|
Part 2
Electronic structure & Ionization Energy page sub-index: 2.1
The modern Periodic Table * 2.2
The electronic structure of atoms (including s p d f
subshells/orbitals/notation) * 2.3 Electron configurations of elements (Z = 1
to 56) * 2.4 Electron configuration and the
Periodic Table * 2.5 Electron configuration of
ions and oxidation states * 2.6 Spectroscopy and
the hydrogen spectrum * 2.7 Evidence of quantum
levels from ionisation energies
Advanced Periodic Table Index *
Part 1
A brief Periodic Table history *
the modern Periodic Table
*
Part 3
Period 1 survey : 1. Hydrogen
:
2. Helium : Summary of
Period 1 : heavier element
formation-stellar nuclear fusion *
Part 7
s-block metals Gps 1/2 Alkali/Alkaline Earth Metals *
Part 10
3d-block Sc-Zn and Transition Metals *
Part 11
Group and Series data summaries
and links to periodicity plots
2.1 The modern
version of the Periodic Table is based on the electronic structure of
atoms
-
With our knowledge
of atomic structure the modern Periodic Table is now laid out in order of
atomic/proton number (Z) and any apparent anomalies sorted out.
-
The atomic/proton number of the nucleus
(Z) decides which element the
atom is, the number of electrons surrounding the nucleus and hence the element's chemistry
which is based on the number of electrons and their arrangement. The
full Periodic Table (Z = 1 to 112) is shown further down
the page with the element symbol, atomic/proton number (Z) and another
version of the Periodic Table (Z = 1 to 56) showing
the electron configuration which is explained in the next
section 2.2.
-
Due to
isotopic mass variations and their nuclear stability, the relative atomic mass does
sometimes go 'up/down' as you proceed through the Periodic Table. See
Part 1 Mendeleev's Periodic
table work.
-
The use and
function of the Periodic Table will never cease! Newly 'man-made'
elements, beyond uranium (Z=92), are being 'synthesised' in nuclear reactors
and cyclotrons.
See GCSE nuclear reactions
and radioactivity
page.
-
We now know the electronic structure of elements and can
understand how the electrons are arranged in principal and
sub-electronic levels and the 'quantum rules' of electron structure
are understood.
-
This knowledge now
allows us to understand why the Periodic Table makes sense in terms
of the known chemistry of the elements, and their subsequent
classification, prior to the discovery and understanding of the
significance of the sub-atomic particles, particularly the
proton and electron and their 'arrangement' in an atom.
-
Mendeleev and his
contemporaries central ideas on classifying elements, despite some
errors and omissions (i.e. not discovered), are now fully vindicated
by our knowledge of the electronic structure of atoms. Mendeleev's
powerful intuition on 'element patterns' was brought to full
fruition by Rutherford and his contemporaries in discovering the
secrets of the atom and quantum physicists elucidating the 'quantum
patterns' of how multi-electron systems function.
-
For the simplified version of
expressing electronic arrangements up to atomic number 20 and the
relationship of the element in the Periodic Table, see the
GCSE Atomic Structure Notes. Its not a bad idea to revise the
basics before getting stuck into the advanced stuff!
2.2 The
'detailed' electronic
structure of the atoms of the various elements
The details
required by different pre-university syllabuses as regards background
theory and orbital knowledge seems to vary quite a lot, so I've done by
best to cater for all of them. If you wish to go straight to working out the s, p,
d electron configuration of an element, click here!
-
How to use the advanced
s, p, d (f) notation for the electron configuration/arrangement of atoms/ions
is outlined below, but no knowledge of quantum mechanics is
required, but you do need to know how to work out electron
arrangements from the rules and a little knowledge of the shape of
orbitals wouldn't go amiss! You do NOT need to know the origin of
the rules or know all about the four quantum numbers, BUT I can't
stand pulling rules out of a hat, so I have given a little
theoretical introduction, if can't stand that, tough!
-
To accurately
describe an electron in an atom requires four quantum numbers which
arise from solutions to the elaborate mathematical equations of quantum mechanics,
which describe the exceedingly complex wave behaviour of electrons.
-
These four quantum numbers
arise from solutions to the complex equations which describe the wave
and quantised behaviour of electrons surrounding the nucleus.
-
The first three
quantum numbers have 0 or +/- integer values and the fourth one is +/-
1/2)
-
The Pauli
exclusion principle states that no electron in an atom can have the
same four quantum numbers, i.e. at least one must differ from
electron to electron for a single atom.
-
The four
quantum numbers are:
-
The principal quantum energy level
number n or shell (n = 1,2, 3 ...), often just referred to as 'the
level'. It is important to think of this as the principal
energy level, i.e. the principal quantum level an electron can
occupy.
-
The
subsidiary/azimuthal/angular quantum number, l, this defines the
'spatial' type
of sub-shell orbital, (l = 0 to n-1). often just
referred to as 'the sub-level or more specifically the
s/p/d/f sub-level' (see orbital diagrams later). Again, it is
important to think of this as a sub-energy level of an
electron.
-
For s
orbital (l = 0), p orbital
(l = 1), d orbital (l = 2) diagrams below, and
for the f orbital (l = 3).
-
For a
given principal quantum number the order of energy of the
sub-level is s < p < d < f.
-
The magnetic
or spatial orientation
(of the orbital) quantum number,
m,
in terms of x,y,z axis (m = -l ... 0 ... l)
-
where l
= the azimuthal quantum number 2. above and allows for each
principal quantum level n, one s orbital for n = 1, 2, 3 etc.,
three p orbitals per for n = 2, 3, 4 etc., 5 d orbitals for n
=3, 4, 5 etc. and seven orbitals for n = 4, 5, 6 etc.
-
See the
orientation of the three p type orbitals and the five d type
orbitals.
-
The electron's spin,
s,
which has the value of +1/2 or -1/2
and can be envisaged as the electron spinning
clockwise/anti-clockwise in a full individual orbital.
-
Electrons
possess spin and if an orbital is filled then the pair of
electrons must have opposite spins (spin-paired).
-
This due to
Pauli exclusion principle, which states that no electron can
have the same four quantum numbers, since the other three
quantum numbers would be the same for a specific orbital, it is
the spin quantum number which will differ (+/- 1/2).
-
The principal
quantum electronic energy levels (n) can be split into sub-levels denoted
by s, b, d and f depending on the number of electrons in the
'system'.
-
The 'space' in
which the electron exists with its particular quantum level energy is
called the atomic orbital and each type, s, p, d or
f has its
own particular 'shape' or 'shapes'.
-
Each individual atomic
orbital can 'hold' a maximum of two electrons.
-
s, p and d orbital
diagrams.
-
Don't worry too
much about all the 'quantum' details above, the important
features to
appreciate are described below.
-
To sum up
'numerically' from the quantum
number rules, for the principal quantum number n ...
-
Each atomic
orbital can hold a maximum of two electrons.
-
For each
principal quantum level n, the following rules apply ...
-
for n = 1,
there is just
one sub-shell: 1s, maximum of 2 electrons,
-
for n = 2 there are two sub-shells: 1 x 2s atomic orbital and 3 x 2p orbitals, maximum of 2
+ 6 = 8 electrons,
-
for n
= 3 there are three sub-shells: 1 x 3s,3 x 3p
orbitals and 5 x 3d orbitals, maximum of 2 + 6 + 10 = 18 electrons,
-
for n
= 4 there are four sub-shells: 1 x 4s,3 x 4p
orbitals, 5 x 4d orbitals and 7 x 4f orbitals, maximum of 2 + 6 + 10
+ 14 = 32 electrons.
-
However the order of filling is not this simple (see below,
with visual diagrammatic help).
-
The arrangement of
electrons in the shells and orbitals is called the electronic
configuration or electron arrangement/structure and is written out in a particular sequence.
-
The orbital
electrons are denoted in the form of e.g. 2p3
-
means there
are three electrons (super-script number 3)
-
in the p sub-shell (the lower case letter)
-
and in the second principal quantum
level/shell (prefix number 2).
-
The quantum
levels and associated orbitals are filled according to the
Aufbau
Principle which states that an electron goes into the lowest available
energy level providing the following 'sub-rules' are obeyed.
-
The Pauli exclusion principle
states that no two electrons can have the same four quantum numbers.
-
Hund's Rule of
maximum multiplicity states that, as far as is possible, electrons will
occupy orbitals so that they have parallel spins. This means if a set of
sub-shell orbitals of the same energy level e.g. a 2p or 3d set, each
orbital will be singly occupied before pairing (to minimise electron
repulsion within a single atomic orbital, i.e. a lower energy state than
paired electron orbitals and unoccupied orbitals.
-
The orbitals are
filled in a definite order to produce the system of lowest energy and
any electron will go into the lowest available energy level.
-
The order of 'filling'
for an electron configuration is shown in the diagram below.
-
It uses is
a simple diagrammatic
convention to show an atomic orbital as a box.
-
Electrons are
shown as
half-arrows (up/down to represent the different spin quantum number s),
see the 2nd diagram.
-
-
The order of
filling (up to atomic number Z = 36, H to Kr) is 1s 2s 2p 3s 3p
4s 3d 4p, up to a total 36 electrons from Z = 1 to 36 i.e. the
order of increasing energy of the subshell or energy sub-level.
-
Note the 'quirk' in order
for filling the 3d sub-shell energy level (see
also the diagram below).
-
Until atomic number 21 (Sc) is reached, the
3d level is too high in energy and the electrons go into the 4s level
and then the 3d level is filled from Sc to Zn.
-
This, and
other 'quirks' I'm afraid, are a feature of the quantum
complexity of multi-electron systems, so just learn the rules
and get on with life!
-
After Z=30, the 'filling' of
the 4p level begins with Ga (Z=31) to Kr (Z=36). After Z=36, and up to Z=56,
so after 4p the
filling order is, 5s 4d 5p 6s, thus completing period and starting
period 6 (and also repeating the pattern of filling in period 4
including a 2nd block of metals, the 4d block.
-
The diagram
for vanadium (Z=23), 1s22s22p63s23p63d34s2
is shown below.
-
 *
-
Just a thought
experiment do the following ...
-
The
table below shows how
they are written out up to Z = 56 and a few others and note the orbital order when
writing out.
-
They are
written out in strict order of principal quantum number 1, 2, 3 etc. and each
principal quantum number is followed
by the s, p or d sub-levels etc., and this is irrespective of
the order of filling, i.e. when writing out the configuration, you
ignore the 3d filling 'quirk' described above.
-
Also in the table,
some
are written out in box diagram format, each box represents an orbital
with a maximum of two electrons of opposite spin (shown by the
opposing arrows). Note the electrons only pair up when all
sub-orbitals are filled separately with a single electron (this minimises electron pair
repulsion within an orbital).
-
Elements with one
or two outer s electrons, and no outer p or d electrons etc., are called
s-block
elements (Groups 1 and 2). -
Elements with at least one outer p electron
are called p-block elements (Groups 3 to 8/0).
-
Elements where the
highest available d
sub-shell is being filled are called d-block elements (*Transition
Metals) and similarly elements where the highest available f sub-shell is being filled are called
f-block elements (the Lanthanides and Actinides).
-
Quantum theory dictates
that electrons can only have certain specific 'quantised' energies and any
electronic level change requires a specific energy change.
-
Any electron will occupy
the lowest available energy level according to the
Aufbau principle (previously described).
-
The order of 'filling' up
to atomic number 56 from the lowest to highest quantum level is ...
-
Writing out
electron configurations for atoms: -
To work out an
electron arrangement for an atom, you start with the atomic number, then
'fill in' the levels and sub-levels according to the rule.
-
The electron
configuration is written out in order of, firstly, the principal quantum
energy level, then within this level in s, p, d, f order and
the total number
of electrons in each sub-energy level is shown as a super-script.
-
Example 1. sodium,
Na, Z = 11
-
1s filled (2e) 9e
left, 2s filled (2e) 7e left, 2p filled (6e) 1e left, last e goes into
the 3s level. According to the notation rule this is written as ...
-
1s22s22p63s1
(2.8.1 in simplified shell notation)
-
Example 2. vanadium,
V, Z = 23
-
1s filled (2e) 21e
left, 2s filled (2e) 19e left, 2p filled (6e) 13e left, 3s filled (2e)
11e left, 3p filled (6e) 5e left, 4s filled (2e) 3e left, last 3e go
into 3d level. According to the notation rule this is written as ...
-
1s22s22p63s23p63d34s2
(2.8.11.2 in simplified shell notation)
-
Example 3. bromine,
Br, Z = 35
-
Filling in the first
18e as described
in example 2. will give an argon structure (1s22s22p63s23p6),
which can be abbreviated to [Ar], the next 2e go into the 4s level (15e
left), the next 10e go into the 3d level, the final 5e go into the 4p
level.
-
[Ar]3d104s24p5
(2.8.18.7 in simplified notation) Note the use of 'noble gas notation' as an abbreviation for all the
filled inner sub-shells making up the equivalent of noble gas electron
arrangement, and will not include the 'outer electrons').
2.3
List of the Electronic Configuration of Elements 1 to 56 using the
advanced notation
The rules of how to assign
electrons in multi-electron atoms to the appropriate quantum levels
is 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).
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.
The energy level filling
order up to Z = 56 is 1s 2s 2p 3s 3p 4s 3d
4p (for Z = 1 to 36) 5s 4d 5p 6s 4f 5d (for Z = 37 to
56)
|
Atomic
number Z and the element name and symbol |
Electron configuration |
Box diagram of outer
electron orbitals
(representing the superscripted electron numbers beyond the inner noble gas core
in [He/Ne/Ar/Kr] which is never involved in chemical bonding/reactions) |
Symbol,
group/series/block and Comments |
|
1
Hydrogen, H |
1s1 |
1s |
H, no
Gp really, a bit unique! |
|
2
Helium, He |
1s2 = [He] |
2s |
He, Gp
0/18 Noble Gas, |
|
3
Lithium, Li |
1s22s1
(simple notation: 2.1) |
[He]2s2p |
Li,
s-block, Gp1 Alkali Metal, |
|
4
Beryllium, Be |
1s22s2
(2.2) |
[He]2s2p |
Be,
s-block, Gp2 Alkaline Earth Metal, |
|
5 Boron, B |
1s22s22p1
(2.3) |
[He]2s2p |
B,
p-block, Gp3/13 |
|
6
Carbon, C |
1s22s22p2
(2.4) |
[He]2s2p |
C,
p-block, Gp4/14, |
|
7
Nitrogen, N |
1s22s22p3
(2.5) |
[He]2s2p |
N,
p-block, Gp5/15, |
|
8
Oxygen, O |
1s22s22p4
(2.6) |
[He]2s2p |
O,
p-block, Gp6/16, |
|
9
Fluorine, F |
1s22s22p5
(2.7) |
[He]2s2p |
F,
p-block, Gp7/17 Halogen, |
|
10 Neon, Ne |
1s22s22p6 = [Ne]
(2.8) |
[He]2s2p |
Ne,
p-block, Gp 0/18 Noble Gas, |
|
11
Sodium, Na |
1s22s22p63s1
(2.8.1) |
[Ne]3s3p |
Na, Gp1
Alkali Metal, |
|
12
Magnesium, Mg |
1s22s22p63s2
(2.8.2) |
[Ne]3s3p |
Mg,
s-block, Gp2 Alkaline Earth Metal, |
|
13
Aluminium, Al |
1s22s22p63s23p1
(2.8.3) |
[Ne]3s3p |
Al,
p-block, Gp3/13, |
|
14
Silicon, Si |
1s22s22p63s23p2
(2.8.4) |
[Ne]3s3p |
Si,
p-block, Gp4/14, |
|
15
Phosphorus, P |
1s22s22p63s23p3
(2.8.5) |
[Ne]3s3p |
P,
p-block, Gp5/15, |
|
16
Sulphur, S |
1s22s22p63s23p4
(2.8.6) |
[Ne]3s3p |
S,
p-block, Gp6/16, |
|
17
Chlorine, Cl |
1s22s22p63s23p5
(2.8.7) |
[Ne]3s3p |
Cl,
p-block, Gp7/17 Halogen, |
|
18
Argon, Ar |
1s22s22p63s23p6
= [Ar] (2.8.8) |
[Ne]3s3p |
Ar,
p-block, Gp 0/18 Noble Gas, |
|
19
Potassium, K |
1s22s22p63s23p64s1
(2.8.8.1) |
[Ar]3d 4s4p |
K,
s-block, Gp1 Alkali Metal, |
|
20
Calcium, Ca |
1s22s22p63s23p64s2
(2.8.8.1) |
[Ar]3d4s4p |
Ca,
s-block, Gp2 Alkaline Earth Metal, |
|
21
Scandium, Sc |
1s22s22p63s23p63d14s2 |
[Ar]3d 4s4p |
Sc, 3d
block, not true Transition Metal |
|
22
Titanium, Ti |
1s22s22p63s23p63d24s2 |
[Ar]3d 4s4p |
Ti, 3d
block, a true Transition Metal |
|
23
Vanadium, V |
1s22s22p63s23p63d34s2 |
[Ar]3d 4s4p |
V, 3d
block, a true Transition Metal |
|
24
Chromium, Cr |
1s22s22p63s23p63d54s1 |
[Ar]3d 4s4p |
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]3d 4s4p |
Fe, 3d
block, a true Transition Metal |
|
27
Cobalt, Co |
1s22s22p63s23p63d74s2 |
[Ar]3d 4s4p |
Co, 3d
block, a true Transition Metal |
|
28
Nickel, Ni |
1s22s22p63s23p63d84s2 |
[Ar]3d 4s4p |
Ni, 3d
block, a true Transition Metal |
|
29
Copper, Cu |
1s22s22p63s23p63d104s1 |
[Ar]3d 4s4p |
Cu, 3d
block, a true Transition Metal |
|
30 Zinc, Zn |
1s22s22p63s23p63d104s2 |
[Ar]3d4s4p |
Zn, 3d
block, not true Transition Metal |
|
31
Gallium, Ga |
[Ar]3d104s24p1 |
[Ar]3d4s4p |
Ga,
p-block, Gp3/13, |
|
32
Germanium, Ge |
[Ar]3d104s24p2 |
[Ar]3d4s4p |
Ge,
p-block, Gp4/14, |
|
33
Arsenic, As |
[Ar]3d104s24p3 |
[Ar]3d4s4p |
As,
p-block, Gp5/15, |
|
34
Selenium, Se |
[Ar]3d104s24p4 |
[Ar]3d4s4p |
Se,
p-block, Gp6/16, |
|
35
Bromine, Br |
[Ar]3d104s24p5 |
[Ar]3d4s4p |
Br,
p-block, Gp7/17 Halogen, |
|
36
Krypton, Kr |
[Ar]3d104s24p6
= [Kr] |
[Ar]3d4s4p |
Kr,
p-block, Gp 0/18 Noble Gas, |
|
37
Rubidium, Rb |
[Kr]5s1 |
[Kr]5s |
Rb,
s-block, Gp1 Alkali Metal, |
|
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 true Transition Metal |
|
40
Zirconium, Zr |
[Kr]4d25s2 |
[Kr]4d5s |
Zr, 4d
block, a true Transition Metal |
|
41
Niobium, Nb |
[Kr]4d45s1 |
[Kr]4d 5s |
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 true Transition Metal |
|
49
Indium, In |
[Kr]4d105s25p1 |
[Kr]4d5s5p |
In,
p-block, Gp3/13, |
|
50 Tin, Sn |
[Kr]4d105s25p2 |
[Kr]4d5s5p |
Sn,
p-block, Gp4/14, |
|
51
Antimony, Sb |
[Kr]4d105s25p3 |
[Kr]4d5s5p |
Sb,
p-block, Gp5/14, |
|
52
Tellurium, Te |
[Kr]4d105s25p4 |
[Kr]4d5s5p |
Te,
p-block, Gp6/16, |
|
53
Iodine, I |
[Kr]4d105s25p5 |
[Kr]4d5s5p |
I,
p-block, Gp7/17 Halogen, |
|
54
Xenon, Xe |
[Kr]4d105s25p6
= [Xe] |
[Kr]4d5s5p |
Xe,
p-block, Gp 0/18 Noble Gas, |
|
55
Caesium, Cs |
[Xe]6s1 |
[Xe]6s |
Cs,
s-block, Gp1 Alkali Metal, |
|
56
Barium, Ba |
[Xe]6s1 |
[Xe]6s |
Ba,
s-block, Gp2 Alkaline Earth Metal,
|
|
57 Lanthanum, La |
[Xe]5d16s2 |
[Xe]5d6s |
La, start of the Lanthanide metal 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 |
2.4 Electron
configuration and the Periodic Table
| Pd |
s block |
3d/4d blocks of Transition Metals (Periods 4/5), the 1st/10th
are NOT true
transition elements, they have no partially filled d shell in an
ion. |
p block |
| Gp1 |
Gp2 |
Gp3 |
Gp4 |
Gp5 |
Gp6 |
Gp7 |
Gp0 |
| 1 |
1H
1s1
|
2He
1s2 |
| 2 |
3Li [He]2s1 |
4Be
[He]2s2 |
The electronic structure of Elements
1 to 56, ZSymbol,
Z = atomic or proton
number = total electrons in neutral atom,
[He] = 1s2, [Ne] = 1s22s22p6,
[Ar] = 1s22s22p63s23p6, [Kr] = 1s22s22p63s23p63d104s24p6
Between Groups 2 and 3 (13)
are the d-blocks and f-blocks where the quantum energy level
rules permit their inclusion and electron filling. Periods 4 and
5 have 18 elements each, including the 3d and 4d blocks
respectively. |
5B
[He]2s22p1 |
6C
[He]2s22p2 |
7N
[He]2s22p3 |
8O
[He]2s22p4 |
9F
[He]2s22p5 |
10Ne
[He]2s22p6 |
| 3 |
11Na
[Ne]3s1 |
12Mg
[Ne]3s2 |
13Al
[Ne]3s23p1 |
14Si
[Ne]3s23p2 |
15P
[Ne]3s23p3 |
16S
[Ne]3s23p4 |
17Cl
[Ne]3s23p5 |
18Ar
[Ne]3s23p6 |
| 4 |
19K
[Ar]4s1 |
20Ca
[Ar]4s2 |
21Sc
[Ar] 3d14s2 |
22Ti
[Ar] 3d24s2 |
23V
[Ar] 3d34s2 |
24Cr
[Ar] 3d54s1 |
25Mn
[Ar] 3d54s2 |
26Fe
[Ar] 3d64s2 |
27Co
[Ar] 3d74s2 |
28Ni
[Ar] 3d84s2 |
29Cu
[Ar] 3d104s1 |
30Zn
[Ar] 3d104s2 |
31Ga
[Ar] 3d104s24p1 |
32Ge
[Ar] 3d104s24p2 |
33As
[Ar] 3d104s24p3 |
34Se
[Ar] 3d104s24p4 |
35Br
[Ar] 3d104s24p5 |
36Kr
[Ar] 3d104s24p6 |
| 5 |
37Rb [Kr]5s1 |
38Sr
[Kr]5s2 |
39Y [Kr] 4d15s2 |
40Zr
[Kr] 4d25s2 |
41Nb
[Kr] 4d45s1 |
42Mo
[Kr] 4d55s1 |
43Tc
[Kr] 4d55s2 |
44Ru
[Kr] 4d75s1 |
45Rh
[Kr] 4d85s1 |
46Pd
[Kr] 4d10 |
47Ag
[Kr] 4d105s1 |
48Cd
[Kr] 4d105s2 |
49In
[Kr] 4d105s25p1 |
50Sn
[Kr] 4d105s25p2 |
51Sb
[Kr] 4d105s25p3 |
52Te
[Kr] 4d105s25p4 |
53I
[Kr] 4d105s25p5 |
54Xe
[Kr] 4d105s25p6 |
| 6 |
55Cs [Xe]6s1 |
56Ba
[Xe]6s2 |
4f-block (14) and 5d-block
(10) 32 elements in period 6 including the Lanthanide Series of
Metals. |
81Tl [Xe] 5d106s26p1 |
82Pb [Xe]
5d106s26p2 |
83Bi [Xe] 5d106s26p3 |
84Po [Xe] 5d106s26p4 |
85At
[Xe] 5d106s26p5 |
86Rn
[Xe] 5d106s26p6 |
| 7 |
87Fr [Rn]7s1 |
88Ra
[Rn]7s2 |
5f-block and 6d-block
including the Actinide Series of Metals. |
|
|
|
|
|
|
-
Note:
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('8 ') can also be numbered as Groups 13 to 18 to fit in with the
actual number of vertical columns of elements. 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.
-
Groups of
elements:
-
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.
-
Blocks of
elements:
-
The s-block
consists of Groups 1 and 2 where the only outer electrons are in an s
sub-energy level orbital.
-
The p-block
correspond to Groups 3 to 8 where the three p sub-energy level orbitals are being
filled.
-
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 3d block. Similarly on period 5 there is a 4d block
where the 4d sub-shell level is filled.
-
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.
-
Series of
elements:
-
The 1st
Transition Metals and other 'horizontal blocks' are sometimes called
a 'series' but they are better described as the '3d block' or '3d
series of elements' 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! (for more details see
???? working on!)
-
The
full Periodic Table is shown below.
| Pd |
s
block |
3d to
6d blocks of Transition Metals (Periods 4
to 7), note that the 1st (d1) and 10th (d10) are NOT true
transition elements. |
p block |
| Gp1 |
Gp2 |
Gp3 |
Gp4 |
Gp5 |
Gp6 |
Gp7 |
Gp0 |
| 1 |
1H
Note: (i) H does not readily fit into any group, (ii) He not
strictly a 'p' element but does belong in Gp 0/8
|
2He |
| 2 |
3Li |
4Be |
The
full IUPAC modern Periodic Table of Elements (ZSymbol, z = atomic or proton
number) |
5B |
6C |
7N |
8O |
9F |
10Ne |
| 3 |
11Na |
12Mg |
13Al |
14Si |
15P |
16S |
17Cl |
18Ar |
| 4 |
19K |
20Ca |
21Sc |
22Ti |
23V |
24Cr |
25Mn |
26Fe |
27Co |
28Ni |
29Cu |
30Zn |
31Ga |
32Ge |
33As |
34Se |
35Br |
36Kr |
| 5 |
37Rb |
38Sr |
39Y |
40Zr |
41Nb |
42Mo |
43Tc |
44Ru |
45 | |
s atomic orbital
p orbitals
d
atomic orbitals