INORGANIC Part 2
sub–index: 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
s–block Groups 1/2 Alkali Metals/Alkaline Earth Metals *
Part 8
p–block 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 sub–indexes 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 of a neutral
atom 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 'man–made'
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
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/IGCSE Atomic Structure Notes.

Its not a bad idea to revise the
basics before getting stuck into the advanced stuff!

BUT, HOW DO WE GET TO THE MODERN
PERIODIC TABLE?  read on
2.2
Orbitals and the electronic structure of the atoms
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.

Orbital
diagram notes:

The diagrams are
NOT to scale and are somewhat simplified.

These are from
theoretical calculations based on the probability functions of the
peculiar behaviour of electrons from the deep realms of quantum
mechanics! Don't worry about it!

These mathematical
functions giving rise to an electron probability distribution e.g.
illustrated by the pictures below of s, p and d orbitals.

They only give a
very approximate representation of electron density.

Each orbital, that is the
space a particular quantum level occupies, can hold a maximum of two
electrons of opposite spin quantum number (+/– ^{1}/_{2})

Quantum physicists
would say that these picture are not real, its all matrix
mathematics really, BUT chemists like pictures, and pictures can
often help students understand difficult concepts and most
importantly, use the concepts to describe chemical systems and
predict properties of atoms and molecules etc.

s atomic orbital diagram

s orbitals
have a spherical shell shape and the faint dark blue circle represents
in cross–section, the region of maximum electron density.

Only one s orbital exists for
each principal quantum number denoted by 1s, 2s, 3s etc.

s sublevels have one orbital

*

p orbitals
diagram

p orbitals
are pairs of 'dumb–bells' aligned along the x, y and z axis at 90^{o}
to each other.

There are three p
orbitals for each principal quantum number from 2 onwards denoted by 2p,
3p and 4p etc.

e.g. 2p can be
composed of 2p_{x}, 2p_{y} and 2p_{z} if all
three orbitals for a particular principal quantum number are occupied.

If a p sub–shell is
full it holds a maximum of 3 x 2 = 6 electrons.

There is no 1p
because quantum rules do not allow this.

p sublevels have three orbitals

*

d
atomic orbital diagrams

d orbitals
have complex shapes, I say no more except their relative alignment is
important in explaining the origin of
colour in transition metal complexes.

There are five d
orbitals for each principal quantum number from 3 onwards denoted by 3d,
4d, 5d etc.

If a d sub–shell is
full it contains a maximum of 5 x 2 = 10 electrons.

There are no 1d or
2d quantum levels, the quantum rules do not permit these.

d sublevels have five orbitals

f orbitals
– orbital shapes not relevant at this level, the first is the 4f
level and there are 7 orbitals holding a maximum of 7 x 2 = 14
electrons if the sub–shell is full.

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).

A summary of how many electrons
can occupy a particular level (first four principal quantum levels
only)

Principal quantum energy level 
type of quantum sublevel orbital 
maximum electrons in quantum sublevel 
maximum electrons in principal quantum level 
1 
s 
2 

2 
s 
2 
8 
p 
6 
3 
s 
2 
18 
p 
6 
d 
10 
4 
s 
2 
32 
p 
6 
d 
10 
f 
14 
 BUT, this is not necessarily the order in
which they are filled from principal quantum level 3 onwards, so beware
and read on!

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, electron structure or electron configuration and is
written out in a particular sequence.

sp d f orbital notation are used in
writing out electron configurations for chemical elements and their ions

The orbital
electrons are denoted in the form of e.g....

(i) 1s^{2}

means there are two electrons (super–script number
^{2})

in the s subshell orbital
(lower case letter)

and in the first principal quantum
level/shell (prefix number 1).

(ii) 2p^{3}

means there
are three electrons (super–script number ^{3})

in the p sub–shell orbitals (the lower case letter)

and in the second principal quantum
level/shell (prefix number 2).

(iii) 3d^{7}

means there
are seven electrons (super–script number ^{7})

in the d sub–shell orbitals (the lower case letter)

and in the third principal quantum
level/shell (prefix number 3).

(iv) 5f^{11}

means there
are eleven electrons (super–script number ^{11})

in the f sub–shell orbitals (the lower case letter)

and in the 5th principal quantum
level/shell (prefix number 5).



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. for 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 before
filling a higher level orbital.

This is known as the aufbau
principle.

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 'spin' box.

Electrons are
shown as
half–arrows (up/down to represent the different spin quantum number s),
see the 2nd diagram.


A note on the 'spin boxes'
representing subshell orbitals:

and
would represent a singly and doubly occupied s orbital

would represent three empty p subshell orbitals, referred to as a
vacant orbital

would represent two halffilled p orbitals

would represent three full p subshell orbitals

would represent five empty d orbitals of a particular d subshell

would represent three halffilled d orbitals and two vacant.

would represent three completely filled d orbitals and tow halffilled

would represent five completely filled d orbitals of a particular d
subshell

I've used this style of diagram on
the next page were the electron configurations of elements 1 to 58 are
listed.

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 'quantum 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) 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.


Above is the electron spin box
diagram showing all the empty levels available from 1s to 6s.

Shown below are various electron
spin box diagrams derived from the above diagram and employing the
aufbau principle, Pauli's exclusion principle and Hund's rule.

carbon (Z = 6) 1s^{2}2s^{2}2p^{2} (group 4
element)

neon (Z = 10) 1s^{2}2s^{2}2p^{6} (group 0/18
noble gas element)

silicon (Z = 14) 1s^{2}2s^{2}2p^{6}3s^{2}3p^{2}
(group 4 element)

argon (Z = 18) 1s^{2}2s^{2}2p^{6}3s^{2}3p^{6}
(group 0/18 noble gas element)

calcium (Z = 18) 1s^{2}2s^{2}2p^{6}3s^{2}3p^{6}4s^{2}
(group 2 metal element)

vanadium (Z=23) 1s^{2}2s^{2}2p^{6}3s^{2}3p^{6}3d^{3}4s^{2}
(3d block transition metal)

For just a thought
experiment, do the following ...

'Empty' the 3d
level of electron arrows and you get the diagram for calcium (Z = 20).

Fill up
completely the
3d and 4p boxes with arrows and you get krypton (Z = 36)

and learn to fill up anything
in between!

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 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).

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 0, modern notation
Groups 13 to 18).

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 i.e. energy
quanta absorbed or energy quanta emitted.

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

BUT first, how to
work out the electron arrangement from the atomic/proton number,

AND then how to write
out the electron configuration.

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.

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 sub–shells making up the equivalent of noble gas electron
arrangement, and will not include the 'outer electrons').

–
TOP OF PAGE
How do you work out the electron arrangement configuration for 1
Hydrogen, H ? How do you work out the electron arrangement
configuration for 2
Helium, He ? How do you work out the electron arrangement
configuration for 3
Lithium, Li ? How do you work out the electron arrangement
configuration for 4
Beryllium, Be ? How do you work out the electron arrangement
configuration for 5 Boron, B ? How do you work out the electron arrangement
configuration for 6
Carbon, C ? How do you work out the electron arrangement
configuration for 7
Nitrogen, N ? How do you work out the electron arrangement
configuration for 8
Oxygen, O ? How do you work out the electron arrangement
configuration for 9
Fluorine, F ? How do you work out the electron arrangement
configuration for 10 Neon, Ne ? How do you work out the electron arrangement
configuration for 11
Sodium, Na ? How do you work out the electron arrangement
configuration for 12
Magnesium, Mg ? How do you work out the electron arrangement
configuration for 13
Aluminium, Al ? How do you work out the electron arrangement
configuration for 14
Silicon, Si ? How do you work out the electron arrangement
configuration for 15
Phosphorus, P ? How do you work out the electron arrangement
configuration for 16
Sulphur, S ? How do you work out the electron arrangement
configuration for 17
Chlorine, Cl ? How do you work out the electron arrangement
configuration for 18
Argon, Ar ? How do you work out the electron arrangement
configuration for 19
Potassium, K ? How do you work out the electron arrangement
configuration for 20
Calcium, Ca ? How do you work out the electron arrangement
configuration for 21
Scandium, Sc ? How do you work out the electron arrangement
configuration for 22
Titanium, Ti ? How do you work out the electron arrangement
configuration for 23
Vanadium, V ? How do you work out the electron arrangement
configuration for 24
Chromium, Cr ? How do you work out the electron arrangement
configuration for 25
Manganese, Mn ? How do you work out the electron arrangement
configuration for 26 Iron, Fe ? How do you work out the electron arrangement
configuration for 27
Cobalt, Co ? How do you work out the electron arrangement
configuration for 28
Nickel, Ni ? How do you work out the electron arrangement
configuration for 29 ? How do you work out the electron arrangement
configuration for
Copper, Cu ? How do you work out the electron arrangement
configuration for 30 Zinc, Zn ? How do you work out the electron arrangement
configuration for 31
Gallium, Ga ? How do you work out the electron arrangement
configuration for 32
Germanium, Ge ? How do you work out the electron arrangement
configuration for 33
Arsenic, As ? How do you work out the electron arrangement
configuration for 34
Selenium, Se ? How do you work out the electron arrangement
configuration for 35
Bromine, Br ? How do you work out the electron arrangement
configuration for 36
Krypton, Kr ? How do you work out the electron arrangement
configuration for 37
Rubidium, Rb ? How do you work out the electron arrangement
configuration for 38
Strontium, Sr ? How do you work out the electron arrangement
configuration for 39
Yttrium, Y ? How do you work out the electron arrangement
configuration for 40
Zirconium, Zr ? How do you work out the electron arrangement
configuration for 41
Niobium, Nb ? How do you work out the electron arrangement
configuration for 42
Molybdenum, Mo ? How do you work out the electron
arrangement configuration for 43
Technetium, Tc ? How do you work out the electron arrangement
configuration for 44
Ruthenium, Ru ? How do you work out the electron arrangement
configuration for 45
Rhodium, Rh ? How do you work out the electron arrangement
configuration for 46
Palladium, Pd ? How do you work out the electron arrangement
configuration for 47
Silver, Ag ? How do you work out the electron arrangement
configuration for 48
Cadmium, Cd ? How do you work out the electron arrangement
configuration for 49
Indium, In ? How do you work out the electron arrangement
configuration for 50 Tin, Sn ? How do you work out the electron arrangement
configuration for 51 ? How do you work out the electron arrangement
configuration for
Antimony, Sb ? How do you work out the electron arrangement
configuration for 52
Tellurium, Te ? How do you work out the electron arrangement
configuration for 53
Iodine, I ? How do you work out the electron arrangement
configuration for 54
Xenon, Xe ? How do you work out the electron arrangement
configuration for 55
Caesium, Cs ? How do you work out the electron arrangement
configuration for 56
Barium, Ba ? How do you work out the electron arrangement
configuration for 57 Lanthanum, La ? How do you work out the electron
arrangement configuration for 58 Cerium, Ce ? chemistry revision notes how to work out the electron
configuration of an element in terms of s p d f orbital quantum levels
AS AQA GCE A level chemistry how to work out the electron configuration
of an element in terms of s p d f orbital quantum levels AS Edexcel GCE
A level chemistry how to work out the electron configuration of an
element in terms of s p d f orbital quantum levels AS OCR GCE A level
chemistry how to work out the electron configuration of an element in
terms of s p d f orbital quantum levels AS Salters GCE A level chemistry
how to work out the electron configuration of an element in terms of s p
d f orbital quantum levels US grades 11 & 12 chemistry how to work out
the electron configuration of an element in terms of s p d f orbital
quantum levels notes for revising how to work out the electron
configuration of an element in terms of s p d f orbital quantum levels s
p d f spdf orbital notes
for AQA AS chemistry, s p d f spdf orbital notes
for Edexcel AS chemistry, s p d f spdf orbital notes for OCR AS chemistry A,
s p d f spdf orbital notes for OCR Salters AS chemistry B,
s p d f spdf orbital notes for AQA A level chemistry, s p d
f spdf orbital notes for Edexcel A level chemistry,
s p d f spdf orbital notes for OCR A level chemistry
A, s p d f spdf orbital notes for OCR Salters A
level chemistry B s p d f spdf orbital notes for US Honours grade 11 grade
12 s p d f spdf orbital notes for preuniversity chemistry courses, spdf
notation for 1 Hydrogen, H ?, spdf notation for 2 Helium, He ?, spdf
notation for 3 Lithium,
Li ?, spdf notation for 4 Beryllium, Be ?, spdf notation for 5 Boron, B ?,
spdf notation for 6 Carbon, C
?, spdf notation for 7 Nitrogen, N ?, spdf notation
for 8 Oxygen, O ?, spdf notation for 9
Fluorine, F ?, spdf notation for 10 Neon, Ne ?, spdf
notation for 11 Sodium, Na ?, spdf notation for 12
Magnesium, Mg ?, spdf notation for 13 Aluminium, Al ?, spdf
notation for 14 Silicon, Si ?, spdf notation for 15 Phosphorus, P ?, spdf
notation for 16 Sulphur, S ?, spdf notation for 17 Chlorine,
Cl ?, spdf notation for 18 Argon, Ar ?, spdf notation for 19 Potassium, K ?,
spdf notation for 20
Calcium, Ca ?, spdf notation for 21 Scandium, Sc ?, spdf
notation for 22 Titanium, Ti ?, spdf notation for 23 Vanadium, V ?, spdf
notation for 24 Chromium, Cr ?, spdf notation for 25 Manganese,
Mn ?, spdf notation for 26 Iron, Fe ?, spdf notation for 27 Cobalt, Co ?,
spdf notation for 28 Nickel, Ni
?, spdf notation for 29 ?, spdf notation for Copper, Cu ?,
spdf notation for 30 Zinc, Zn ?, spdf notation for 31
Gallium, Ga ?, spdf notation for 32 Germanium, Ge ?, spdf
notation for 33 Arsenic, As ?, spdf notation for 34 Selenium, Se ?, spdf
notation for 35 Bromine, Br ?, spdf notation for 36 Krypton,
Kr ?, spdf notation for 37 Rubidium, Rb ?, spdf notation for 38 Strontium, Sr ?,
spdf notation for 39
Yttrium, Y ?, spdf notation for 40 Zirconium, Zr ?, spdf
notation for 41 Niobium, Nb
?, spdf notation for 42 Molybdenum, Mo ?, spdf
notation for 43 Technetium, Tc ?, spdf notation for
44 Ruthenium, Ru ?, spdf notation for 45 Rhodium, Rh ?, spdf
notation for 46 Palladium, Pd
?, spdf notation for 47 Silver, Ag ?, spdf notation for 48 Cadmium, Cd ?,
spdf notation for 49 Indium, In
?, spdf notation for 50 Tin, Sn ?, spdf notation for 51 ?,
spdf notation for Antimony, Sb ?, spdf notation for 52
Tellurium, Te ?, spdf notation for 53 Iodine, I ?, spdf
notation for 54 Xenon, Xe ?, spdf notation for 55
Caesium, Cs ?, spdf notation for 56 Barium, Ba ?, spdf
notation for 57 Lanthanum, La ?, spdf notation for
58 Cerium, Ce
TOP OF PAGE
















Website content © Dr
Phil Brown 2000+. All copyrights reserved on revision notes, images,
quizzes, worksheets etc. Copying of website material is NOT
permitted. Exam revision summaries & references to science course specifications
are unofficial. 
