of the atom does not alter much in moving from Cr to Cu.²

            Near the end of the period, there is a slight increase in the atomic radii; this is due to the fact that near the end of the series, the electron - electron repulsions between added electrons in the same orbitals are greater than the attractive forces due to the increased nuclear charge. This result in the expansion of the electron cloud and thus the atomic radius increases.

Sc 1.44

Ti 1.32

V 1.22

Cr 1.17

Mn 1.17

Fe 1.17

Co 1.16

Ni 1.15

Cu 1.16

Zn 1.25

*The elements in the 3^rd group of the d-block show the expected increase in size Sc ----Y-----La.

*The elements in the 4^th to 12^th group of d – block the atomic radii of the first and second transition elements show expected increases but  radii of the second  and third transition series are almost same;  (why?) this is due to the fact in the atoms of the second transition series, the number of shells increases, their atomic radii are larger than those of the elements of the first transition series but in second and third transition metals between  lanthanum and hafnium there are 14 lanthanide elements are present, in which the antepenultimate 4f shell of electrons is filled. There is a gradual decrease in size of the 14- lanthanide elements from Ce to Lu. This is called lanthanide contraction. Thus the lanthanide contraction cancels almost exactly the normal size increases on descending a group of transition elements.

Ionic Radii:-

“The effective distance from the center of the nucleus of the ion upto which it exerts its influence on its electronic cloud is called ionic radii.”  Cation always has smaller radius and an anion is always larger than it parent atom.

*In case of isoelectronic ions, i.e., ions having same number of electrons but different nuclear charge, ionic radius decreases with increase in nuclear charge.

*Ionic radii follow the same trend as the atomic radii.

Oxidation states

The transition elements exhibit a large number of oxidation states. With the exception of few elements, most of these show variable oxidation states. These different oxidation states are related to the electronic configuration of their atoms. For example, the oxidation states exhibited by the transition elements of  the first series are listed in TABLE.

Different oxidation states of First transition series

Element		Outer electronic   configuration	Oxidation states
Sc	3d14s2		+2,+3
Ti	3d24s2		+2,+3,+4
V	3d34s2		+2,+3,+4,+5
Cr	3d54s1		(+1),+2,+3,(+4),(+5),+6
Mn	3d54s2		+2,+3,+4,(+5),+6,+7
Fe	3d64s2		+2,+3,(+4),(+5),(+6)
Co	3d74s2		+2,+3,(+4)
Ni	3d84s2		+2,+3,+4
Cu	3d104s1		+1,+2
Zn	3d104s2		+2

*Oxidation states with in the brackets are unstable.

Explanation

The existence of the transition elements in different oxidation states means that their atoms can lose different number of electrons. This is due to the participation of inner (n-1) d-electrons in addition to outer ns-electron because, the energies of the ns and (n-1) d-subshells are almost equal. For example, scandium has the electronic configuration of 3d¹4s². It exhibits an oxidation state of +2 when it uses both of its two 4s-electrons for bonding. It can also show oxidation state of +3 when it uses its two s-electrons and one d-electron. 

Important conclusion regarding oxidation states of transition elements

The examination of common oxidation states shown by different transition metals reveals the following facts :

(i)  The variable oxidation states of transition metals are due to participation of inner (n -1) d and outer  n s-electrons. The lowest oxidation state corresponds to the number of ns-electrons. For example, in the first transition series , the lowest oxidation states of Cr (3d⁵4s¹) and Cu(3d¹⁰4s¹) are +1 while for others , it is +2 (3d¹⁰4s²) 

(ii) Except scandium, the most common oxidation state of the first row  transition elements is +2 which arises due to loss of two 4s-electrons. This means that after scandium 3d-orbitals become more stable and therefore , are lower in energy than the 4s-orbitals. As a result, electrons first removed from 4s-orbitals.

(iii) For the first five elements, the minimum oxidation state is equal to the number of electrons in the s-orbitals and the other oxidation states are given by the sum of outer  s - and some or all d-electrons. The highest oxidation state is equal to the sum of the outer s (ns ) and (n -1) d-electrons. For the remaining five elements , the minimum oxidation state is given by the electrons in s-orbital while the maximum oxidation state is not related to their electronic configurations. The highest oxidation state shown by any transition metal is +8.

(iv) In the +2 and +3 oxidation states, the bonds formed are mostly ionic. In the compounds of higher oxidation states (generally formed wth oxygen and fluorine), the bonds are essentially covalent. Thus the bonds in +2 and +3 oxidation states are generally formed by the loss of two or three electrons respectively, while the bonds in higher oxidation states are formed by sharing of d-electrons. For example MnO₄^- (Mn in +7 ) state all the bonds are covalent.

(v)  Within a group, the maximum oxidation state increases with atomic number. For example, iron (Group 8) shows common oxidation states of +2 and +3 but ruthenium and osmium in the same group form compounds in the +4 , +6 and +8 oxidation states.

(vi) Transition metals also form compounds in low oxidation states such as +1 and 0 or negative. The common examples are [Ni(CO)₄], [Fe(CO)₅] in which nickel and iron are in zero oxidation state.

Complex formation tendency of transition metals:

Transition metal ions have very large vacant d- orbitals in their valence shells. As a result of this, lone pairs of electrons from ligands easily fill up the valence shells. When this happens, coordination compounds are formed. Note that this also takes place because d-orbitals have definite shapes and size as opposed to f-orbitals. That's why transition metals form coordinationcompounds while inner transition elements don't.

Magnetic properties of transition metals: 

The magnetic properties of a compound is a measure of the number of unpaired electrons in it. There are two main types of substances :

i) Paramagnetic substances : The substances which are attracted by magnetic field are called paramagnetic substances and this character arises due to the presence of unpaired electrons in the atomic orbitals. 

ii) Diamagnetic substances :  The substances which are repelled by magnetic field are called diamagnetic substances and this character arises due to the presence of paired electrons in the atomic orbitals.

            Most of the compounds of transition elements are paramagnetic in nature and are attracted by the magnetic field.

            The transition elements involve the partial filling of d-subshells. Most of the transition metal ions or their compounds have unpaired electrons in d-subshell (configuration from d¹ to d⁹) and therefore, they give rise to paramagnetic character. The magnetic character is expressed in Bohr magnetonsabbreviated as B.M. The magnetic moments of some ions of the first transition series are given in the following TABLE.

Magnetic moments of ions of first transition series.

Ion	Outer configuration	Number of unpaired electrons	Magnetic moment(mB)
Sc3+	3d0	0	0
Ti 3+	3d1	1	1.75
V3+	3d2	2	2.76
Cr3+	3d3	3	3.86
Cr2+	3d4	4	4.8
Mn2+	3d5	5	5.96
Fe2+	3d6	4	5.10
Co2+	3d7	3	4.4-5.2
Ni2+	3d8	2	2.9 - 3.4
Cu2+	3d9	1	1.8 - 2.2
Zn2+	3d10	0	0

The magnetic moments arise only from the spin of electrons. This can be calculated from the relation :

μ  = √n(n+2 B.M

Where n is the number of unpaired electrons and B.M represents Bohr magneton. It is clear from the Table that as the number of unpaired electron increases from 1 to 5, the magnetic moment and hence the paramagnetic character also increases. After d⁵ configuration, there is decrease in magnetic moment due to decrease in number of unpaired electrons. For example, d⁶ configuration has 4 unpaired electrons, d⁷ configuration has 3 unpaired electrons and so on. 

In addition to paramagnetic and diamagnetic substances, there are a few substances such as iron metal, iron oxide which are highly magnetic (about 1000 times more than ordinary metals). These are called Ferromagnetic substances.

Electronic Configuration of second transition series:

Second (4d) Transition Series (Y-Cd)

At. No.	39	40	41	42	43	44	45	46	47	48
Element	Y	Zr	Nb	Mo	Tc	Ru	Rh	Pd	Ag	Cd
Config.	4d15s2	4d25s2	4d45s1	4d55s1	4d55s2	4d75s1	4d85s1	4d105s0	4d105s1	4d105s2

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References:

1-http://hsc.csu.edu.au/chemistry/options/art/2769/ch984.htm#b1

2-http://www.ucc.ie/academic/chem/dolchem/html/comp/transmet.html

3-http://www.docbrown.info/page07/transition1.htm

4-http://www.newagepublishers.com/samplechapter/001210.pdf

5-http://www.bestchoice.net.nz/html/sa1/main/s148/p3960.htm