Chemistry KS4 science GCSE/IGCSE/AS O Level Revision Notes
Atomic Structure, Isotopes,
Electronic Structure of Atoms
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. The simplified
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.
The link between the Periodic Table and Electronic Structure is described.
The important historic alpha particle scattering experiment is described,
variations of atomic structure diagrams, Allotropes – 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
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
1. to 5. are essential for
ALL KS4 Science GCSE/IGCSE science–chemistry students
Alpha particle scattering *
Atomic (proton) number * Atom structure
* 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)
Appendix 1. The Alpha Particle Scattering Experiment
Appendix 2. Atomic structure diagrams
Appendix 3. Allotropes
– don't confuse with isotopes!
The Mass Spectrometer
(separate page for advanced level students!)
GCSE/IGCSE Foundation Atomic Structure multiple choice QUIZ
Higher Atomic Structure multiple choice QUIZ
gap–fill worksheet on atomic structure
Structure of Atoms
– three fundamental particles
ARE ATOMS? and WHAT DO WE MEAN BY FUNDAMENTAL PARTICLES? (sub–atomic particles)
Compounds are formed when two or more
elements are chemically combined to form a new substance in a reaction
which is not easily reversed ie its difficult to separate a compound back
into its constituent elements.
However, what are elements made up of?
How and why do elements bond together?
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) and two or more different atoms/elements must be
present in a compound.
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.
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.
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
You can think of the mass of
an electron as about 1/2000th of the mass of a
proton or neutron
What can we say about 'A
Portrait of an Atom'?
– an image of what you can't see!
However this diagram, although
more realistic in terms of the real size of the nucleus compared to the atom
as a whole, it is not convenient to give a brief diagrammatic picture of the
composition of an atom like the style of diagram 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
- 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 =
–1, NOT in 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 where the
electrons are, the diameter of the nucleus of protons plus neutrons is about a
ten thousandth of the diameter of an atom!
- Since the nucleus is composed of positive
protons and neutral neutrons, the nucleus itself must be positive.
- 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.
- 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'.
- If electrons are removed from an atom you
get a positive ion and if electrons are added to an atom you get a
- An ion, by definition, cannot be neutral.
Some important evidence for this 'picture' is
obtained from alpha particle scattering experiments (see
- The atomic/proton number (Z) is the number of
protons in the nucleus and is also known as the
proton number of the particular element.
- Each element has its own atomic number, so
all the atoms of a particular element have the same atomic 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
- The mass number (A)
is also known as the nucleon number, that
is the number of particles in the nucleus of a particular atom–isotope
(notes on isotopes – definition and examples).
- Therefore the mass/nucleon number = sum of
the protons plus neutrons in the nucleus.
The neutron number (N) =
mass number – proton/atomic number
In a neutral 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.
- If not, and the atom has an overall
surplus electrical charge, and is then called an ion e.g.
- the positive sodium ion Na+ (11 protons, 10 electrons,
excess positive protons),
- or the negative chloride ion Cl–
(17 protons, 18 electrons, excess negative electrons,
- for more details and examples see
ionic bonding notes.
alternative atomic structure diagrams than the one below are described and
shown in Appendix
- In the example below for lithium–7,
the nuclide notation states that
- before the chemical symbol of
the element Li
- the top left number =
nucleon/mass number = 7
- and the bottom left number =
proton/atomic number = 3
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
(iron–56), 26 protons, 30 neutrons (56 – 26), 26 electrons
atom (cobalt–59), 27 protons, 32 neutrons (59 – 27), 27 electrons
atom (californium–98), 98 protons, 148 neutrons (246 – 98), 98
WHAT ARE ISOTOPES? ARE THEY IMPORTANT?
- Isotopes are atoms
of the same element with different numbers of neutrons
and therefore different masses
(different nucleon/mass numbers).
This gives each isotope of a particular element a different mass or nucleon
number, but, being the same element they have the same atomic or proton
number and are identical
- The phrase 'heavier' or
'lighter' isotope means 'bigger' or 'smaller' mass number.
- There are small physical differences between
the isotopes e.g. the heavier isotope has a greater density or boiling point,
the lighter the isotope the faster it diffuses.
- However, because they have the same
number of protons (proton/atomic number) isotopes of a particular element have the same electronic structure and identical
- Examples of isotopes are illustrated
and described below.
- Caution Note: Do NOT assume the word isotope means
the atom it is
radioactive, this depends on the stability of the nucleus i.e. unstable
atoms (radioactive) might be referred to as radioisotopes.
Many isotopes are extremely
stable in the nuclear sense and NOT radioactive i.e. most of the atoms that make up you and the
world around you!
are the three isotopes of hydrogen with mass numbers of 1, 2 and 3, with 0,
1 and 2 neutrons respectively. All have 1 proton, since all are hydrogen! Hydrogen–1 is the most
common, there is a trace of hydrogen–2 (sometimes called deuterium) naturally but
hydrogen–3 (sometime called tritium) is very
unstable and is used in atomic bombs – nuclear fusion weapons.
- They are sometimes denoted more simply as
1H, 2H and 3H since the
chemical symbol H means hydrogen and therefore must have only one proton.
or 3He and 4He,
are the two isotopes of helium with mass numbers of 3 and 4, with 1 and 2
neutrons respectively but both have 2 protons. Helium–3 is formed in the Sun
by the initial nuclear fusion process. Helium–4 is also formed in the Sun
and as a product of radioactive alpha decay of an unstable nucleus.
- An alpha
particle is a helium nucleus (mass 4, charge +2) and if it picks up two electrons it becomes
stable atoms of the gas helium. For more details see
Radioactivity Revision Notes Part 4
or 23Na and 24Na,
are the two isotopes of sodium with mass numbers of 23 and 24, with 12 and
13 neutrons respectively but both have 11 protons. Sodium–23 is quite stable
e.g. in common salt (NaCl, sodium chloride) but sodium–24 is a radio–isotope
and is a gamma emitter used in medicine as a radioactive tracer e.g. to
examine organs and the blood system.
are the two
nuclear symbols for the two most common and stable isotopes of the element chlorine.
They both have 17 protons in the nucleus and 35–17 = 18 and 37–17 = 20
are the two
nuclide symbols for the two most common and stable isotopes of the element
They both have 35 protons in the nucleus and 79–35 = 44 and 81–35 = 46
neutrons respectively. By coincidence, there are almost exactly 50% of each
isotope present in naturally occurring bromine.
- The isotopes of carbon
trace, unstable radioactive
- Now is an appropriate point to introduce
the concept and definition of relative atomic mass (Ar),
which is required for very accurate quantitative chemistry
- The relative atomic mass of an element is the
average mass of all the isotopes present compared to 1/12th of the mass of
a carbon–12 atom (12C = 12.00000 i.e. the standard).
- When you average the masses of the
isotopes of carbon, taking into account their relative abundance (%),
you arrive at a relative atomic mass of carbon of 12.011, Ar(C)
= 12.011, though at this academic level 12.0 is accurate enough!
- Knowledge of isotopes is important in modern
- Radioactive isotopes are used in medicine to trace aspects of body
chemistry due to their radioactive emissions, and in chemical synthesis as
tracers to follow how a reaction sequence occurs.
- Radioactive isotopes are
used in radiotherapy to kill malignant cancer cells.
- DO NOT CONFUSE ISOTOPES and
ALLOTROPES – see Appendix 3.
The Electronic Structure of Atoms
– rules to be learned
WHAT DO WE MEAN BY Electron
configuration, electronic structure of atoms – arrangement in shells or energy levels?
The electrons are arranged in energy levels or shells around the nucleus and with
'orbits' on average increasing in distance from the nucleus.
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, upto 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 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 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 are
summarised below with reference to the periods of the periodic table and in
order of increasing atomic number.
Period 1 – elements 1 to 2
Period 2 – elements 3 to 10
Period 3 – elements 11 to 18
Period 4 – first two elements
19 to 20
written out as 220.127.116.11 and
18.104.22.168 (1st,2nd,3rd full shells with 2,8,8
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!
For the rest of Period 4 and
other Periods you need a more
electron configuration system upto at least Z=36 using s, p, d and f orbital
notation BUT for advanced 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 Periodic Table
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 [22.214.171.124],
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. = 126.96.36.199 etc. upto Z = 20
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
EXTRA NOTE ON 'ATOMIC' NOTATION –
representation of isotopes of ions
Nuclide notation and ions
(interpretation required for advanced level students only)
isotope ion, 11 protons, 13 neutrons, 10 electrons (one electron
lost to form a positive ion)
isotope ion, 11, protons, 12 neutrons, 10 electrons (one electron
lost to form a positive ion)
sulfur–32 in the form of the sulfide ion, 16 protons, 16 neutrons,
18 electrons (two electrons gained to form the double charged
For more on electron
structure and chemical changes and compound formation see ...
and for more on metal and
non–metal reactivity see
5. 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 ' relative atomic mass'
(the old term was 'atomic
weight'). 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.
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 Group 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
1. The history of the atom concept and the Alpha Particle Scattering Experiment
(some of the ideas described
here go above GCSE/IGCSE level!)
Leucippus and Democritus ~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'.
The Greeks idea was not
forgotten and later revived by Boyle and Newton but with little
progress. However in 1808
that all matter was made up of tiny hard particles called atoms and the
different types of atoms (elements) combined together to give all the
different substances of the physical world. 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
known, and, as it happens, correctly so.
~1897 proposed his 'plum pudding' theory based on the growing evidence
that atoms where themselves composed of even small more fundamental
particles and the mass and charge of the proton and electron. Thompson
envisaged a plumb pudding atom consisting of a positively charged
'pudding' with just enough lighter negatively charged electrons embedded
in it to produce a neutral atom. The positive balancing the negative was
correct but the relative size and nature of the nucleus were not.
Rutherford, Geiger and Marsden
(the latter two were students of Rutherford at Cambridge University) conducted alpha particle scattering experiments (1902–1910,
and described in
detail below). 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.
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. 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.
However the plot of √Kα against atomic weight
(relative atomic mass) gave a zig–zag plot.
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). In 1920
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.
It was not until 1932 that the nature of the neutron was finally deduced
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.
Rutherford and Geiger–Marsden scattering experiments
particle beams are fired on very thin layers of metals (e.g. very fine
gold leaf) some rather surprising results were by
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!)
proportion were deflected slightly
about 1 in 20,000 were 'bounced' back through an angle of
(so whatever was there was substantial BUT not very
From a detailed
mathematical analysis of the scattering results, the only 'model'
which could account for the pattern was an atom of ...
(why most alpha particles passed through!),
positive centre causing deflection
(like charges repel,
alpha particles are positively charged and so are being
repelled by the 'later
to be discovered protons'),
dense centre of similar or greater charge or mass to an
(which we now call the
words, an atom is well represented by the
Bohr model of the atom diagram
near the top of this page were we started!
theories of atomic structure, e.g. the 'plum pudding' model in
which 'protons' and 'electrons' were scattered or arranged
evenly across the atom, were superceded by this 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.
experiments showed that the outer bits could be knocked off
atoms and these had a very tiny mass and a negative charge,
in other words the
page on other experiments with mixed particle beams and
Appendix 2. Atomic structure
diagrams – some variations!
e.g. for the element lithium 73Li
consisting of three protons and four neutrons
3. Allotropes – don't
confuse with isotopes!
WHAT ARE ALLOTROPES?
As explained above,
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.
Oxygen atoms usually form
oxygen molecules (also called dioxygen), BUT they can form an unstable
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 of neutrons,
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.
Appendix 4a Introduction to the Mass Spectrometer (for advanced level students only!)
Appendix 4b The Time of Flight Mass Spectrometer (for advanced level students only!)
These two sections are
now on a separate Advanced Level Chemistry page
for this page: allotropes * Alpha particle scattering * Atomic (proton)
number * Atom structure * Electron * Electron arrangement (examples) *
Electron shell rules * ions * Isotopes * Mass (nucleon) number * Mass
spectrometer (advanced students only) * 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
Revision KS4 Science
GCSE/IGCSE/O level Chemistry Information Study Notes for revising for AQA
GCSE Science, Edexcel GCSE Science/IGCSE Chemistry & OCR 21stC Science, OCR
Gateway Science WJEC gcse science chemistry CCEA/CEA gcse science
chemistry (revise courses equal to US grade 8, grade 9 grade 10) Exam
Revision Tuition A Level Revision Guides for A Level Courses Examinations
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 pre–university students (equal to US grade 11
and grade 12 and AP Honours/honors level courses) for revising science
chemistry courses revision guides
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