Doc Brown's Advanced Level Chemistry - Pre-University Inorganic Chemistry - Periodic Table Revision Notes

Part 2 Electronic structure, spectroscopy & ionisation energies,

electron configurations of the elements for Z = 1 to 58

Section 2.3 Electron configurations for elements of atomic number Z = 1 to 56

Parts 2.1 to 2.2 should be read and studied first!

Sub-index for this page

Part 2.3 uses the rules on assigning electron arrangements, and how the quantum level notation is written out, and using boxes to represent orbitals, as well as the usual written orbital notation, is given for elements Z = 1 to 56.

e.g. 26 Iron, Fe

1s22s22p63s23p63d64s2

[Ar]3d4s4p

Part 2.4 The relationship between electron configuration and the Periodic Table and uses the electron configurations to show how the Periodic Table arises, i.e. an element's position in the Periodic Table and hence its chemistry, is primarily determined by the arrangement of its outer valency electrons.

Part 2.5 shows how to work out the electron configuration of ions, (positive cations or negative anions formed by the loss or gain of valence electrons) and relating electron arrangements to the oxidation states exhibited by selected elements.

Appendix: The electron configuration of all 118 elements of the Periodic Table

See also on other separate pages part ...

2.1 The electronic basis of the modern Periodic Table

2.2 The electronic structure of atoms (including s p d f subshells/orbitals/notation) 

2.6 Spectroscopy and the hydrogen spectrum

2.7 Evidence of quantum levels from ionisation energies

2.8 Emission and absorption spectra of elements

 2.3 List of the Electronic Configuration of Elements Z = 1 to 56 using the advanced s, p, d and f notation

YOU MUST STUDY Parts 2.1 and 2.2 before studying section 2.3 onwards – The rules of how to assign electrons in multi–electron atoms to the appropriate quantum levels were explained in section 2.2.

The list below quotes the ground state electron configurations i.e. the lowest available state according to the Aufbau principle (previously described). The order of filling the electron levels is listed below and also indicated on the diagram below.

Electron Box diagrams of the outer electron arrangement and examples of the simple electron notation (e.g. 2.8.1) are also included, with brief comments in the end right hand column e.g. element symbol, group, series etc.

The electrons–in–boxes notation for subshells: Boxes are used to represent an individual orbital or set of orbitals in the electrons are shown as arrows. The pairs up/down arrows represent a full orbital with electrons of opposite spin and note how the half–filled boxes/orbitals illustrate Hund's rule of maximum multiplicity

Energy level filling order to Z = 56 is 1s 2s 2p 3s 3p 4s 3d 4p (for Z = 1 to 36) then 5s 4d 5p 6s 4f/5d varies (for Z = 37 to 56)

However, when writing out the electron configuration you must write them out in order of strict principal quantum with the accompanying s, p, d, f notation

order of writing out is: 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 6s (up to Z = 58)

BUT the order of orbital filling is: 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s (up to Z = 56)

So at Z = 21 Sc, you start to fill the 3d orbitals (NOT 4p orbitals),

then at Z = 31 Ga you start to fill the 4p orbitals, and this is all you need pre-university! I think?

Atomic number Z and the element name and chemical symbol	Electron configuration Electron arrangement s, p, d & f notation with electron number superscripts (plus some simplified electron arrangements)	Electron spin box diagrams of the outer electron orbitals for the electron configuration of the atom representing the superscripted electrons beyond the inner noble gas core [He/Ne/Ar/Kr], the latter are not involved in chemical bonding or reactions.	Symbol, group/series/block and Comments Gp = Group!
1 Hydrogen, H	1s1	1s	H, no Group really, a bit unique!
2 Helium, He	1s2 = [He]	1s very stable, unreactive, filled shell	He, Group 0/18 Noble Gas
3 Lithium, Li	1s22s1 (simple notation: 2.1)	[He]2s2p empty sub-shell	Li, s–block, Gp1 Alkali Metal, v. reactive
4 Beryllium, Be	1s22s2 (2.2)	[He]2s2p empty sub-shell	Be, s–block, Gp2 Alkaline Earth Metal,
5 Boron, B	1s22s22p1 (2.3)	[He]2s2p	B, p–block, Group 3/13
6 Carbon, C	1s22s22p2 (2.4)	[He]2s2p	C, p–block, Group 4/14
7 Nitrogen, N	1s22s22p3 (2.5)	[He]2s2p	N, p–block, Group 5/15
8 Oxygen, O	1s22s22p4 (2.6)	[He]2s2p	O, p–block, Group 6/16
9 Fluorine, F	1s22s22p5 (2.7)	[He]2s2p	F, p–block, Group 7/17 Halogen
10 Neon, Ne	1s22s22p6 = [Ne] (2.8)	[He]2s2p very stable, unreactive, filled outer shell	Ne, p–block, Group 0/18 Noble Gas
11 Sodium, Na	1s22s22p63s1 (2.8.1)	[Ne]3s3p empty sub-shell	Na, Gp1 Alkali Metal, v. reactive
12 Magnesium, Mg	1s22s22p63s2 (2.8.2)	[Ne]3s3p empty sub-shell	Mg, s–block, Gp2 Alkaline Earth Metal,
13 Aluminium, Al	1s22s22p63s23p1 (2.8.3)	[Ne]3s3p	Al, p–block, Group 3/13
14 Silicon, Si	1s22s22p63s23p2 (2.8.4)	[Ne]3s3p	Si, p–block, Group 4/14
15 Phosphorus, P	1s22s22p63s23p3 (2.8.5)	[Ne]3s3p	P, p–block, Group 5/15
16 Sulfur, S	1s22s22p63s23p4 (2.8.6)	[Ne]3s3p	S, p–block, Group 6/16
17 Chlorine, Cl	1s22s22p63s23p5 (2.8.7)	[Ne]3s3p	Cl, p–block, Group 7/17 Halogen
18 Argon, Ar	1s22s22p63s23p6 = [Ar] (2.8.8)	[Ne]3s3p very stable, unreactive, filled outer shell	Ar, p–block, Group 0/18 Noble Gas
19 Potassium, K	1s22s22p63s23p64s1 (2.8.8.1)	[Ar]3d4s4p	K, s–block, Gp1 Alkali Metal, v. reactive
20 Calcium, Ca	1s22s22p63s23p64s2 (2.8.8.1)	[Ar]3d4s4p	Ca, s–block, Gp2 Alkaline Earth Metal
21 Scandium, Sc	1s22s22p63s23p63d14s2	[Ar]3d4s4p	Sc, 3d block, not a true Transition Metal
22 Titanium, Ti	1s22s22p63s23p63d24s2	[Ar]3d4s4p	Ti, 3d block, a true Transition Metal
23 Vanadium, V	1s22s22p63s23p63d34s2	[Ar]3d4s4p	V, 3d block, a true Transition Metal
24 Chromium, Cr	1s22s22p63s23p63d54s1	[Ar]3d4s4p	Cr, 3d block, a true Transition Metal
25 Manganese, Mn	1s22s22p63s23p63d54s2	[Ar]3d4s4p	Mn, 3d block, a true Transition Metal
26 Iron, Fe	1s22s22p63s23p63d64s2	[Ar]3d4s4p	Fe, 3d block, a true Transition Metal
27 Cobalt, Co	1s22s22p63s23p63d74s2	[Ar]3d4s4p	Co, 3d block, a true Transition Metal
28 Nickel, Ni	1s22s22p63s23p63d84s2	[Ar]3d4s4p	Ni, 3d block, a true Transition Metal
29 Copper, Cu	1s22s22p63s23p63d104s1	[Ar]3d4s4p	Cu, 3d block, a true Transition Metal
30 Zinc, Zn	1s22s22p63s23p63d104s2	[Ar]3d4s4p	Zn, 3d block, not a true Transition Metal
31 Gallium, Ga	[Ar]3d104s24p1	[Ar]3d4s4p	Ga, p–block, Group 3/13
32 Germanium, Ge	[Ar]3d104s24p2	[Ar]3d4s4p	Ge, p–block, Group 4/14
33 Arsenic, As	[Ar]3d104s24p3	[Ar]3d4s4p	As, p–block, Group 5/15
34 Selenium, Se	[Ar]3d104s24p4	[Ar]3d4s4p	Se, p–block, Group 6/16
35 Bromine, Br	[Ar]3d104s24p5	[Ar]3d4s4p	Br, p–block, Group 7/17 Halogen
36 Krypton, Kr	[Ar]3d104s24p6 = [Kr] (2.8.18.8)	[Ar]3d4s4p very stable, filled outer shell	Kr, p–block, Group 0/18 Noble Gas
37 Rubidium, Rb	[Kr]5s1	[Kr]5s	Rb, s–block, Gp1 Alkali Metal, v. reactive
38 Strontium, Sr	[Kr]5s2	[Kr]5s	Sr, s–block, Gp2 Alkaline Earth Metal,
39 Yttrium, Y	[Kr]4d15s2	[Kr]4d5s	Y, 4d block, not a true Transition Metal
40 Zirconium, Zr	[Kr]4d25s2	[Kr]4d5s	Zr, 4d block, a true Transition Metal
41 Niobium, Nb	[Kr]4d45s1	[Kr]4d5s	Nb, 4d block, a true Transition Metal
42 Molybdenum, Mo	[Kr]4d55s1	[Kr]4d5s	Mo, 4d block, a true Transition Metal
43 Technetium, Tc	[Kr]4d55s2	[Kr]4d5s	Tc, 4d block, a true Transition Metal
44 Ruthenium, Ru	[Kr]4d75s1	[Kr]4d5s	Ru, 4d block, a true Transition Metal
45 Rhodium, Rh	[Kr]4d85s1	[Kr]4d5s	Rh, 4d block, a true Transition Metal
46 Palladium, Pd	[Kr]4d10	[Kr]4d5s	Pd, 4d block, a true Transition Metal
47 Silver, Ag	[Kr]4d105s1	[Kr]4d5s5p	Ag, 4d block, a true Transition Metal
48 Cadmium, Cd	[Kr]4d105s2	[Kr]4d5s5p	Cd, 4d block, not a true Transition Metal
49 Indium, In	[Kr]4d105s25p1	[Kr]4d5s5p	In, p–block, Group 3/13
50 Tin, Sn	[Kr]4d105s25p2	[Kr]4d5s5p	Sn, p–block, Group 4/14
51 Antimony, Sb	[Kr]4d105s25p3	[Kr]4d5s5p	Sb, p–block, Group 5/14
52 Tellurium, Te	[Kr]4d105s25p4	[Kr]4d5s5p	Te, p–block, Group 6/16
53 Iodine, I	[Kr]4d105s25p5	[Kr]4d5s5p	I, p–block, Group7/17 Halogen
54 Xenon, Xe	[Kr]4d105s25p6 = [Xe]	[Kr]4d5s5p very stable, filled outer shell	Xe, p–block, Group 0/8/18 Noble Gas
55 Caesium, Cs	[Xe]6s1	[Xe]6s	Cs, s–block, Gp1 Alkali Metal, v. reactive
56 Barium, Ba	[Xe]6s2	[Xe]6s	Ba, s–block, Gp2 Alkaline Earth Metal,
57 Lanthanum, La	[Xe]5d16s2	[Xe]5d6s	La, start of 5d–bock and Lanthanide Series
58 Cerium, Ce	[Xe]4f26s2 not 4f15d16s2	things get a bit less systematic in the f blocks	Ce, 1st of f–block in the Lanthanides Metals
************************	****************************************	*********************************************************	*****************************************************

The electron spin box diagrams can be used to show the full electronic structure e.g.

All based on the right-hand diagram

More 'quantum level quirks'!

A note on two anomalies in the 3d block, namely the transition metals chromium and copper:

Cr is [Ar]3d⁵4s¹ and not [Ar]3d⁴4s²

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

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

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 2.4 Electron configuration and the Periodic Table

Not all the elements are shown but the position of s, p, d and f blocks are shown and explained after the table

It is an element's electron configuration that determines whether it is an s, p, d or f block element.

Instead of 'simply' assigning an element to its group or series because of its chemical properties e.g. formula of similar compounds, we can now assign an element's position in the periodic table using its electron configuration.

An s block element has one or two outer electrons i.e. ... s¹ or ... s²  e.g. group 1 and group 2 metals (shown below)

A p block element has 1-6 outer p electrons beyond the s² orbital i.e. ... s²p¹ to s²p⁶  e.g. group 3/13 across to group 0/18 (shown below)

A d block element has 1 to 10 electrons in the outermost set of d orbitals i.e. ...d¹ to ...d¹⁰  e.g. 3d block of Sc to Zn, 4d block Y to Cd (shown below)

Note that a true transition element has an ion with an incomplete d sub-shell (3d, 4d etc.)

An f block element has 1 to 14 electrons in the outermost set of f orbitals i.e. ...f¹ to ...f¹⁴  e.g.  Ce to Lu (NOT shown below)

This partial periodic table relates an element's electron configuration to the element's position in the periodic table.

You can then see the patterns between an atom's electron arrangement and the group, block or series the element belongs to.

• Note on Group numbers

  • Using 0 to denote the Group number of Noble Gases is very historic now since compounds of xenon known exhibiting a valency of 8.

  • Because of the horizontal series of elements e.g. like the Sc to Zn block (10 elements), Groups 3 to 0 can also be numbered as Groups 13 to 18 to fit in with the actual number of vertical columns of elements and this is the modern trend in periodic table notation.

  • This can make things confusing, but there it is, classification is still in progress!

  • The atomic/proton number, decides which element an atom is and the outer electron structure decides which group/block/series the element belongs to and ultimately its chemistry.

  • The positions of the s, p, d and f blocks are also indicated Periodic Table above and arise from the quantum rules

    • s block elements have an outer shell of just 1 or 2 s electrons i.e. s¹ or s² configuration beyond an inner noble gas configuration (2 per period from period 2 onwards), that is groups 1 and 2. Technically, hydrogen and helium are in the s block, but bare little chemical similarity with the group 1 and group 2 metals.

    • p block elements have an outer electron configuration of s²p¹ to s²p⁶ i.e. elements where the p sub-shell is being filled (6 per period from period 2 onwards), that is groups 3/13 to 0/18.

    • d block elements eg the 3d (Sc-Zn) where the 3d sub-shell is being filled and like wise for the 4d block (Y-Cd),10 elements per block per period from period 4 onwards, the first horizontal blocks of metals which lie between the s block and p block,

    • f blocks elements eg the 4f and 5f blocks where the f sub-shells start being filled (14 elements in each block per period from period 6 onwards).

• The most stable electron configurations

  • When the outer s and p quantum levels and any completely filled inner orbital quantum levels eg 3d or 4f, you get a particularly stable element with minimal chemical reactivity ie you get a Noble Gas element [simple electron notation in ()]

    • Z = 2, helium, 1s² = [He] (2)

    • Z = 10, neon, 1s²2s²2p⁶ = [Ne] (2.8)

    • Z = 18, argon, 1s²2s²2p⁶3s²3p⁶ = [Ar] (2.8.8)

    • Z = 36, krypton, [Ar]3d¹⁰4s²4p⁶ = [Kr] (2.8.18.8)

    • Z = 54, xenon, [Kr]4d¹⁰5s²5p⁶ = [Xe]

    • Z = 86, radon, [Xe]5d¹⁰6s²6p⁶ =[Rn]

• What is the electronic basis of Groups of elements? – their 'electronic classification'

  • For groups 1 to 2, and 'old' 3 to 0/'new' notation 13 to 18 (except He), all the elements in the same vertical column have the same outer electron configuration and therefore will be expected to have a very similar chemistry.

    • This gives the electronic basis for Mendeleev's brilliant conception of the periodic table, ie laying out all the elements in order of 'atomic weight' and lining them up to give vertical columns of chemically and physically similar elements.

  • For the d blocks of Groups 3 to 12, using the 'new' group number notation, the vertical 'group' connection of similar outer electron configuration is consistent except for V/Nb, Fe/Ru, Co/Rh, Ni/Pd where the 3d/4s and 4d/5s pairs of levels are of very similar energy and small differences in outer electron configuration occur.

    • Never–the–less these pairs of elements show strong similarities as part of the justification for denoting the Transition Metals plus Groups 4 to 0 as Groups 3 to 18.

• What is the electronic basis for the 'series of elements'? – their 'electronic classification'

  • The '1st Transition Metals Series' from Sc to Zn, and other 'horizontal blocks' are sometimes called a 'series' but they are better described as the '3d block' or '3d series of elements' (and, 4d block, 4f block – filling of 4f sub–shell etc.), but a horizontal row of elements, unlike the vertical columns of the eight vertical groups.

    • Why is 'block' better than 'series'?

      • The reference to the electronic structure is very important, the word series is a bit vague!

      • Technically, scandium (Sc, Z = 21) and zinc (Zn, Z = 30), are NOT true transition metals BUT they are true 3d block elements!

        • See Detailed Advanced Level Transition Metal Notes

• What is the overall electronic basis for blocks of elements across the whole of the periodic table?

  • The s–block consists of Groups 1 and 2 where the only outer electrons are in an s sub–energy level orbital (no outer p electrons, 2 per period).

    • Technically, the same applies to hydrogen and helium in period 1.

  • The p–block corresponds to Groups 3 to 0 (old notation) or Groups 13 to 18 (new notation) where the three p sub–energy level orbitals are being filled (6 per period).

  • Starting with period 4, where the first of the d sub–shells is low enough in energy to be filled, there are ten elements 'inserted' between groups 2 and 3, the so–called d blocks of ten elements (the 1st block, the 3d block Sc–Zn is on Period 4).

    • Therefore Sc to Zn form the head elements of Groups 3 to 12 using the 'new' group number notation.

    • Similarly on period 5 there is a 4d block where the 4d sub–shell level is filled.

    • So 10 d block elements per period are now permitted\under the quantum number rules.

  • Starting with cerium (Z=58, period 6), see in full table below, there is a further insertion of fourteen elements where the seven f–orbital sub–shell is being filled after the first of the d–block metals and similarly with thorium (Z=90) in period 7 and these are known as the f blocks (14 per period where permitted).

• The full Periodic Table is shown below without the electron configurations, but including the old/new group number notation.

Notes:

The Noble Gases have been referred to as Group 0 because they were believed not to form compounds with other elements.

However, since 1961, many compounds of xenon have been prepared including xenon(VIII) oxide, XeO₄, thus attaining the expected maximum possible oxidation state based on the number of electrons in the outer shell, so Group 18 seems most appropriate to use these days for advanced level chemistry courses.

The d block elements are sometimes referred in terms of their vertical columns as Groups 3 to 12, and the subsequent p–block group columns as Groups 13 to 18.

The s p d f blocks are shown in the Periodic Table above.

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 2.5 Electronic configuration of ions and oxidation states

How do you work out the electron arrangement of ions? How do you work out the electron configuration of ions?

In what order to you remove electrons for positive ions? In what order do you add electrons for negative ions?

• The electron configuration of ions:

  • Beware in quoting the configurations for simple ions where, although the order of removal is basically the reverse of the order for filling the energy levels, there is one important exception you should know.

    • For d/transition blocks/series the 4s electrons are 'removed' before the 3d electrons and similarly the 5s electrons are 'removed' before the 4d electrons.

  • Positive ions (cations) are formed by electron loss and the order of removal is the reverse of the order the full electron configuration is written out e.g.

    • sodium atom Na = 1s²2s²2p⁶3s¹ , sodium ion Na⁺ = 1s²2s²2p⁶ = [Ne]

    • calcium atom is Ca = 1s²2s²2p⁶3s²3p⁶4s², calcium ion Ca²⁺ = 1s²2s²2p⁶3s²3p⁶ = [Ar]

    • iron atom Fe =  [Ar]3d⁶4s², iron(II) ion Fe²⁺ = [Ar]3d⁶, and iron(III) ion Fe³⁺ = [Ar]3d⁵

    • germanium atom Ge = [Ar]3d¹⁰4s²4p², germanium(II) ion Ge²⁺ = [Ar]3d¹⁰4s², germanium(IV) ion Ge⁴⁺ = [Ar]3d¹⁰

    • For the s block elements, the maximum positive oxidation state is governed by the number of electrons removed to give a complete a noble gas structure in the covalent or ionic bonding situation.

  • Negative ions (anions) are formed by electron gain and the filling order rule is continued e.g.

    • chlorine Cl = [Ne]3s²3p⁵, chloride ion Cl^– = [Ne]3s²3p⁶ = [Ar]

    • oxygen: O = [He]2s²2p⁴, oxide ion O^2– = [He]2s²2p⁶ = [Ne]

    • phosphorus: P = [Ne]3s²3p³, phosphide ion P^3– = [Ne]3s²3p³ = [Ar]

    • For the p block groups 3/13 to 0/18 the maximum negative oxidation state is governed by the number of electrons needed to complete a noble gas structure in the covalent or ionic bonding situation (restricted to -1 to -4).

• Further comments on oxidation state and electronic structure:

  • For more details see notes on oxidation state and redox reactions.

  • The maximum oxidation state is often, but not always, limited by an inner full noble gas structure with or without a full d/f sub–shell.

  • The maximum oxidation state from Group 1 Alkali Metals (+1) to Group 0/18 Noble Gases (+8) is numerically equal to the number of outer electrons, i.e. those beyond an inner noble gas core or inner noble gas plus a full d/f sub–shell e.g.

    • The maximum oxidation states for Groups 1 to 0 in old notation (now s-block Groups 1-2 and p-block 13-18)

    • Group 1, outer electron configuration s¹, eg sodium is +1 in sodium chloride NaCl

    • Group 2, outer electron configuration s², eg magnesium is +2 in magnesium oxide MgO

    • Group 3 (13), outer electron configuration s²p¹, eg aluminium is +3 in aluminium fluoride AlF₃

    • Group 4 (14), outer electron configuration s²p², eg silicon is +4 in silicon dioxide SiO₂

    • Group 5 (15), outer electron configuration s²p³, eg phosphorus is +5 in the phosphate(V) ion PO₄^3–

    • Group 6 (16), outer electron configuration s²p⁴, eg sulfur is +6 in sulfur trioxide SO₃

    • Group 7 (17), outer electron configuration s²p⁵, eg chlorine is +7 in the chlorate(VII) ion ClO₄^–

    • Group 0 (18), outer electron configuration s²p⁶ (except He, just s²), eg xenon is +8 in xenon(VIII) oxide XeO₄

      • The maximum oxidation state pattern for the d blocks is a bit more complicated and the trend goes through a maximum e.g.

      • the maximum observed oxidation states for the 3d block and transition metal series:

      • Sc (+3), Ti (+4), V (+5), Cr (+6), Mn (+7), Fe (+3, maybe +6?), Co (+3), Ni (+3), Cu (+3), Zn (+2)

      • The maximum oxidation state from scandium to manganese (Sc-Mn) equates to the presence of 4 to 7 3d and 4s electrons beyond the inner noble gas core of [Ar], the 'theoretical' ion would have a 3d⁰ structure.

  • Atoms or ions with the same electron configuration are referred to as isoelectronic eg

    • Na⁺, Mg²⁺, Al³⁺ and Ne are all 1s²2s²2p⁶ and you might say the three ions are isoelectronic with neon.

    • N^3-, O^2- and F^- are also all isoelectronic with neon.

    • Zn²⁺ and  Ge⁴⁺ are both [Ar]3d¹⁰ and so isoelectronic

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