Introduction to Transition Elements

Definition of transition elements: (What is a transition metal?)

Transition metals are as metals which have partially filled d or f sub-shells in the neutral atom or in any of their usual positive oxidation states.

They can be classified into:

a) The main transition metals (those of the d-block)

b) The inner transition elements (those of the f-block).

Main transition elements

Main transition elements can be defined as those elements at least one of whose simple ions has an outer-shell of d-electrons.

Inner Transition Elements (f-block transition elements)

They are elements with partially filled f-sub-shells in common oxidation states. The f-block transition elements are in the two rows at the bottom of the periodic table.

i) Lanthanides (rare earths):

Comprise the fourteen elements having partially filled 4f-orbitals

ii) Actinides:

Comprise the fourteen elements having partially filled 5f-orbit

Electronic configuration of transition metals (d-block)

The general formula of electronic configuration of all transition elements are nS² (n-1)d^1-10

First transition or 3d series:

Element	Symbol	Atomic No.	Electronic Configuration
Scandium	Sc	21	[Ar] 3d¹4s²
Titanium	Ti	22	[Ar] 3d²4s²
Vanadium	V	23	[Ar] 3d³4s²
Chromium	Cr	24	[Ar] 3d⁵4s¹
Manganese	Mn	25	[Ar] 3d⁵4s²
Iron	Fe	26	[Ar] 3d⁶4s²
Cobalt	Co	27	[Ar] 3d⁷4s²
Nickel	Ni	28	[Ar] 3d⁸4s²
Copper	Cu	29	[Ar] 3d¹⁰4s¹
Zinc	Zn	30	[Ar] 3d¹⁰4s²

Second transition or 4d-series :

Element	Symbol	Atomic No.	Electronic Configuration
Yttrium	Y	39	[Kr] 4d¹5s²
Zirconium	Zr	40	[Kr] 4d²5s²
Niobium	Nb	41	[Kr] 4d⁴5s¹
Molybdenum	Mo	42	[Kr] 4d⁵5s¹
Technetium	Tc	43	[Kr] 4d⁵5s²
Ruthenium	Ru	44	[Kr] 4d⁷5s¹
Rhodium	Rh	45	[Kr] 4d⁸5s¹
Palladium	Pd	46	[Kr] 4d¹⁰5s⁰
Silver	Ag	47	[Kr] 4d¹⁰5s¹
Cadmium	Cd	48	[Kr] 4d¹⁰5s²

Third transition or 5d-series:

Element	Symbol	Atomic No.	Electronic Configuration
Lanthanum	La	57	[Xe] 5d¹6s²
Hafnium	Hf	72	[Xe] 4f¹⁴5d²6s²
Tantalum	Ta	73	[Xe] 4f¹⁴5d³6s²
Tungsten	W	74	[Xe] 4f¹⁴5d⁴6s²
Rhenium	Re	75	[Xe] 4f¹⁴5d⁵6s²
Osmium	Os	76	[Xe] 4f¹⁴5d⁶6s²
Iridium	Ir	77	[Xe] 4f¹⁴5d⁷6s²
Platinum	Pt	78	[Xe] 4f¹⁴5d¹⁰6s⁰
Gold	Au	79	[Xe] 4f¹⁴5d¹⁰6s^!
Mercury	Hg	80	[Xe] 4f¹⁴5d¹⁰6s

Anomalous Configuration:

(1) Chromium, Cr (Z=24) has the configuration [Ar] 3d⁵4s¹rather than the expected configuration [Ar] 3d⁴4s²

(2) Similarly, Cu (Z = 29), Ag (Z = 47), Pd (Z 46) prefer to have [Ar]3d¹⁰ 4s¹ , 4d¹⁰ 5s¹, 4d¹⁰ 5s⁰ Outer configurations respectively rather than 4d⁹5s² , 4d⁹5s², 4d⁸5s² respectively.

(3) In third transition series, Au has 5d¹⁰6s¹ outer configuration rather than

the expected configuration of 5d⁹ 6s².

Explanation for “out of sequence” or anomalous configuration:

The reason, for the abnormal configuration is that half-filled (eg. d⁵) and fully-filled (i.e.d¹⁰) levels are associated with extra stability.

The anomalous configuration of Copper:

This because the sub-levels 4s and 3d, are close in energy and as the atomic number (i.e. nuclear charge and attraction increases) the difference in energy between 4s and 3d decreases so that 3d becomes lower in energy than 4s. and take the configuration 3d¹⁰4s¹.

The anomalous configuration of Chromium:

For Cr, the 3d and 4s sub-levels are close enough (i.e. 3d and 4s are almost degenerate) that Hund’s rule prevails, and the six electrons apparently prefer to occupy singly the 4s and five 3d orbitals with spins unpaired rather than pairing in a 4s² 3d⁴ configuration.

Sample Problem. State a qualitative rule which can explain anomalous

configurations. Justify with examples.

Solution: The qualitative rule is that there is a particular stability associated with orbitals that are half-filled and filled with electrons. Examples are:

Cr (Z = 24) (Ar)¹⁸ 3d⁵ 4s¹(3d half filled)

Cu (z = 29) (Ar)¹⁸ 3d¹⁰ 4s¹ (3d completely filled)

“It may be that the more symmetrical charge distribution for the half-filled and fully-filled orbitals is the reason for these favored and preferred configurations.”

General properties of transition elements:

The general characteristics of transition metals can be summarized in the following points

1) the physical properties of transition metals shows that they characterize by hardness (hard metal), high density, high melting and boiling temperatures, conductivity of heat and electricity (conduct both heat and electricity), malleability (they are malleable) and ductility (they are ductile), metallic luster (lustrous) and they form alloys.

2) Many are electropositive enough to displace hydrogen from acid solutions (low ionization energies). This means that they are better reducing agents (give electrons more easily) than hydrogen in acid solutions (They have reducing properties).

3) With few exceptions, d-block metals exhibit multiple oxidation state (or variable valency) in their compounds.

4) With few exceptions, ions and compounds of the transition metals are colored in one or more of their oxidation states.

5) Regarding the magnetic properties, these elements form many paramagnetic compounds in which there are unpaired electrons, (with few exceptions).

6) Many of these elements and their compounds act as catalysts for chemical reactions they have strong catalytic properties.

7) One of the most important properties of transition elements is the ability to form coordination compounds or complex ions

8) Because of the similarity of their atomic radii, they are able to form alloys between themselves as well as with other metals in periodic table.

Detailed Study of the Properties of Transition Elements:

1. Electronic configuration of transition metals.

As mentioned before, Copper and Chromium have out of sequence configurations

2. Density of transition metals.

The densities of the transition metals are higher than those of the s-block elements of the same period. This is attributable to the increase in the nuclear masses and to the lower atomic volumes which result from the outer electrons being pulled in by the increased nuclear charge.

3. Metallic Character transition metals.

All the transition metals are metals; they are hard, malleable and ductile.

The presence of unfilled d-orbitals favors both covalent and metallic bonding.

These metals are good conductors of heat and electricity.

4. Melting points and boiling points transition elements.

The transition metals have much higher melting points and boiling points than other metals.

High melting points are characteristic of incomplete filling of electron shells (unfilled d-orbitals), with the result that high stability has to be achieved by the forming of a large number of metallic bonds

Note the following points:

a) Zinc, Cadmium, Mercury have very low melting points as they have fully-filled d-orbitals.

b) Calcium (filled 3d sub-shell) and Mn (half-filled 3d sub-shell, have lower

melting points than adjacent members of the first transition series.

The melting points of the transition metals increase from Sc to Cr and then decreases from Fe, Cu to Zn. The same pattern is observed for 2^nd (4d) and 3^rd transition series.

As we move across a row the melting points increase, reaching a maximum at the Group V or VI elements, after which the melting points decrease.

The reason of melting point trend:

The melting point and boiling point depend upon the strength of metal bonding, which depends on the number of unpaired electrons in the metal atoms

At beginning of a period of transition metals, there is one unpaired d-electron. The number of unpaired d-electron increases across a period until group VIB, after which the electrons begin to pair.

5. Ionization Energy of transition elements.

The first ionization energies are higher than s-block due to the increase in the nuclear charge, thus outer electrons of the d-block elements are more firmly held.

This suggests that all transition metals are typical metals of moderate to weak electropositivity. They are less, electropositive than s-block elements.

The I.E. increases from left to right across a row of transition elements.

6. Electronegativity values.

This electronegativity increases from left to right in the periodic table.

Because they are more electronegative than s-block elements, it is no surprise that covalency features more in the properties of the transition elements. i.e. the tendency to form covalent compounds increases as the atom gets smaller i.e. from Sc to Cu.

7. Standard oxidation potential and reducing properties.

Most of the transition metals have higher oxidation potential than hydrogen (taken as zero) and can displace H₂ from dilute solution of acids such as HCl, H₂SO₄. On the other hand, metals like Cu, Ag, Au, Hg have lower oxidation potential than H₂ and cannot displace H₂ from dilute HCI, H₂SO₄

Despite their high oxidation potentials (except Cu, Ag, Au, Hg), the transition metals are poor reducing agent. This is attributed to their high heats of vaporization, high ionization energies which more than offset the high heats of hydration of their small transition elements.

8. Variable oxidation state.

The transition metals show multiple oxidation states in their compounds.

The cause of showing different oxidation states is:

The 4s and 3d (in general ns and (n -l)d orbitals are very close in energy. Hence electrons from ns (e.g. 4s) and a variable number of electrons from inner (n -1)d electrons (e.g. 3d) can be used for compound formation.

Note the following:

(1) The maximum oxidation state is equal to the total number of 3d and 4s electrons e.g. Mn (3d⁵4s²) shows highest oxidation state of +7.

(2)

a. The higher oxidation states are acidic; they tend to form acidic oxides and largely covalent compounds e.g.

b. The lower oxidation states are basic; the lower states tend to, form basic oxides and compounds which possess a higher degree of ionic character. The basic character decreases from left to right across the series.

c. The intermediate oxidation states are amphoteric, e.g. MnO₂, Fe₂O₃,Fe₃O₄.

(3)

a. The stability of the highest oxidation states for the first transition series elements decreases from left to right across the series i.e. from Sc to Zn.

b. The elements of the 2nd and 3rd transition series generally are stable with higher oxidation states than are the elements of the first series.

c. Compounds of titanium (IV) are harder to reduce than those of Vanadium (V), which in turn are harder to reduce than those of Cr (VI) or Mn (VII).

Stability of oxidation states:

The factors which influence the stability of various oxidation states are:

A) Ionization energies: The 2^nd I. E. of copper is unusually high; this fact explain the existence of the unusual +1 oxidation state of copper.

B) Electronic structure: Mn (II), and Fe (Ill) are more stable than expected due to half- filled shell.

C) Nature of the ligands: Certain types of ligand stabilize certain oxidation states.

D) Ambient Conditions: Solvent, temperature, pH etc.

Question about transition metals part 2

MCQ Questions with answers in inorganic chemistry

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