Periodic Table - Transition Metals - Introduction - Doc Brown's Chemistry  Revising Advanced Level Inorganic Chemistry Periodic Table Revision Notes

Part 10. Transition Metals and the 3d–block - main index, introduction and data table

The 3d block of metals consists of the ten elements ₂₁Sc to ₃₀Zn (to fill the 3d sub-shell), but the true transition elements run from ₂₂Ti to ₂₉Cu, because they can form an ion with an incomplete d sub–shell and are know as the 1st transition metal series of elements.

Sub-index for this page

10.1 (a) A very brief introduction to transition metals

10.1 (b) Electronic structure and the position of the 3d block and transition metals in the periodic table

10.1 (c) Comparison of certain properties of the 3d block of metals and other elements on period 4

10.2 (d) Data Table: Summary of selected properties – concentrating only on the 3d–block series

10.2 (e) Five general physical/chemical characteristics related to electron configurations of 3d block metals

INDEX of ALL my 3d block and transition metal pages (23 pages in all)

INORGANIC Part 10 3d block TRANSITION METALS sub–index:

10.1–10.2 Introduction 3d–block Transition Metals (this page)

The individual chemistry pages of the 3d block elements including transition metals

10.3 Scandium * 10.4 Titanium * 10.5 Vanadium * 10.6 Chromium * 10.7 Manganese * 10.8 Iron

10.9  Cobalt * 10.10 Nickel * 10.11 Copper * 10.12 Zinc * 10.13 Other Transition Metals e.g. Ag and Pt

Various appendix pages on transition metals

Appendix 1. Hydrated salts, acidity of hexa–aqua ions and patterns explained

Appendix 2. Introduction to complexes and ligands

Appendix 3. Complexes and isomerism

Appendix 4. Electron configuration & colour theory

Appendix 5. Redox equations, feasibility, E^ø for a reaction

Appendix 6. Catalysis - heterogeneous and homogeneous examples explained

Appendix 7. Redox equations - reactions involving change in oxidation state of a transition metal

Appendix 8. Stability constants of complex ions and entropy changes

Appendix 9. Colorimetric analysis to determine a complex ion formula

Appendix 10 3d block – extended data tables

Appendix 11 Some 3d–block compounds, complexes, oxidation states and an electrode potential chart

Appendix 12 Hydroxide complex precipitate 'pictures', formulae and equations

See also absorption spectra of transition metals

10.1. Introduction to the 3–d block and 1st transition metal series

10.1 (a) A very brief introduction to the 3d block of elements

The 3d block includes the first series of transition elements and make up part of period 4 of the periodic table

A transition element is a metal whose atoms or at least one ion has a partially filled d subshell.

Scandium and zinc are not true transition metals but are still 3d block elements.

Typical properties of transition elements:

Coloured compounds and ions.

Form complex ions

The transition metal or its compounds show catalytic properties.

Variable oxidation states within the elements chemistry and can range from +1 (Cu) up to +7 (Mn).

High melting point.

High density.

High tensile strength.

Show paramagnetism.

Sub-index for this page

10.1 (b) Electronic structure and the position of the 3d block and transition metals in the periodic table

The d block series of metals lie between the s block of groups 1-2 and p block of group 3/13 to group 0/8/18.

Pd

s block

d blocks (3d block manganese) and f blocks of metallic elements

p block elements

Gp1

Gp2

Gp3/13

Gp4/14

Gp5/15

Gp6/16

Gp7/17

Gp0/18

1

1H

2He

2

3Li

4Be

The modern Periodic Table of Elements ZSymbol, z = atomic or proton number 3d block of metallic elements: Scandium to Zinc focus on manganese

5B

6C

7N

8O

9F

10Ne

3

11Na

12Mg

13Al

14Si

15P

16S

17Cl

18Ar

4

19K

20Ca

21Sc [Ar]3d14s2 scandium

22Ti [Ar]3d24s2 titanium

23V  [Ar] 3d34s2 vanadium

24Cr [Ar] 3d54s1 chromium

25Mn    [Ar]   3d54s2 manganese

26Fe [Ar] 3d64s2 iron

27Co [Ar] 3d74s2 cobalt

28Ni [Ar] 3d84s2 nickel

29Cu [Ar] 3d104s1 copper

30Zn [Ar] 3d104s2 zinc

31Ga

32Ge

33As

34Se

35Br

36Kr

5

37Rb

38Sr

39Y

40Zr

41Nb

42Mo

43Tc

44Ru

45Rh

46Pd

47Ag

48Cd

49In

50Sn

51Sb

52Te

53I

54Xe

6

55Cs

56Ba

57,58-71

72Hf

73Ta

74W

75Re

76Os

77Ir

78Pt

79Au

80Hg

81Tl

82Pb

83Bi

84Po

85At

86Rn

7

87Fr

88Ra

89,90-103

104Rf

105Db

106Sg

107Bh

108Hs

109Mt

110Ds

111Rg

112Cn

113Nh

114Fl

115Mc

116Lv

117Ts

118Os

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Element	Electron configuration	Electron spin box diagrams of the outer electron orbitals.	Comments Gp = Group!
19 Potassium	1s22s22p63s23p64s1	[Ar]3d4s4p	K, s–block, Gp1 Alkali Metal, v. reactive
20 Calcium	1s22s22p63s23p64s2	[Ar]3d4s4p	Ca, s–block, Gp2 Alkaline Earth Metal
21 Scandium	1s22s22p63s23p63d14s2	[Ar]3d4s4p	Sc, 3d block, not a true Transition Metal
22 Titanium	1s22s22p63s23p63d24s2	[Ar]3d4s4p	Ti, 3d block, a true Transition Metal
23 Vanadium	1s22s22p63s23p63d34s2	[Ar]3d4s4p	V, 3d block, a true Transition Metal
24 Chromium	1s22s22p63s23p63d54s1	[Ar]3d4s4p	Cr, 3d block, a true Transition Metal
25 Manganese	1s22s22p63s23p63d54s2	[Ar]3d4s4p	Mn, 3d block, a true Transition Metal
26 Iron	1s22s22p63s23p63d64s2	[Ar]3d4s4p	Fe, 3d block, a true Transition Metal
27 Cobalt	1s22s22p63s23p63d74s2	[Ar]3d4s4p	Co, 3d block, a true Transition Metal
28 Nickel	1s22s22p63s23p63d84s2	[Ar]3d4s4p	Ni, 3d block, a true Transition Metal
29 Copper	1s22s22p63s23p63d104s1	[Ar]3d4s4p	Cu, 3d block, a true Transition Metal
30 Zinc	1s22s22p63s23p63d104s2	[Ar]3d4s4p	Zn, 3d block, not a true Transition Metal
31 Gallium	[Ar]3d104s24p1	[Ar]3d4s4p	Ga, p–block, Gp3/13, Boron group
32 Germanium	[Ar]3d104s24p2	[Ar]3d4s4p	Ge, p–block, Gp4/14, Carbon group
33 Arsenic	[Ar]3d104s24p3	[Ar]3d4s4p	As, p–block, Gp5/15, Pnictogen
34 Selenium	[Ar]3d104s24p4	[Ar]3d4s4p	Se, p–block, Gp6/16, Chalcogen
35 Bromine	[Ar]3d104s24p5	[Ar]3d4s4p	Br, p–block, Gp7/17 Halogen
36 Krypton, Kr	[Ar]3d104s24p6	[Ar]3d4s4p	Kr, p–block, Gp 0/8/18 Noble Gas

• The change in electron configuration across period 4 of the periodic table.

  • A d sub-shell can hold a maximum of 10 electrons.

  • The 3d block corresponds to filling of the 3d sub-shell orbitals from scandium to zinc (Z = 21 to 30, 10 elements).

• It is the particular electronic structure of the transition metals within each d block that facilitates chemical properties significantly different from elements in other parts of the periodic table.

• Also, within the d blocks themselves, there are many similarities between the elements regarded as true transition metals e.g. titanium to copper.

  • For a more details see section 10.2 (e) General chemical characteristics and electron configurations of transition metals

  • The differences between elements within the 3d block are less obvious than e.g. the period 4 elements in the p block (from gallium to krypton).

  • Each successive member of the 3d block gains one more electron that goes into a 3d orbital, these are not outermost shell electrons because the 4s sub-shell is filled first - electron configurations quoted above in the context of their position in the periodic table.

• The elements scandium to zinc (Z = 21 to 30) are known as the 3d block of elements or 3d–block of metals because here the first of the possible d sub–shells is progressively filled (3d–block – first row of the d–blocks of ten elements).

  • See rules on how to work out electron configurations

• The true transition elements run from ₂₂Ti to ₂₉Cu, because they can form an ion with an partially filled 3d sub–shell and are know as the 1st transition metal series of elements.

• The transition elements are group of industrially important metals mainly due to their strong inter–atomic metallic bonding giving them generally high melting/boiling points and high tensile strength.

• These–called 'transition metal characteristics' arise from behaviour of the d sub–shell energy level electrons but scandium and zinc are not true transition metals i.e. Ti to Cu are the real transition elements (reasoning later).

• Note that physically, zinc is low melting and a lower tensile strength compared to the others in the 3d block.

• Although scandium is physically typical of a transition metal e.g. high melting point and high tensile strength,

  • chemically, scandium only forms a single and colourless triple charged ion ([Ar]3d⁰ for the Sc³⁺ ion),

  • with no incomplete 3d shell in the ion, so scandium shows little of the general characteristics associated with transition metals.

• Therefore, a similar argument applies to the single and colourless doubly charged ion ([Ar]3d¹⁰ for the Zn²⁺ ion),

  • so zinc does not show the typical characteristics of transition metal chemistry,

  • again with no incomplete 3d shell in the ion, zinc shows little of the general characteristics associated with transition metals.

  • e.g. scandium or zinc do not exhibit variable oxidation state, coloured complex ions, catalytic properties of the metal or ion.

  • See a definition of a transition metal below.

• Therefore probably the best definition of a transition metal is an element which forms at least one ion with partially filled d sub–shell containing at least one electron.

  • How this relates to variable oxidation state and coloured complex ions is elaborated further in section 10.2 and the subsequent sections on the individual metals (links below) and some of the. Zinc (Zn²⁺, [Ar]3d¹⁰) and scandium (Sc³⁺, [Ar]3d⁰) cannot meet this criteria.

• The presence of the partially–filled d sub–shells of electrons gives transition elements properties which are not in general possessed by the main group elements, namely Groups 1–2 and 3/13 to 0/8/18, BUT, there are similarities with other metals.

  • Watch out for the modern group notation of 1 to 18, a move on from 1 to 0/8.

PLEASE NOTE the following about my 3d-block elements - transition metals notes:

• All the reactions are described with visual observations and full ionic equations whether redox reactions or not.

• I have made extended use of standard electrode potentials to indicate not only the relative oxidising/reducing power of a half–cell reaction, but also to argue for the thermodynamic feasibility of a reaction.

  • However, if you have done little on using electrode potentials you should still be able to follow the reaction descriptions and the accompanying equations.

  • Electrode potentials (half–cell potentials) and how to work out E^Ø_reaction (which must be a +ve voltage for the reaction to be feasible) are all explained in Equilibria Part 7 Redox Equilibria and Electrode Potentials.

  • How to use half–cell reactions to work out full redox equations is explained on Redox Reactions Part II section 6 Constructing redox equations.

• In the latest Periodic Table convention, the 3d–block metallic elements are considered the 'head elements' of Groups 3–12.

  • Groups 1–2 remain unchanged but Groups 3–7 and 0 become Groups 13–18. I tend to retain the Groups 3–7 and 0/8 convention for the moment but future is 13–18 and I often indicate both!

  • The table shows the 'electronic' vertical connections of groups 3 to 12 of the even more modern periodic table.

There are actually many 'vertical' chemical similarities in a 'classic' periodic table way of thinking to justify this latest 'numbering' of the Periodic Table. e.g.

• In most cases the three elements quoted above, per vertical column, have the same outer electron configuration.

• 'Modern Group 3': Scandium and yttrium are very similar with a relatively simple M³⁺ ion chemistry.

• 'Modern Group 10': Nickel, palladium and platinum are good hydrogenation catalysts.

  • They all tend to form more square planar complexes than other transition elements.

• 'Group 11': Copper, silver and gold are relatively unreactive metals in terms of corrosion.

  • They form linear complexes like the cationic, [Ag(NH₃)₂]²⁺ or the anionic [CuCl₂]^– and [Au(CN)₂]^–.

  • All three are extremely good conductors of heat and electricity.

• 'Modern Group 12': Zinc and cadmium chemistry is mainly about the M²⁺ ion.

• From modern 'Group 3 to 7' the maximum known oxidation state (albeit in some pretty unstable compounds at times) is equal to the 'new' group number i.e. Sc/Y/La (+3) to Mn/Tc/Re (+7).

• The discontinuity of atomic/proton number from lanthanum to hafnium on period 6 is due to the insertion of the 4f–block elements ₅₈Ce to ₇₁Lu.

• When 3d block elements form ions, the 4s electrons are 'lost' first e.g.

  • Iron's electron configuration is 1s²2s²2p⁶3s²3p⁶3d⁶4s²  or [Ar]3d⁶4s², and

  • the iron(II) ion is 1s²2s²2p⁶3s²3p⁶3d⁶  or  [Ar]3d⁶

  • and the iron(III) ion is  1s²2s²2p⁶3s²3p⁶3d⁵  or  [Ar]3d⁵

Sub-index for this page

10.1 (c) Comparison of selected properties of the 3d block of metals and other elements on period 4

That is for Z = 1 to 38 particularly the preceding Group 1 metal potassium and the Group 2 metal calcium.

The electronic structures have been already discussed in 10.1 (b)

10.1 (b) Electronic structure and the position of the 3d block and transition metals in the periodic table

• Graphs of element properties, focus on Period 4

  • Periodicity plots for elements Z = 1 to 38 Look for Z = 21 (Sc) to 30 (Zn)

  • Survey of selected properties of Period 4 of the periodic table

• Melting/boiling points:

  • Generally higher than other elements in period 4.

• 1st ionisation energy:

  • The 3d block 1st ionisation energies tend to increase from left to right and fit in with the general pattern for period 4.

• Pauling electronegativity:

  • The 3d–block values range from a relatively low 1.3 to 1.9 and fit in with the general pattern of increasing value across period 4.

•   Atomic radius:

  • 3d–block elements have similar values and significantly less than for potassium and calcium, though there is a steady decrease from Sc to Mn with increase in atomic number.

• Electrical conductivity: The 3d–block are quite good conductors of electricity and very good in the case of copper (ditto silver Ag below Cu), otherwise generally less than potassium and calcium, but obviously much greater electrical conductors than the semi-metals and non-metals to the right of zinc on period 4.

• Density: 3d–block range from 3.0 to 8.9 g/cm³ and significantly more than for potassium (0.86) and calcium (1.5) and the semi-metals and non-metals of the p block (Ga to Kr).

• Periodicity plots for elements Z = 1 to 96 if you want to look for the 4d and 5d blocks!

Other comparison points of the elements titanium to copper (true transition metals) with nearby metals.

Potassium (+1), calcium (+2) and scandium (+3) only have one oxidation state in compounds, whereas Ti to Cu have compounds in at least at least three oxidation states, even if some are not very stable!

Sub-index for this page

10.2. General information and data table for the 3d block metals Sc–Zn

10.2 (d) Data Table 1 – summary of selected properties – concentrating only on the 3d–block series

Elect. pot. = standard electrode potential data for 3d block transition elements (all metals) (E^Ø at 298K/25^oC, 101kPa/1 atm, 1M solutions.)

na = data not applicable to that particular 3d block transition metal

CLICK for a more detailed data table 2 summary

• The transition metals are the most important structural metals for industry due to their strength arising from the strong inter–atomic forces (see metal bonding and alloy structure).

• The strong bonding is due to small ionic radii and at least 3 delocalised 3d or 4s electrons contributing to the bonding which accounts for their high tensile strength, malleability (can be readily beaten into shape) and ductility (can be drawn into wire).

• They are silvery–grey solids apart from the dark orange of copper.

• See the

  • discussion on electron configurations in section (b)

  • graphs of the physical properties in section (c)

• They generally have high melting/boiling points and densities and readily mix with themselves or other elements to give a huge variety of alloys with a wide range of uses based on varied hardness, strength, malleability and anti–corrosion properties. 

• There is a general, but small, contraction of the atomic/ionic radii across the series as the atomic/proton number rises, i.e. an increasing positive attractive force on the outer electrons of the same sub–shells (3d and 4s).

• The 3d block metals are more dense than the metallic elements to the left or the semi-metals and non-metals to the right.

• Although the 3d block elements conduct electricity, as expected for most metals, there is considerable variation e.g. scandium is relatively poor, whereas copper is an extremely good electrical conductor - it is also the same for thermal conduction.

Sub-index for this page

10.2b. (e) General chemical characteristics of the 3d block, including transition metals

Wherever possible, their properties are related to their electronic structure.

The 1st ionisation energies and electronegativities are neither particular high or particularly low for the 3d block metals and are of no great consequence at pre-university level chemistry e.g. the compounds of 3d block metals can involve both ionic or covalent bonding.

However, the subsequent ionisation energies (2nd, 3rd etc.) pattern for each element are important in helping to understand the oxidation states that are possible compared to the elements on the left and right - note particularly the restriction of only +1 and +2 oxidation states for group 1 and group 2 metals respectively.

21 Scandium, Sc	1s22s22p63s23p63d14s2	[Ar]3d4s
22 Titanium, Ti	1s22s22p63s23p63d24s2	[Ar]3d4s
23 Vanadium, V	1s22s22p63s23p63d34s2	[Ar]3d4s
24 Chromium, Cr	1s22s22p63s23p63d54s1	[Ar]3d4s
25 Manganese, Mn	1s22s22p63s23p63d54s2	[Ar]3d4s
26 Iron, Fe	1s22s22p63s23p63d64s2	[Ar]3d4s
27 Cobalt, Co	1s22s22p63s23p63d74s2	[Ar]3d4s
28 Nickel, Ni	1s22s22p63s23p63d84s2	[Ar]3d4s
29 Copper, Cu	1s22s22p63s23p63d104s1	[Ar]3d4s
30 Zinc, Zn	1s22s22p63s23p63d104s2	[Ar]3d4s

The chemistry of the 3d block metals is dominated by the behaviour of the 3d electrons.

The 3d block corresponds to the filling of the 3d sub–shell of electrons, best appreciated by the 'box diagrams' of their electron structure.

Each half–arrow is an electron, which tend to singly occupy the sub–orbitals as much as possible to minimise repulsion (Hund's Rule of maximum multiplicity). 

The electron arrangements are those that gives the lowest total energy.

To minimise repulsion between pairs of electrons in an orbital, they occupy the 3d sub-shell orbitals singly if possible (Sc to Ni).

The outer electrons of the neutral atoms are either in the 3d or 4s sub–shell, but only in 3d sub-shell in ions.

The 4s sub–shell is initially filled by potassium [Ar]4s¹ and calcium [Ar]4s² before start to fill the 3d sub-shell.

The electron arrangement for each element from Sc to Zn is also given at the start of each individual metal section in terms of s, p and d notation.

All 10 elements, Sc to Zn are 3d block elements (the filling of the 3d sub–shell) BUT a true transition metal element is one in which there is a partially filled d sub–shell, i.e. a d shell holding at least one electron in one or more chemically stable ions (Ti to Cu).

For 3d block metals this means at least one stable ion with the configuration within the range [Ar]3d¹ e.g. Ti³⁺ to [Ar]3d⁹ e.g. Cu²⁺ and so excludes scandium and zinc which do not have an incomplete sub-shell in any stable ion.

Zinc only forms Zn²⁺, [Ar]3d¹⁰ and scandium only forms Sc³⁺, [Ar]3d⁰, so neither can meet this criteria for a true transition metal. See theory of colour in transition metal complexes.

There are two apparent anomalies in the electron configuration sequence from left to right as the 3d sub–shell energy level is filled.

The energies of the 3d and 4s orbitals are quite similar, so in two cases the alternative configuration with singly filled 4s orbital is just slightly lower in energy, though is of no significance in terms of their 'transition metal chemistry'.

Cr is [Ar]3d⁵4s¹ and not [Ar]3d⁴4s²

and Cu is [Ar]3d¹⁰4s¹ and not [Ar]3d⁹4s²

The electron arrangements are those that gives the lowest total energy.

because an inner half–filled or fully–filled 3d sub–shell seem to be a little lower in energy, and marginally more stable.

The total number of outer 3d/4s electrons is equal to the maximum oxidation state from Sc (+3) to Mn (+7) and there are many stable compounds exhibiting these maximum oxidation states.

After Mn there is significantly less stability of species with the metal in oxidation states above +3 for Fe and Co, and above +2 for Ni, Cu and Zn (only +2 is exhibited).

The four 'classic' chemical characteristics and a note on their general physical properties

(but NOT unique to transition metals)

(1) Complex formation:

A complex ion is  formed when a transition metal ion is surrounded by ligands which dative covalent bond with it by donating lone pairs of electrons into vacant d orbitals. The d electrons of the metal ion do not take part in ligand bonding.

Appendix 2  offers a more detailed introduction to complex ions of transition metals as well as numerous examples 'en route' particularly from Ti to Cu.

A summary of some important definitions - all explained in more detail with examples in Appendix 2.

(i) A ligand is a neutral molecule or ion that forms a co-ordinate (dative covalent) bond with a central transition metal atom or ion by donation of a pair of electrons.

The ligands can be monodentate, bidentate or polydentate, depending on how many lone pairs of bonding electrons can be donated per ligand molecule.

(ii) A complex is a central metal atom or ion (often a transition metal) surrounded by, and bonded to, a number of ligands.

The complex can be charged or neutral e.g. [Cu(H₂O)₆]²⁺  or   [PtCl₂(NH₃)₂Cl₂]⁰

(iii) The co-ordination number is the number of co-ordinate (dative) covalent bonds to the central metal atom or ion of the specific complex.

Do NOT assume this equals the actual number of ligand molecules.

(2) Formation of coloured ions:

 Appendix 4  offers an introduction to the origin of the colour in transition metal complex ions as well as examples 'en route' from colourless 'non–transition' Sc³⁺ complexes, to coloured Ti^{II, III, IV} to Cu^II 'true transition' complexes and finally explaining colourless 'non–transition' Zn²⁺ complexes at the end of the 3d–block.

The simplified diagram ignores the true complex structure of aqueous ions.

The above diagram illustrates some of the colours of hydrated 3d-block ions, often [M(H₂O)₆]ⁿ⁺, where n = 2 or 3, and the electron configuration of the central metal ion (Mⁿ⁺) of the complex in terms of written convention and 'electron boxes' for the 3d electrons.

Only the true transitions can exhibit a coloured ion with a partially filled 3d set of orbitals.

Note that the 4s electrons are remove first.

Remember [Ar] is shorthand for 1s²2s²2p⁶3s²3p⁶

(3) Variable and maximum oxidation states – variable valency:

• From Sc to Mn the maximum oxidation state is determined by the total maximum number of 3d and 4s electrons. After that, things get very complicated but the maximum tends to fall down to +2 for zinc after +3 for Fe and Co (there are some higher oxidation state species, but not that common and not that stable in aqueous media).

  • See Table of examples of compounds/complexes for different oxidation states of the 3d–block elements.

• The relative ease of oxidation state change for Ti to Cu AND the maximum oxidation state formed by Sc to Mn, is partly explained by considering the ionisation energies involved and a comparison with Group 1, 2 and 3/13 metals too.

  • In the sequences below the atoms and ionised species are all in the gaseous state as is the convention for ionization energy data.

  • The energies (kJmol^–1) required to remove the next most loosely bond electron to give the next more highly charged ion (the next higher oxidation state) are shown as a sequence.

  • Only for the first example, potassium, are the full formal equations shown.

  • The successive enthalpy of ionisation sequences for Group 1 (potassium), Group 2 (calcium), the 3d–block (e.g. titanium) and Group 3 (gallium) are now considered for period 4 (in kJmol^-1).

  • Gp1: K_(g) == +418 ==> K⁺_(g) == +3070 ==> K²⁺_(g) 

    • they would be formally written as:

      • for the 1st ionisation enthalpy: K_(g) – e^– ===> K⁺_(g)  

      • and for the 2nd ionisation enthalpy: K⁺_(g) – e^– ===> K²⁺_(g)  

      • What you notice is the huge increase in the 2nd ionisation enthalpy of potassium and hence very prohibitive for potassium to have a valency of 2 i.e. to form ionic compounds containing a K²⁺ ion.

  • Gp2: Ca_(g) == +590 ==> Ca⁺_(g) == +1150 ==> Ca²⁺_(g) == +4940 ==> Ca³⁺_(g) 

    • I'll set out other equations just using M to represent the metal and refer back to these equations e.g.

      • 3rd ionisation enthalpy:  M²⁺_(g)  -  e^-  ===> M³⁺_(g)

      • 4th ionisation enthalpy:  M³⁺_(g)  -  e^-  ===> M⁴⁺_(g)

      • 5th ionisation enthalpy:  M⁴⁺_(g)  -  e^-  ===> M⁵⁺_(g)

      • So, when M is Ca, there is a huge increase in the 3rd ionisation enthalpy, so Ca³⁺ containing compounds are energetically very unfavourable.

  • 3d–block: e.g. Ti_(g) == +661 ==> Ti⁺ == +1310 ==> Ti²⁺ == +2720 ==> Ti³⁺ == +4170 ==> Ti⁴⁺ == +9620 ==> Ti⁵⁺

    • For titanium the big increase occurs at the 5th ionisation energy, but compounds with oxidation state up to +4 are formed and stable.

    • A similar argument applies to manganese with compound up to +7 oxidation state, after that, things are not so simple.

    • So for transition metals it is energetically reasonable favourable to exhibit a variety of oxidation states with the more gradual increase in ionisation enthalpies.

  • Gp3/13: Ga_(g) == +577 ==> Ga⁺ == +1980 ==> Ga²⁺ == +2960 ==> Ga³⁺ == +6190 ==> Ga⁴⁺ 

    • With the group 3/13 element, there is a huge increase at the 4th ionisation energy, so it is energetically very unfavourable for gallium to exhibit a valency of 4 or oxidation state of 4 in its compounds.

    • The argument applies to the rest of the p block elements to you observe a maximum valency or oxidation state of + to +7 for group 4/14 to group 7/17 elements.

• So, for Groups 1, 2 and 3, the ionisation energy dramatically rises after the outer shell of s or p electrons are removed, i.e. a very stable electronic noble gas structure [Ar] for Groups 1 and 2 and [Ar]3d¹⁰ for the p block elements.

  • This gives a maximum positive stable oxidation state equal to the group number (old numbers 1 to 8).

  • The energy required (very endothermic) to make Na²⁺, Ca³⁺ and Ga⁴⁺ is too high to be compensated by exothermic bond formation with other elements like oxygen or chlorine etc.

  • Also note that intermediate lower oxidation states Ca⁺ and Ga²⁺ are not very stable either - but outside the scope of this page.

  • I'm afraid ionisation energies and electron arrangements are not the only factors to be considered, you also need to study the Born Haber Cycle in some detail to prove this, but not here and not usually on a pre–university course!

• For the transition metals, at first, the successive ionisation energies rise relatively gradually, due to the 3d/4s electron levels being of similar energy.

  • When all the outer s and d electrons are removed to leave an [Ar] core, there is a dramatic rise as an electron must be removed from the inner very stable noble gas (argon) core in the case of the 3d block of metals.

  • Therefore Ti has a maximum oxidation state of +4, but +2 and +3 species are also formed, but NOT +5 compounds.

  • This does mean however, across the 3d–block, there is the potential for very high oxidation states if there are enough 3s and 3d electrons that can be energetically favourably removed or become involved in stable bonding e.g. Mn has a maximum oxidation state of +7 by 'removing  *  or 'sharing' the outer 3d⁵4s² electrons. (see extended data table).

    • Similarly you can argue that the maximum oxidation states for vanadium would be +5 and chromium +6, as is indeed is the case!

    • After manganese, things get complicated and there is a general decrease from Mn (+7) to Zn (+2) in the maximum possible higher oxidation states, and many higher oxidation state compounds of Fe, Co, Ni and Cu are unstable and uncommon.

  •  *  Of course e.g. in manganese (VII) compounds, 7 electrons are not removed to give an Mn⁷⁺ ion, but, unlike calcium and gallium, true transition metals form many stable compounds of the 'intermediate' oxidation states e.g. manganese forms +2, +3, +4, +6 as well as +7 oxidation state compounds.

    • From scandium to manganese, the maximum positive oxidation state is numerically equal to the total number of 3d and 4s electrons (see table below), but after that, things are not so simple and you tend to get a gradual decline in maximum oxidation state from iron to zinc.

  • This is due to closeness of the energies of the 3d sub–shell electrons and the stabilising influence of ligand molecules like water or ammonia and ligand ions like chloride or cyanide.

    • Vacant 3d orbitals (and 4s/4p orbitals too) can accept pairs of electrons to for stable dative covalent bonds.

    • I'm afraid arguments for the characteristic variable oxidation states of transition metals based on ionisation energies and the similar energies of the 3d orbitals is a bit limited, but better than nothing!

    • I have done detailed notes on oxidation state/oxidation number and redox reactions – a lot of which come up in transition metal chemistry.

    • A summary of some of the possible 'major' and 'minor' oxidation states is given in the table below.

A clear trend in maximum oxidation state from scandium to manganese equalling the number of outer electrons from 3d¹4^s2 to 3d⁵4s².

From iron to zinc there is a general trend in decreasing maximum oxidation state, but not as clear a pattern as Sc to Mn.

The lower oxidation states are more often found in simple ionic compounds e.g.

Cr³⁺ in Cr₂O₃,  Mn²⁺ in MnCl₂, Cu²⁺ in CuSO₄ (though the salts might be hydrated)

The 'simple' cations can be considered to be at the centre of (often) octahedral complex ions e.g.

hydrated hexaaqua ions [M(H₂O)₆]ⁿ⁺  where M = the transition metal and n is usually 2 or 3.

Examples of the electronic structure of the central metal ion are shown below,

The electron configuration of the central metal ion (Mⁿ⁺) of the complex is shown in terms of written convention and 'electron boxes' for the 3d electrons.

Only the true transitions can exhibit a coloured ion with a partially filled 3d set of orbitals.

Note that the 4s electrons are remove first and remember [Ar] is shorthand for 1s²2s²2p⁶3s²3p⁶

For a more detailed discussions see:

Appendix 11 Some 3d–block compounds, complexes, oxidation states & electrode potentials

Appendix 5. Redox equations, feasibility, E^ø

Appendix 7. How to balance redox equations

 

(4) Catalytic activity by the elements and their compounds:

Many examples are quoted in the detailed notes and the theory of heterogeneous and homogeneous catalysis is outlined in Appendix 6.

(i) The surfaces of transition metals like iron, nickel and copper are good for the temporary absorption of substrate molecules, enabling strong covalent bonds to be broken, hence facilitating the reaction e.g. nickel catalyses the hydrogenation of alkenes (heterogeneous catalysis).

(ii) The catalytic properties of transition metal compounds usually involve temporary changes in oxidation state of the metal ion and  can be heterogeneous catalysis with a solid or homogeneous if a soluble catalyst complex ion.

In all cases, a reaction pathway of lower activation energy is facilitated.

(5) Physical characteristics of the transition metals

Transition elements tend to be dense, have high melting/boiling points, durable and hard wearing with a high tensile strength - the latter explaining their wide use as a very useful strong structural materials - further improvements from the use of alloys (see links below).

This is due to the strong metallic bonding between the atoms in the crystal structure of metals - in fact the lattice consists of metal immobile ions held together by delocalised electrons.

Transition metals can release a pool of delocalised electrons from both the inner shell (e.g. 3d) and outer shell (e.g. 4s) to contribute to the metallic bond.

This contrasts with the s-block elements of groups 1 and 2 where there metallic bond can only rely on the outer s electrons and hence generally have lower densities, melting/boiling points and tensile strength.

Alloying a metal

1. Pure metal and the regular arrangement of atoms in a metal lattice.

2. When metals are pure, the layers of atoms can slip over each other without the overall bonding being broken.

This is why the are malleable (hammered into shape) and ductile (drawn into wire).

3.Adding another element, metal or non-metal, disrupts this 'slip' effects so the alloy is stronger than the pure metal.

In the diagram above the blue circles can represent a larger metal atom with a radius > iron, and the purple circles a non-metallic atom with a radius < iron atoms.

To avoid over-repetition, you should also read ...

The chemical bonding in metals - giant lattice structure

Explaining the properties of metals using the metallic bonding model

Alloys: Theory, improved design and problems using metals e.g. fatigue and corrosion

General basic GCSE level notes on transition metals

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