can't we trust atoms? Because they make up everything!
Atomic Structure, Isotopes,
Electronic Structure of Atoms, Atomic Structure Experiments, ionisation -
(and a brief
mention of allotropes in appendix 3.)
Chemistry KS4 science GCSE/IGCSE/O Level/A Level Revision Notes
Structure of Atoms – three fundamental particles
'Portrait of an atom' and particle size comparisons
1c. Atoms, sub-atomic particles,
nuclide notation & diagrams
Isotopes – definition,
nuclide notation and examples
Electronic Structure of Atoms
– rules and the Periodic Table
Which electron arrangements are stable and which are not?
The Periodic Table and Electronic Structure
Relative Atomic Mass
details on a separate calculation page
History of atomic structure models, alpha particle scattering experiment
Atomic structure diagrams
– don't confuse with isotopes!
The effect of
ionizing radiation on atoms formation of positive ions
Introduction to the mass spectrometer
(only for advanced A level students!)
Some abbreviations used: A =
mass number, Ar = relative atomic mass, Z = atomic number/proton number
Alpha particle scattering *
Atomic (proton) number
* Electron arrangement (examples)
* Electron shell rules
* ions * Isotopes
* Mass spectrometer (advanced students only) * Neutron
Nuclide symbol notation
* Periodic Table (and electron structure)
Periodic Table (its general structure)
Atomic structure and radioactivity
This page describes the
structure of atoms in terms of the three fundamental sub–atomic particles –
protons, neutrons and electrons. Isotopes are defined with examples and
nuclide notate (e.g. top right picture) is explained. Early theories and
models of atomic structure are described and explained including the Bohr
theory of the atom. The history of the development of the atomic model is
described in detail, explaining at each stage why the theory of what an atom
is, had to be changed in the light of new evidence.
electronic structure of atoms is explained via the rules on filling shells
with electrons. Which electron arrangements are stable? and Why? and which
electron arrangements are unstable giving rise to very reactive elements
like the Group 1 Alkali Metals and Group 7 Halogen elements.
The link between the Periodic Table and Electronic Structure is explained
and described with diagrams of the periodic table and electronic structure.
The important historic alpha particle scattering experiment is described,
variations of atomic structure diagrams, There is a section explaining what
allotropes are, don't confuse with
isotopes! and, on a separate page for advanced level chemistry students, the
mass spectrometer is explained with annotated diagrams and explanatory
These notes on atomic structure are
designed to meet the highest standards of knowledge and understanding required
for students/pupils doing GCSE chemistry, IGCSE chemistry, O
Level chemistry, KS4 science courses and a primer for A Level chemistry courses.
These revision notes should prove useful for the new AQA, Edexcel and OCR
GCSE (9-1) courses.
Revision notes history of atomic
structure models, isotopes definition explained, working out numbers of electrons, protons, neutrons,
revising for A level AQA GCSE chemistry, A level Edexcel GCSE chemistry, OCR
GCSE 21st century science, A level OCR GCSE Gateway science GCSE 9-1
Structure of Atoms
– three fundamental particles
An element consists of one type of atom only.
Therefore, elements are the simplest substances that we can use and
investigate in chemistry because an element cannot be split into other substances (unlike
Each element has identical atoms (except for isotopes,
different numbers of neutrons - explained later) which are physically and
chemically identical and each element has its own unique physical and
Ever element has its own unique chemical symbol which is
used to denote elements in the periodic table, in chemical formulae and
chemical equations e.g. hydrogen is H, copper Cu, chlorine Cl or potassium
K. The symbol is a single capital letter (upper
case e.g. C, N, O, F, C, P etc.) or a capital letter followed by a lower
case letter (e.g. Cu, Fe, Cl, Br, Li etc.).
ARE ATOMS? and WHAT DO WE MEAN BY FUNDAMENTAL PARTICLES? (sub–atomic particles)
is the smallest particle of a substance, an element, which
can have its own characteristic properties AND cannot be split into
However, why do we have different elements?
Is an atom the simplest particle we need to know about to
In order to answer these questions we must
look a bit deeper into the fundamental structure of matter, that is
everything around you!
Atoms are the smallest particles of matter whose
properties we study in Chemistry.
Every element or compound is comprised of
atoms. All the atoms are the same in the structure of an element
(ignoring isotopes - different numbers of neutrons, see later) and two or more different atoms/elements must be
present in a compound.
element has its own chemical symbol (carbon C, oxygen O, sodium Na
etc.), which with added numbers (e.g. right), can be used to indicate
the composition of an atom in terms of protons, electrons and neutrons.
All of this will be explained in detail below
Initially, once the concept of an atom was
established, it was assumed that atoms were indestructible and not divisible
into smaller particles, but merely combined in different proportions to give
the range of compounds we know about e.g. Dalton's atom model.
However from experiments done in the late
19th and early 20th century it was deduced that atoms
made up of three fundamental or sub–atomic particles called protons, neutrons and
electrons, which are listed below with their
relative masses and electrical charges.
the history of the development of the atomic model is described in detail.
WHAT ARE THE CHARACTERISTIC PROPERTIES OF THESE
WHAT IS THE NUCLEUS? WHAT ARE NUCLEONS?
The three fundamental
particles of which atoms are composed
The table gives the relative
mass and electric charge of the three sub–atomic particles known as the
proton, neutron and electron
the nucleus, a nucleon
In the nucleus, a nucleon
1/1850 or 0.00055
NOT a nucleon. Electrons are arranged in energy levels or shells
in orbit around the nucleus
Protons and neutrons are much heavier
You can think of the mass of
an electron as about 1/2000th of the mass of a
proton or neutron, so, a pretty small mass BUT they occupy most of the space of
You should also realise because of the relatively small mass of
the electrons most of an atom's mass is in the nucleus.
values of 1/1836 quoted for the relative mass of an electron, but don't
worry about it, there are different ways/scales on which an electron's mass
has been calculated.
The actual mass of a proton or neutrons
is ~1.67 x 10-27 kg (~1.67 x 10-24 g)
The mass of an electron is ~9.1 x 10-31 kg
(~9.1 x 10-28 g)
The mass of an atom varies
from about 1 x 10-20 to 1 x 10-18 kg (1 x 10-23 to 1 x 10-21 g)
depending on the element
The radius of the nucleus ranges from about 1
x 10-16 to 1 x 10-14 m (1 x 10-7
to 1 x 10-5 nm) depending on the
The diameter of atoms varies from about 1 x 10-10
to 5 x 10-10 m (0.1 to 0.5 nm) depending on the
Generally speaking the radius of an atom is
about 10,000 times that of the nucleus!
A typical relatively small molecule would be
no bigger than ~1 x 10-10 to 1 x 10-9 m
(~0.1 to 1 nm)
TOP OF PAGE and
What can we say about 'A
Portrait of an Atom'?
Images of what you
normally can't see! Size comparison of atoms, molecules and cells!
The picture of 'atomic structure,
illustrated below is the result of many developing 'atomic theory' backed up by
successive generations of experimental results. This is the best picture we have
(at least for GCSE and A Level chemistry courses!).
However this diagram, which is
based on the Bohr model of atomic structure, although
more realistic in terms of the real size of the nucleus compared to the atom
as a whole, so it is not convenient to give a brief diagrammatic picture of the
composition of an atom.
The central nucleus of
protons and neutrons (most of the mass of an atom) is extremely small even compared to the size of an
atom. The rest of the 'almost empty space' of an atom is occupied by the
negative electrons, held by, and moving around the positive nucleus in their
energy levels or 'shells'.
The electrons are also pretty tiny in
mass too, compared to a proton or neutron, but the volume the electrons and
their energy levels occupy, determines the size of the atom, but an atom mainly
empty space with the nucleus at the centre!
Bohr theorised the negative electrons can
only exist in certain specific energy levels (shells) held in place by the
positive nucleus (see section on the history of development of the atomic
model). These are shown in the above diagram, but fully explained later on
All of these theories must, and have been, backed up by
repeated and varied experiments.
As each new experiment was/is done, it must
support the current theory or the theory needs to be modified to take into
account new discoveries.
Some of these important experiments are described
further down the page.
Even new experimental findings written up in research
papers should be thoroughly peer reviewed, that is checked by
scientists of at least equal academic ranking to the researchers. That's
how science works!
The size of an atom compared to other
The size of an individual atom is around
0.1 nm or 1 x 10-10 m
and the size of the nucleus itself is only
about 1 x 10-14 m, about 1/10000th of the radius of the atom.
(nm = nanometres, 1 nm = 10-9
Most of the mass is in the centre of
the atom, that is the nucleus,
which has a radius of around 1/10000th of the
This means the radius of the nucleus
is about 1 x 10-14 m (0.00001 nm), pretty small !!, but still consists
of most of the mass of an atom!!!
If you look at the table of size
comparison below, its not until you get to a human hair can we see
clearly something with the naked eye. The width of a human hair is
approximately 106 times that of an atom (a million times
bigger) and 1010 times bigger than a nucleus (ten thousand
million times bigger). You can of course see cells under examination
with an optical microscope, but these are over 500 000 times bigger than
an individual atom. You can, however, observe atoms using an electron
comparison data table of particles sizes/dimensions
of atoms, molecules, nanoparticles and other 'things'
typical small protein
||silver or titanium dioxide nano-particles
typical virus e.g. cold virus
||a typical carbon nanotube
typical eukaryotic cell
width of human hair
|Size in nm - diameter or length
0.3 x 0.6
longest length or diameter m
| 2 x
6 x 10-10
3 x 10-10
0.5 to 1.0
na means not applicable
TOP OF PAGE and
1c. Atoms, sub-atomic particles,
nuclide notation and diagrams
The number of protons in
the nucleus of an atom decides what element that atom is.
e.g. if the atom has 3
protons in the nucleus, it cannot be anything except lithium!
Elements consist of one
type of atom only determined by the proton number (atomic number).
ALL terms explained below
Some more concise and
handy styles to show the atomic composition of the same lithium atom
- What sub–atomic particles make up atoms?
What is their mass and charge?
- The diagram
above of a 'portrait of an atom' gives some idea on the
structure of an atom (sometimes called the Bohr Atomic Model), it also includes some important definitions and
notation used to describe atomic structure
- REMINDER of the three fundamental particles you need to
know are ...
- proton: particle mass = 1, electric charge = +1, the
charged particle in the nucleus
- neutron: particle mass = 1, charge = 0,
uncharged particle in the nucleus
- electron: particle mass = 1/1850 ~1/2000,
electric charge =
- Electrons are NOT in the nucleus but exist in electronic energy levels around the nucleus
(a sort of orbit, often described as a shell, see later).
- The nucleus of protons and neutrons is tiny,
even compared to the tiny atom!
- So most of the volume of an atom is empty space, BUT it
is where the
tiny electrons are.
- In fact the diameter of the nucleus of protons plus
neutrons is about a ten thousandth of the diameter of the whole atom!
- Since the nucleus is composed of positive
protons and neutral neutrons, the nucleus itself must be positive.
- A neutral atom carries no overall charge
because the number of positive protons equals the number of negative
electrons (both singly charged + and -), and this information is given by the atomic/proton number.
- A neutral helium atom has 2 protons and 2 electrons and
a uranium atom has 92 protons and 92 electrons. 2+ balances 2- and 92+
balances 92-, easy!
- Since a neutron is electrically neutral, the number
of neutrons in the nucleus of an atom cannot affect the total positive
nuclear charge of the protons or the number of negative electrons in the
- Protons and neutrons are the 'nucleons'
or 'sub–atomic' particles present
in the minute positive nucleus and the negative electrons are held by the positive
protons in 'orbits' called energy levels or shells.
Some important evidence for this 'picture' is
obtained from alpha particle scattering experiments (see
- Although the nucleus must be positive because
of the positive protons (neutrons are neutral) an individual atom is neutral
because the number of electrons equals the number of protons – so the
charges 'cancel out'.
- An ion particle carries an overall surplus
electric charge - positive or negative - so how are they formed?
- If electrons are removed from an atom you
get a positive ion from excess positive charge, and, if electrons are added to an atom, you get a
negative ion from excess negative charge.
- An ion, by definition, cannot be neutral and the
number of protons cannot equal the number of electrons.
atomic number (Z) is the
protons in the nucleus and is also known as the
proton number of
the particular element and it is this number that decides what element a
particular atom is.
- Each element has its own atomic number, so
all the atoms of a particular element have the same atomic number.
In a neutral atom, the number of electrons
equals the atomic/proton number.
- It is the proton/atomic number (Z) that determines the
number of electrons an element has, its specific electron structure and
therefore the specific identity of a particular element in terms of its
physical and chemical properties.
- It cannot be overemphasised that it is the
electronic structure that determines the chemical character of an
element, hence the proton/atomic number determines everything about a
mass number (A)
is also known as the
is the number of particles in the nucleus of a particular atom–isotope
(notes on isotopes – definition and examples).
The neutron number (N) =
mass number – proton/atomic number
In an individual atom the number of protons (+)
equals the number of electrons (–), that is the number of positive charges
is equal to the number of negative charges to make the atom neutral.
- In the example
in the diagram for lithium–7,
73Li is its nuclide atomic notation
If the proton number and electron number are different, the atom has an overall
surplus or deficiency of electrical charge, resulting in an electrically
charged particle called an ion e.g.
- If an atom loses 1 or 2 electrons, the
protons produce an excess of 1 or 2 units of positive charge.
- The excess positive charge on the positive ion
is written as + or 2+ etc.
- If an atom gains 1 or 2 electrons, the
extra electrons produce an excess of 1 or 2 units of negative charge.
- The excess negative charge on the negative ion
is written as - or 2- etc.
- So, reading the symbols below, as above, BUT, now taking into account
the electrical charge on the ion ..
sodium ion Na+,11 protons, 10 electrons (11-1),
12 neutrons (23-11)
positive magnesium ion Mg2+,12 protons, 10 electrons
(12-2), 12 neutrons (24-12)
the negative chloride ion Cl–,17 protons, 18 electrons
(17+1), 20 neutrons (37-17)
negative sulfide ion S2-,16 protons, 18 electrons (16+2),
16 neutrons (32-16)
- for more details and examples see
ionic bonding notes.
The electrons are
arranged in specific energy levels according to a set of rules (dealt
This description of an
atom consisting of the relatively minute nucleus of protons and neutrons
surrounded by electrons in particular shells or energy levels is sometimes
referred to as the Bohr Model of the atom, after the great Danish
scientist Niels Bohr (1885–1962), one of the brilliant founders of modern
Other examples of
interpreting the nuclide notation and definition reminders:
Top left number is the nucleon number or
mass number (A = sum of protons + neutrons =
Bottom left number is the atomic number or
proton number (Z = protons in nucleus)
Electrons = protons if
the atom is electrically neutral i.e. NOT an ion.
The neutron number N = A
– Z i.e. mass/nucleon number – atomic/proton number
Therefore from the following
'full' atomic symbols, assuming we are dealing with electrically
neutral atoms, the number of sub-atomic particles for the following
atoms will be as follows ...
atom (isotope cobalt–59), mass
59, 27 protons, 32 neutrons (59 – 27), 27 electrons
atom (isotope californium–246), mass 246, 98 protons, 148 neutrons (246 – 98),
So, at this point we had
better explain, slightly belatedly, what isotopes are!
TOP OF PAGE and
ISOTOPES explained, nuclide notation
of nuclide symbols and how to
WHAT ARE ISOTOPES? ARE THEY IMPORTANT?
TOP OF PAGE and
More on the Electronic Structure of Atoms
WHAT DO WE MEAN BY the electron
configuration?, electronic structure of atoms?
that is what is the arrangement of
electrons in the shells or energy levels?
What is the relationship between an
atom's electronic structure and its position in the Periodic
The Bohr model of the atom in its more
elaborate form involves the maximum numbers of electrons that each shell or
energy level can hold and how the shells are progressively filled with
electrons from atom to another with increase in proton/atomic number.
The electrons are arranged in energy levels or shells around the nucleus and with
'orbits' on average increasing in distance from the nucleus.
Electrons in an atom occupy the lowest available energy levels (the
innermost available shells).
The lowest energy levels are
always filled first, you can think of the lower the shell, the nearer the
nucleus, and numbered 1, 2, 3 etc. as the shell gets further from the
Each electron in an atom is in a
particular energy level (or shell) and the electrons must occupy the
lowest available energy level (or shell) available nearest the nucleus.
When the level is full, the next electron goes into the next highest level (shell) available.
There are rules to learn about the maximum number of electrons allowed in each shell and you have to be able to work out the arrangements for the first 20
elements (for GCSE students,
up to at least 36 for Advanced level students).
The 1st shell
can contain a maximum of 2 electrons (electrons 1–2)
The 2nd shell can contain a maximum of 8 electrons
The 3rd shell also has a maximum of 8 electrons
The 19th and 20th electrons go into the 4th
shell, (required limit of GCSE chemistry knowledge).
Remember the total electrons
to be arranged equals the atomic/proton number for a neutral atom.
If you know the atomic (proton) number,
you know it equals the number of electrons in a neutral atom, you then apply
the above rules to work out the electron arrangement (configuration).
For elements 1 to 20 the electron
arrangements/configurations are written out in the following manner:
- Note that each number represents the number of
electrons in a particular shell, dots or commas are used to separate the numbers of
electrons in each shell. They are written out in order of increasing average
distance from the positive nucleus which holds these negative electrons in
their energy levels (shells).
- The electron configurations or electron
summarised below with reference to the periods of the periodic table and in
order of increasing atomic number up to proton number 20.
Period 1 – elements 1 to 2
Period 2 – elements 3 to 10
Period 3 – elements 11 to 18
are denoted by 2,8,1 to 2,8,8 (1st
shells full with 2 & 8 electrons, ie 2.8)
2,8,3 would mean two electrons in
the 1st level (shell), eight in the 2nd level (full shell) and three
electrons in the 3rd outer level (shell).
Period 4 – first two elements
19 to 20
are written out as 2,8,8,1 and
2,8,8,2 (1st, 2nd, 3rd full shells with 2,8,8
2,8,8,1 would mean two electrons
in the 1st level (shell), eight in the 2nd level, eight electrons in the 3rd
level (shell) and one electron in the 4th outer level (shell).
Reminder – this is as far as
GCSE students need to know, after that things get more complicated, BUT only
for advanced level students!
For example, after element 18,
the 3rd shell can hold a maximum of 18 electrons!
The above is summarised in the
The first 20 elements of the
I've repeated the same 20 elements of the
periodic table showing simple diagrams of their electronic
structure and group numbers. Group 1 are the alkali metals, Group 7 are the
halogens and group 0 are the noble gases.
The electron shell arrangements are quoted in
numbers e.g. 2,4 for C (carbon) but you need to be able to draw electron
diagrams showing the electronic structure of the atom.
In the above table, check out the atomic
number, the lower subscript on the element symbol, and apply the rules, and
hopefully it makes sense.
Some examples of electronic diagrams are given below
and GCSE/IGCSE/O level students need to be able to work and draw the
electronic structures of the first 20 elements.
You should notice that the
number of shells used equals the period number of the element in the
They can be all worked by the
'shell filling' rules described above.
For the rest of Period 4 and
other Periods you need a more
electron configuration system up to at least Z=36 using s, p, d and f orbital
notation BUT this is for advanced A level chemistry
Examples: diagram, symbol or name of element (Atomic Number = number of
protons and the number of electrons in a neutral atom), shorthand electron arrangement
and a diagram to help you follow the numbers.
Filling 1st shell, electron level 1
2 elements only,
Period 1 of the
Filling 2nd shell, electron level 2
3 of the 8 elements of Period 2
Filling 3rd shell, electron level 3
of the 8 elements of Period 3
The first 2 elements of the 4th shell
to Kr [220.127.116.11],
start of Period 4
Only the first 2 of the 18 elements of Period 4 are shown
above, the rule for 3rd shell changes from element 21 Sc onwards
(studied at Advanced level, so GCSE students don't worry!)
A few more 'snappy' examples –
given atomic number, work out electron configuration (abbreviated to e.c.)
Z = 3 e.c. 2,1 or
Z = 7 e.c. = 2,5 or Z = 14 e.c. = 2,8,4 or Z =
19 e.c. = 2,8,8,1 etc. up to Z = 20
TOP OF PAGE and
4. Which electron arrangements are stable
and which are not?
Both atoms and ions are considered
WHY ARE SOME ELECTRON ARRANGEMENTS ARE MORE
STABLE THAN OTHERS?
WHICH ELECTRON ARRANGEMENTS ARE THE MOST
STABLE AND WHICH ELECTRON ARRANGEMENTS THE LEAST STABLE?
HOW DO ELECTRON ARRANGEMENTS RELATE TO THE
REACTIVITY OF CHEMICAL ELEMENTS?
When an atom has its outer
level full to the maximum number of electrons allowed, the atom is particularly
stable electronically and very unreactive.
This is the situation
with the Noble Gases: He is , neon is [2,8]
and argon is [2,8,8] etc.
There atoms are the most
reluctant to lose, share or gain electrons in any sort of chemical
interaction because they are so electronically stable.
For all elements most
of their chemistry is about what outer electrons do or don't!
and [2,8,8] etc. are known as the 'stable Noble Gas
arrangements', and the atoms of other elements try to attain
this sort of electron structure when reacting to become more stable.
More details on
Electron configuration notes for
Advanced Level Chemistry Students
The most reactive metals
have just one outer electron.
These are the Group
1 Alkali Metals, lithium [2,1], sodium [2,8,1],
With one outer shell
electron, they have one more electron than a stable Noble Gas
So, they readily lose the
outer electron when they chemically react to try to form (if
possible) one of the
stable Noble Gas electron arrangements – which is why atoms react in
the first place!
When Group 1 Alkali Metal atoms lose an electron they form a positive ion because
the positive proton number doesn't change, but with one negative
electron lost, there is a surplus of one + charge e.g.
sodium atom ==>
Na ==> Na+
is [2.8.1] ==> [2.8] electronically
in fundamental particles
[11p + 11e] ==> [11p + 10e]
IONS are atoms
or group of atoms which carry an overall electrical charge i.e.
not electrically neutral.
The most reactive
non–metals are just one electron short of a full outer shell.
These are the Group 7
Halogens, namely fluorine [2,7], chlorine [2,8,7] etc.
These atoms are one
electron short of a stable full outer shell and seek an 8th outer
electron to become electronically stable – yet again, this is why
They readily gain an
outer electron, when they chemically react, to form one of the
stable Noble Gas electron arrangements either by sharing electrons
(in a covalent bond) or by electron transfer forming a singly
charged negative ion (ionic bonding) e.g.
chlorine atom ==>
Cl ==> Cl–
is [2.8.7] ==> [2.8.8]
in fundamental particles
[17p + 17e] ==> [17p + 18e]
proton number of Cl doesn't change but the chloride ion carries one
extra negative electron to give the surplus charge of a single – on
For more on electron
structure and chemical changes and compound formation see ...
and for more on metal and
non–metal reactivity see
TOP OF PAGE and
The Periodic Table and
Electronic Structure – more patterns!
Selected Elements of the Periodic Table are shown below
atomic number and chemical symbol.
HOW DOES AN ELEMENT'S ELECTRON ARRANGEMENT
RELATE TO ITS POSITION IN THE PERIODIC TABLE?
The elements are laid out in order of
Atomic Number – that is the number of protons in the nucleus.
It is important to realise
that the 'chemical structure' of the periodic Table
(shown above), that is the chemical similarity of vertical groups 'like'
elements (apart from the
was known well before the electronic structure of atoms was understood.
In other words the elements are laid
out in vertical columns (groups) and horizontal rows (periods) so that
chemically (usually) VERY similar elements appear under each other – and there
is a very good electronic structure reason for this!
However, it wasn't understood why they behaved in the same way chemically
e.g. similar compound formulae and reactions etc. nor was it understood at
first why Noble Gases were so unreactive towards other elements.
the electronic structure of atoms was understood, 'electronic' theories
could then be applied to explain the chemical similarity of elements in a
vertical Group of the Periodic Table.
Originally they were
laid out in order of 'atomic weight'
(now called relative atomic mass). This is not correct for some elements now that we know their
detailed atomic structure in terms of protons, neutrons and electrons, and of
course, their chemical and physical properties in more accurate and extensive
Argon (at. no.
18, electrons 2,8,8) has a relative atomic mass of 40. Potassium (at.
no. 19, electrons 2,8,8,1) has a relative atomic mass of 39. BUT Argon, in terms
of its physical, chemical and electronic properties is clearly a Noble Gas in
Group 0. Likewise, potassium is clearly an Alkali Metal in Group 1.
Hydrogen, 1, H, does not readily fit into any group
A Group is a vertical column of chemically and physically similar elements
Group 1 The Alkali Metals (Li, Na, K etc.)
with one outer electron (one more than a Noble Gas structure),
Group 7 The Halogens (F, Cl, Br, I etc.)
with seven outer electrons (one short of a Noble Gas arrangement)
and Group 0 The Noble Gases (He, Ne, Ar etc.). The group number equals the number of electrons in the outer shell (e.g. chlorine's electron arrangement is 2.8.7, the second element down Group 7 on period 3).
A Period is a horizontal row of elements with a variety of properties (left to right goes from metallic to non–metallic elements. All the elements use the same number of electron shells which equals the period number (e.g. sodium's electron arrangement 2.8.1, the first element in Period 3).
The ten elements Sc to Zn are
called the Transition Metals Series and form part of a
period between Group 2 and Group 3 from Period 4 onwards.
Below are the electron arrangements
for elements 1 to 20 set out in
format (Hydrogen and The Transition metals etc. have been omitted). When you move down to the next period you start to fill in the next shell according to the maximum electrons in a shell rule (see previous section).
NOTE: In the most
modern periodic table notation Groups 3–7 and 0 are numbered Groups 3 to 18.
the same outer electron structure
– check out the last number from element 3 onwards.
The first element in a period has one outer
electron (e.g. sodium Na 2.8.1), and the last element has a full outer shell
argon Ar 2.8.8)
Apart from hydrogen (H, 1) and helium (He, 2) the last electron number is the group number
(in the old notation) and the number of shells used is equal to the Period
The periodicity of
elements i.e. the repetition of very chemically similar elements in a group is due to the
repetition of a
More GCSE/IGCSE notes on the Periodic Table
electronic explanations of
Advanced Level Chemistry –
electron configurations/arrangements and the Periodic Table
TOP OF PAGE and
1. The history of the atom concept, development of atomic structure models and the famous alpha particle
Leucippus and Democritus and others ~400 BC, wondered what was the result of
continually dividing a substance i.e. what was the end product or
smallest bit i.e. what was left that was indivisible – the word atomic
is from Greek adjective meaning 'not divisible'.
(some of the theory ideas described
here goes above GCSE/IGCSE level!)
The development of the 'atomic model'
is an excellent example of how new experiment evidence initiates the
need to change an existing scientific model or even come up with a
different theoretical model, in this case for atomic structure.
You should appreciate that knowledge and understanding of atomic
structure has evolved over time and as a scientific model of atomic
structure, it must explain current experimental observations and pose
questions for future investigations.
Any further change in an atomic
structure hypothesis e.g. because of new evidence, must be
re-tested out in the laboratory and the results checked by other research
groups from around the world. Plus a transparent peer group review of
any research paper to be published, which means scientists having their
work checked by other scientists. I
they don't agree, somebody has got something wrong or made false claims
or just done the experiment badly! If further experimental checks don't back
up a hypothesis, it must be modified or abandoned in favour of other
When a hypothesis is backed up by
experimental evidence from different scientists and lots of
cross-checking of results e.g. the Bohr theory of the atom, it becomes
an accepted theoretical model of an atom. Even to this day the atomic
structure model is still developing e.g. the hypothesis that neutrons
and protons are made of quarks which are held together by gluons, but at
this point I'm out of my depth! Nothing stands still in scientific
theory, even with atomic structure, and the hypotheses of modern quantum
physicists about the structure of the nucleus I find pretty 'whacky',
actually, pretty incomprehensible!
The ancient Greeks thought that
everything was made of four basic 'elements' - air, earth, fire and
water. However, at the height of the great classical Greek
The Greeks idea was not
forgotten and later revived by Boyle and Newton but with little
However, in 1808
at the beginning of the 19th century,
that all matter was made up of tiny hard particles/spheres called
Dalton also proposed (correctly) the theory
that different types of atoms (elements)
combined together to give all the different substances of the physical
world (all which of course is true, except for the 'hard solid
He also produced the first list of 'atomic
weights' (we now call relative atomic masses) on a scale based on
hydrogen – given the arbitrary value of 1 since it was lightest element
known, and, as it happens, correctly so.
He was incorrect by stating that atoms were indivisible, because we now know that atoms consist of
electrons, protons and neutrons and that atoms can be 'taken apart' by
ionisation or nuclear changes e.g.
Until the discovery of the electron, atoms were thought of as hard
indivisible spheres, but brilliant 'JJ' Thompson changed all that.
New experimental evidence led to a
new scientific model of the atom
with the discovery of the electron, recognised as the
first known 'sub-atomic particle', change was on the way.
around 1897 proposed his 'plum pudding model' theory
(picture on right) based on the growing evidence
that atoms were themselves composed of even smaller more fundamental
particles like the electron i.e. atoms
were not hard indivisible spheres, so the solid sphere had to go.
He based his model on experiments that showed that atoms
contained even smaller negatively charged particle called electrons which could
be removed from atoms using a vacuum tube and applying a high potential
difference (voltage) to a very low pressure gas. He showed that the mass
of an electron was much less than the mass of an atom and that it
had a negative electric charge. Therefore the
'hard indivisible sphere' model of an atom was wrong.
From his experiments Thomson
envisaged a plumb pudding atom consisting of a positively charged
'pudding' (a sort of ball of positive charge) with just enough lighter negatively charged electrons embedded
in it to produce a neutral atom. Note that both the positive charges and
negative charges are evenly distributed through the sphere of the
atom (shown later by Rutherford and Bohr etc. to be completely wrong).
The idea of positive particles balancing the negative
correct but the relative size and nature of the nucleus and distribution
of electrons were not, BUT it was a more advanced model.
Rutherford, assisted by Hans Geiger and Ernest Marsden
(the latter two were students of Rutherford at Cambridge
University) conducted alpha particle scattering experiments
(1902–1910, and described in detail below). The famous alpha particle experiment in 1909 was designed to test the
plum pudding theory of JJ Thomson.
By 1911, these experiments established
(i) minute nature of the nucleus even
compared to the size of an atom.
(ii) the nucleus was positive and
the positive charge varied from element to element.
(iii) the positive charge was concentrated in the nucleus and able
to deflect other positive particles e.g. alpha particles.
famous Rutherford and Geiger–Marsden alpha particle scattering experiment
particle beams are fired on very thin layers of metals (e.g. very fine
gold leaf) some rather surprising results were
scientists of the early 20th century.
By using a 360o
charged particle detection system it was found that ...
particles passed through un–deflected
(as if there was nothing there!), this was expected.
This was predicted from JJ Thompson's plumb pudding
model, but all the alpha particles were expected to pass
through or to be slightly detected (observation 2.
below), though NOT big deflections.
proportion were deflected slightly
something there!), this again was not unexpected.
about 1 in 20,000 were 'bounced' back through an angle of
in other words were reflected backwards, a
totally unexpected result
and quite shocked the experimenters - not what they were
expecting. This was because the JJ Thompson's plumb
predicted the positive charge was spread out and not
sufficiently concentrated to cause, for some alpha
particles, a 180o deflection! So, whatever was
there, was substantial in mass and positive charge to cause
the repulsion 'bounce' of the positive alpha particles,
BUT what it was (the 'nucleus') it wasn't very big!
These results made it quite plain the
JJ Thompson plumb pudding model was in some way wrong ie the positive
charge was NOT spread throughout the volume of an atom, therefore a new
model must be proposed to take into account the new results.
From a detailed mathematical
analysis of the scattering experiment results, the only 'atomic model'
which could account for the pattern was an atom consisting of ...
(which is why most alpha particles passed through undeflected),
thus completely contradicting JJ Thompson's 'plum pudding'
model. Other experiments showed that the electrons were orbiting
in energy levels around the nucleus, but occupying virtually no
significant volume in themselves as particles.
a relatively minute
positive centre (the nucleus) causing deflection
(like charges repel,
alpha particles are positively charged and so were being
repelled by the 'later
to be discovered' positive protons in the nucleus), we now
know the nucleus is positive due to protons,
dense centre of similar or greater charge or mass to an
(which we now call the nucleus),
so most of the mass of an atom was in the central nucleus, we
know the mass is made of protons and neutrons.
3. Most of the atom is mainly empty space with a cloud of
negative electrons moving around the dense relatively massive positive nucleus.
Putting these three points together formed the
basis of the modern picture of the 'nuclear atom', in other
words the nuclear atomic model.
there was still a puzzle to solve - why didn't the negative
electrons collapse into the nucleus?
The great physicist Niels Bohr suggested the electrons orbited in
energy levels (shells) that prevented electrons from being attracted
into the nucleus - but would experiments confirm this theory?
Later experiments did show that
electrons are arranged in energy levels, sort orbits around the nucleus,
ideas first proposed by scientists such as
Niels Bohr adapted the nuclear model by suggesting that
electrons orbit the nucleus at specific distances from the nucleus and
each orbit was a specific electron energy
level - a fixed electronic energies.
The theoretical calculations of Bohr agreed with experimental
Later experiments led to the idea that the positive charge
of any nucleus could be subdivided into a whole number of smaller
particles, each particle having the same amount of positive charge.
name proton was given to these particles.
This was a necessary extension and modification to the Rutherford model
of an atom, because this model could not account for why the electrons
were not attracted to the nucleus.
suggested that the negative electrons can only exist in certain
specific energy levels (shells) at fixed distances from the
nucleus and held in place by the positive
nucleus. This theory added too, and complimented the Rutherford model of the atom to gives a
reasonably complete picture of an atom (at least for this
e.g. on the right the 'Bohr'
electronic diagram for sodium.
Bohr envisaged the electrons
orbiting the nucleus in specific energy levels (or fixed shells) at specific distances
from the central nucleus with nothing in between. In other words the
electrons have sufficient energy to keep away from the nucleus and
be confined to these specific energy levels. The negative charge of the
electrons was still balanced by the positive charge (protons) of the nucleus.
So now, as far as we can tell (at GCSE/A level anyway) an atom is quite well represented by the
Bohr model of the atom
(picture below) which moves the Rutherford nuclear model another
of a 'Bohr' lithium atom
AND, most importantly,
experimental results matched a theoretical mathematical model of
simple 'electronic' atoms like hydrogen.
So by now, earlier theories of atomic structure, e.g. the 'plum pudding'
model in which 'protons' and 'electrons' were scattered or arranged
evenly across the atom, were superseded by the nuclear model of
Rutherford and subsequently this was superseded by Bohr's electronic
energy level model.
It was the only
model that could explain the scattering of the high speed
alpha particles by a small dense and positive atomic centre
AND the behaviour of electrons.
Experiments had shown that the outer bits could be knocked off
atoms and these had a very tiny mass and a negative charge,
in other words the
Further experiments showed that the nucleus (partly) consisted of
positive particles with the same mass and charge as an ionised
hydrogen atom, that is a proton
(mass 1, charge +1).
for GCSE level)
In 1913 Moseley
studied the X–rays emitted by highly energised–ionised atoms and from
the X–ray spectra of elements (the K alpha line,
Kα) he was
able to deduce the electric charge of the nucleus which we now know is
equal to the atomic number of protons in the nucleus.
Moseley showed that when atoms were bombarded with cathode rays
(electrons) X–rays where produced which he investigated with an X-ray
It was found that the square root of
the highest energy emission line (called the K alpha line, Kα) gave a
linear plot with the apparent atomic number Z (it
wasn't known yet that this was the proton number),
Z = constant x √Kα
but the plot of √Kα against atomic weight
(relative atomic mass) gave a zig–zag plot, suggesting this 'atomic
number' was far more important the 'atomic weight' of an element in
terms of the atom's fundamental structure.
K alpha line, Kα
is due to an electronic transition of the inner most electron
nearest the nucleus.
(ii) Sadly, Moseley was
killed in action during the First World War at Gallipoli in
1915, a great loss to science as well as his family and friends.
(iii) We now know that
Moseley's 'atomic number' is in fact the number of protons in
the nucleus (atomic number = proton number).
(iv) By 1898, thanks to the
German scientist Wilhelm Wien, the hydrogen ion (proton) had
been identified as the simplest basic unit of positive charge.
By 1925, later experiments by Rutherford and others, identified
the number Z as the value of the positive charge of the nucleus
and that it equated to Z protons in the nucleus - hence the
atomic number = proton number = Moseley's Z value.
However, there was
still the problem of why the atomic mass and atomic number where
different i.e. in the case of the lighter elements, the atomic weight
was often about twice the atomic number.
In 1919 Aston developed a
cathode ray tube i.e. like those used by Wien and Thompson etc. into a
'mass spectrograph', which we now know as a
mass spectrometer GCSE–AS
atomic structure notes.
This showed that atoms of the same
element had different masses but there was no experimental evidence that
they had different atomic numbers (which of course they didn't). These
different atoms of the same element were called isotopes.
Rutherford suggested there might be a 'missing' neutral particle and in
1932 Chadwick discovered the neutron by bombarding beryllium atoms with
alpha particles which produced a beam of neutrons.
These were shown to have a
relative mass of 1 (same as a proton) and were electrically
neutral and quite penetrating into matter. This penetration and
lack of charge had made them difficult to detect.
Prior to this, Rutherford and
others had conducted experiments to show that the smallest
particle in an atom was equivalent to a hydrogen atom without
its electron, that is the proton.
was not until 1932 that the nature of the neutron
was finally deduced by
and he showed that the nucleus also contained an electrically neutral
particle of similar mass to a proto, and this completely explained the nature of isotopes and backed up the
ideas from Moseley's work that the fundamentally important number that
characterises an element is its atomic number and NOT the atomic mass
(or mass number).
The neutron discovery, ~20 years after the discovery of the nucleus,
completed the 'modern' picture and theory of the composition of an atom
in terms of the three principal sub-atomic particles - which is
sufficient for the needs of us chemists!
Advanced level note on the discovery of neutrons:
Chadwick bombarded a thin metal foil of beryllium atoms (94Be)
with alpha particles (42He) and this produced
a highly penetrating radiation that was unaffected by electrical or
42He ====> 126C
This 'neutral' penetrating radiation was eventually identified as a
beam of neutrons!
The neutron wasn't detected directly, the
neutron beam knocks protons out of atoms which could be
detected. Its very difficult to detect a neutral particle
because it doesn't do anything in electric or magnetic fields.
There are no deflection observations to put into a mathematical
model to work out the mass or charge of the deflected particle
e.g. from the sort of experiments JJ Thompson did.
Note the carbon-12 atom produced balances the nuclear equation in
terms of mass (9 + 4 = 12 + 1) and positive proton nuclear charge
(4+ + 2+ = 6+).
page on other experiments with mixed particle beams and
Atomic structure and radioactivity
TOP OF PAGE and
Appendix 2. Atomic structure
diagrams – some variations!
e.g. for the element lithium 73Li
consisting of three protons and four neutrons
TOP OF PAGE and
3. Allotropes – don't
confuse with isotopes!
WHAT ARE ALLOTROPES?
isotopes are atoms
of the same element with different masses due to different numbers of
neutrons in the nucleus. Same protons and electrons.
number 6 = 6 protons = carbon, but there can be 6, 7 or 8 neutrons giving
isotopes of carbon–12, 13 or 14.
As explained above,
Oxygen atoms usually form
oxygen molecules (also called dioxygen), BUT they can form a very
molecule O3 ozone (also called trioxygen). The mass of the
oxygen atoms in each of the molecules is mainly 16 (99.8%), and about 0.2%
of two other stable isotopes of masses 17 and 18. Whatever isotope or isotopes
make up the molecule, it doesn't affect the molecular structure or
the respective chemistry of
the O2 or O3 molecules.
However, what sometimes confuses the issue
is the fact that oxygen O2 and ozone O3 are
examples of allotropes.
defined as different forms of the same element in the same physical state.
The different physical
allotropic forms arise from different arrangements of the atoms and
molecules of the element and in the case of solids, different
They are usually
chemically similar but always physically different in some way e.g.
(oxygen, dioxygen) and O3 (ozone, trioxygen) are both gases
but have different densities, boiling points etc.
diamond and buckminsterfullerene are all solid allotropes of the element
carbon and have significantly different physical and in some ways
chemical properties! (details
on bonding page)
Rhombic and monoclinic sulphur have
different geometrical crystal structures, that is different ways of packing the
sulphur atoms (which are actually both made up of different packing
arrangements of S8 ring molecules). They have different solubilities and melting points.
There is also a 3rd unstable allotrope of sulfur called
made by pouring boiling molten sulphur into cold water which forms a
black plastic material consisting of chains of sulphur atoms
It doesn't matter which isotopes make up the structure of any of an
element's allotropes described above, so to summarise by one example
oxygen–16, 17 or 18 are isotopes
of oxygen with different nuclear structures due to different numbers
but they behave chemically in an identical manner.
and O3 are different molecular structures of the same element
in the same physical state and are called allotropes
irrespective of the isotopes that make up the molecules.
Allotropes can have different physical and chemical properties.
TOP OF PAGE and
Appendix 4. The effect of
ionizing radiation on atoms
Ionisation by EM radiation of atoms to
form a positive ion
The higher energy uv, and both X-ray
and gamma radiations have enough energy to cause complete ionisation of an atom.
Ionisation has already been
explained on this page, but a reminder won't go amiss!
An atom is ionised if it
completely loses one or more electrons to form a positive ion.
The process is illustrated
below for the formation of a monopositive ion of sodium from a
In terms of electrons, the change is
== high energy uv/X-ray/gamma ray photon ==>
This represents the ionisation of a sodium
atom to form a positive sodium ion and a free electron:
Na ==> Na+ + e¯
(electron configuration change of sodium from 2.8.1 ==> 2.8, as in chemistry
In this case the incoming EM radiation must
have sufficient energy to promote the electron all the way up the energy levels
until it is completely free of the attraction of the positive nucleus - so the
atom has been ionised.
A positive ion is formed because there are
now less negative electrons on the atom than positive protons - so there is a
surplus of positive charge on the atom.
The more electrons 'knocked off' the
bigger the positive charge on the ion.
i.e. the charge on the ion can be +, 2+, 3+
etc. by knocking off 1, 2 or 3 electrons etc. The more electrons knocked off,
the bigger the positive charge on the ion.
(Note: Using X-rays, you can knock off all 92
electrons from a uranium atom, element 92, 92U, to form the U92+
ion, but this far too extreme for GCSE students!, but very exciting to
More on ionisation and ionising
radiation in the
Alpha, beta & gamma radiation - properties of 3 types of radioactive
nuclear emission & symbols
,dangers of radioactive emissions - health and safety issues and ionising radiation
gcse physics revision notes
sources, types, properties, uses (including medical) and dangers
TOP OF PAGE and
5. Mass Spectrometer
(now on a separate page for advanced level students!)
See also Atomic structure
other associated web pages
GCSE/IGCSE Foundation Atomic Structure multiple choice QUIZ
Higher Atomic Structure multiple choice QUIZ
gap–fill worksheet on atomic structure
structure and elements"
2nd Word-fill quiz
and definitely NOT GCSE/IGCSE
pages on atomic structure
Advanced Level Chemistry notes on electronic structure – s, p, d orbitals
A Level notes on electron configurations of elements & the periodic table
A Level Notes on mass spectrometers, mass spectrometry and relative atomic
TOP OF PAGE and
9-1 chemistry: Appreciate that new experimental evidence may lead to a
scientific model being changed or replaced.
Before the discovery of the electron atoms were
thought to be tiny spheres that could not be
The discovery of the electron led to the plum
pudding model of the atom. The plum pudding
model suggested that the atom was a ball of
positive charge with negative electrons embedded
The results from the Rutherford and Marsden’s
alpha scattering experiments led to the plum
pudding model being replaced by the nuclear
Niels Bohr adapted the nuclear model by
suggesting that electrons orbit the nucleus at
specific distances. The theoretical calculations of
Bohr agreed with experimental observations.
Later experiments led to the idea that the positive
charge of any nucleus could be subdivided into a
whole number of smaller particles, each particle
having the same amount of positive charge. The
name proton was given to these particles.
In 1932 the experimental work of James
Chadwick provided the evidence to show the
existence of neutrons within the nucleus. This historical context provides an opportunity
for you to show an understanding of why
and describe how scientific methods and
theories develop over time. You should be able to: Be able to describe why the new evidence
from the scattering experiment led to a change in the atomic
model. Details of these experiments are not required. Be able to describe the difference between the plum
pudding model of the atom and the nuclear
model of the atom. Details of experimental work supporting the Bohr
model are not required.
Details of Chadwick’s experimental work are not
required. Know the relative electrical charges of the particles
in atoms: proton, neutron and electron.
In an atom, the number of electrons is equal to
the number of protons in the nucleus. Atoms
have no overall electrical charge.
The number of protons in an atom of an
element is its atomic number. All atoms of a
particular element have the same number of
protons. Atoms of different elements have
different numbers of protons. You should be able to use the atomic
model to describe atoms. Appreciate that atoms are
extremely small, having a radius of about
0.1 nm (1 x 10-10 m).
The radius of a nucleus is less than 1/10 000 of
that of the atom (about 1 x 10-14 m). Almost all of the mass of an
atom is in the nucleus. You must know the relative masses of
protons, neutrons and electrons. The sum of the protons and neutrons
in an atom is its mass number. Atoms of the same element can have
different numbers of neutrons; these atoms are called isotopes of
that element. Atoms can be represented as shown symbolically
(upper left = mass number, lower left = atomic number). You should be able to calculate the
numbers of protons, neutrons and electrons in
an atom or ion, given its atomic number and
Be able to use SI units and the prefix nano. Be able to recognise expressions in standard form.
Be able to estimate the size and scale of atoms. he electrons in an
atom occupy the lowest available energy levels (innermost available
shells). The electronic structure of an atom can be represented by
numbers or by a diagram (on left). For example, the electronic
structure of sodium is 2,8,1 or showing two electrons in the lowest
energy level, eight in the second energy level and one in the third
energy level. You may answer questions in terms of either energy
levels or shells.
You should be able to work out and represent the
electronic structures of the first twenty
elements of the periodic table in both forms.
ocr A gcse 9-1
physics P1.1a Be able to describe how and why the atomic model has changed over time including
Thomson, Rutherford (alongside Geiger and Marsden) and Bohr models. Check out
the timeline showing the development of atomic theory
and discussion of the different roles played in developing the atomic model and how
different scientists worked together.
P1.1b Be able to describe the atom as a positively charged nucleus surrounded by negatively
charged electrons, with the nuclear radius much smaller than that of the atom
and with almost all of the mass in the nucleus. P1.1c Know the typical size (order of magnitude) of atoms and small molecules
typically 1 x 10-10 m
for this page: allotropes * Alpha particle scattering * Atomic (proton)
number * Atom structure * Electron * arrangement (examples) *
Electron shell rules * ions * Isotopes * Mass (nucleon) number * Neutron * Neutron number * Nuclide
symbol notation * Periodic Table (and electron structure) * Proton *
stable/unstable electron arrangements * mass electric charge proton neutron
electron fundamental sub–atomic particles *1. The Structure of Atoms – three
fundamental particles 2. Isotopes – definition and examples 3. The
Electronic Structure of Atoms – rules to be learned 4. Which electron
arrangements are stable and which are not? 5. The Periodic Table and
Electronic Structure – more patterns! 1. to 5. are essential for ALL
GCSE/IGCSE science–chemistry students Appendix 1. The Alpha Particle
Scattering Experiment * 2. Atomic structure diagrams – some variations! 3.
Allotropes – don't confuse with isotopes! 4. The Mass Spectrometer (separate
page for advanced level students!) AND Atomic Structure Crossword Puzzle *
ANSWERS! * multi–word gap–fill worksheet on atomic structure GCSE/IGCSE
Foundation Atomic Structure lower tier multiple choice quiz or
GCSE/IGCSE Higher Atomic Structure higher tier multiple choice quiz
What does a nucleus of an atom consist
of? what are isotopes? A Level GCSE IGCSE revision
notes on atomic structure information on atomic structure KS4 Science atomic
structure GCSE chemistry guide
notes on atomic structure for schools colleges academies science course tutors images
pictures diagrams of apparatus for atomic structure investigations word balanced
symbol equations of atomic structure science chemistry revision notes on
atomic structure revising
the chemistry of atomic structure help in chemical understanding of
atomic structure description
of atomic structure experiments for chemistry courses university courses in chemistry
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chemistry atomic structure USA US grade 8 grade 9 grade10 atomic
structure chemistry explanations of atomic structure chemistry revision
notes atomic nucleus AQA science GCSE chemistry proton particle Edexcel
science GCSE chemistry neutron particle OCR 21st century science GCSE
chemistry masses of sub-atomic particles OCR Gateway science GCSE chemistry
electric charges of sub-atomic particles WJEC gcse science chemistry atomic
models CCEA/CEA gcse science chemistry history of developments in
atomic structure notes for revising models of atoms atomic structure of the
elements gcse chemistry revision
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chemistry what is the structure of an atom? what is the mass
and charge of a proton? what is the mass and charge of a neutron? what is
the mass and charge of an electron? what are isotopes? definition of proton
number, what are isotopes? what are sub-atomic particles? diagram of an
atom, atomic structure diagrams, diagrams of isotopes, give examples of
isotopes, definition of mass number, definition of atomic number, atomic
notation for elements, nuclide notation for isotopes, isotopes of hydrogen,
isotopes of helium, atomic notation for ions, isotopes of bromine, isotopes
of chlorine, isotopes of sodium, diagrams of the isotopes of carbon,
electron arrangements of isotopes, rules on electron arrangements, history
of atomic models, Dalton's theory, experiments and theories of JJ Thompson,
explanation of Rutherford's alpha particle experiment, nuclear atom theory,
electrons in orbits, explaining what we mean by allotropes, allotropes of
oxygen, allotropes of sulfur, what is the relative size of an atom? how
small is an electron? how small is a proton? how small is a neutron? what is
the relative size of the nucleus
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See also Atomic structure