INORGANIC Part 2
subindex: 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) * 2.3
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
Level Inorganic Chemistry Periodic Table Index *
Part 1
Periodic Table history
* Part 2
Electron configurations, spectroscopy,
hydrogen spectrum,
ionisation energies *
Part 3
Period 1 survey H to He *
Part 4
Period 2 survey Li to Ne * Part
5 Period 3 survey Na to Ar *
Part 6
Period 4 survey K to Kr and important trends
down a group *
Part 7
sblock Groups 1/2 Alkali Metals/Alkaline Earth Metals *
Part 8
pblock Groups 3/13 to 0/18 *
Part 9
Group 7/17 The Halogens *
Part 10
3d block elements & Transition Metal Series
*
Part 11
Group & Series data & periodicity plots * All
11 Parts have
their own subindexes near the top of the pages
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 electron configuration.

The
full Periodic Table (Z = 1 to 112) is shown
in section 2.4 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 introduced and 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.

The use and
function of the Periodic Table will never cease! Newly 'manmade'
elements, beyond uranium (Z=92), are being 'synthesised' in nuclear reactors
and cyclotrons.
See GCSE/IGCSE nuclear reactions
and radioactivity
pages

We now know the electronic structure of elements and can
understand how the electrons are arranged in principal and
subelectronic 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 subatomic 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 multielectron 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/IGCSE Atomic Structure Notes.
2.2
Orbitals and the electronic structure of the atoms
The details
required by different preuniversity 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 subshell orbital, (l = 0 to n1). often just
referred to as 'the sublevel or more specifically the
s/p/d/f sublevel' (see orbital diagrams later). Again, it is
important to think of this as a subenergy 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
sublevel 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/anticlockwise in a full individual orbital.

Electrons
possess spin and if an orbital is filled then the pair of
electrons must have opposite spins (spinpaired).

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 sublevels 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 subshell: 1s, maximum of 2 electrons,

for n = 2 there are two subshells: 1 x 2s atomic orbital and 3 x 2p orbitals, maximum of 2
+ 6 = 8 electrons,

for n
= 3 there are three subshells: 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 subshells: 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).

How do we work out
electron the arrangement of an atom?

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. 2p^{3}

means there
are three electrons (superscript number ^{3})

in the p subshell (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 'subrules' 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
subshell 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
halfarrows (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 sublevel.

Note the 'quirk' in order for
filling the 3d subshell 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 multielectron 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) and finishes with 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), 1s^{2}2s^{2}2p^{6}3s^{2}3p^{6}3d^{3}4s^{2}
is shown below.

*

Just a thought
experiment do the following ...

The
table in Part 2.3 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 sublevels 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).

Elements with one
or two outer s electrons, and no outer p or d electrons etc., are called
sblock
elements (Groups 1 and 2). 
Elements with at least one outer p electron
are called pblock elements (Groups 3 to 8/0).

Elements where the
highest available d
subshell is being filled are called dblock elements (*Transition
Metals) and similarly elements where the highest available f subshell is being filled are called
fblock 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 sublevels according to the rule.

Example 1. sodium,
Na, Z = 11

1s filled (2e) 9e's
left, 2s filled (2e's) 7e left, 2p filled (6e's) 1e left, last electron goes into
the 3s level.

According to the notation rule this is written as ...

1s^{2}2s^{2}2p^{6}3s^{1}
(2.8.1 in simplified shell notation)

Example 2. vanadium,
V, Z = 23

1s filled (2e's) 21e's
left, 2s filled (2e's) 19e left, 2p filled (6e's) 13e's left, 3s filled (2e's)
11e's left, 3p filled (6e's) 5e's left, 4s filled (2e's) 3e's left, last 3e's go
into 3d level.

According to the notation rule this is written as ...

1s^{2}2s^{2}2p^{6}3s^{2}3p^{6}3d^{3}4s^{2}
(2.8.11.2 in simplified shell notation)

Example 3. bromine,
Br, Z = 35

Filling in the first
18e's as described
in example 2. will give an argon structure (1s^{2}2s^{2}2p^{6}3s^{2}3p^{6}),
which can be abbreviated to [Ar], the next 2e's go into the 4s level (15e's
left), the next 10e's go into the 3d level, the final 5e's go into the 4p
level.

[Ar]3d^{10}4s^{2}4p^{5}
(2.8.18.7 in simplified notation)

Note the use of 'noble gas notation' as an abbreviation for all the
filled inner subshells making up the equivalent of noble gas electron
arrangement, and will not include the 'outer electrons').


Exam Revision Tuition A Level
Revision Guides for A Level Courses Examinations for GCE Advanced Level Theoretical
Physical Chemistry A level Revision Notes for GCE Advanced
Subsidiary Level AS Advanced Level A2 IB
Revise AQA GCE Chemistry OCR GCE Chemistry Edexcel GCE Chemistry Salters
Chemistry CIE Chemistry, WJEC GCE AS A2 Chemistry, CCEA/CEA GCE AS A2 Chemistry revising courses for preuniversity students
(equal to US grade 11 and grade 12 and AP Honours/honors level courses)
Website
content copyright © Dr W P Brown 20002012 All rights reserved
on
revision notes, puzzles, quizzes, worksheets, xwords etc. * Copying of website
material is not permitted * I do not personally endorse the adverts 
but they do pay for the site!
Alphabetical Index for Science
Pages Content
A
B C D
E F
G H I J K L M
N O P
Q R
S T
U V W
X Y Z 