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Brown's Chemistry Clinic
GCSE-KS4-IGCSE-AS
Science-CHEMISTRY revision-information notes on
Atomic Structure, Isotopes,
Periodic Table and Electronic Structure of Atoms
Atomic Structure page CONTENTS
1. The
Structure of Atoms - fundamental particles
2. Isotopes
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!
Appendix 1. The Alpha Particle Scattering Experiment
Appendix 2. Atomic structure diagrams - some
variations!
Appendix 3. Allotropes - don't confuse with isotopes!
Appendix 4 The Mass Spectrometer
(for advanced level students)
See also on other GCSE pages
Chemical
Bonding Notes and diagrams * Radioactivity
and nuclear reaction notes * Periodic Table notes
and atomic structure quizzes GCSE
Foundation m/c quiz or GCSE
Higher m/c quiz and a
multi-word
fill exercise on atomic structure * Advanced
Level Chemistry Students - electron configurations/arrangements and the
Periodic Table
KEYWORDS-phrases for
this page:
allotropes *
Alpha particle scattering *
Atomic (proton) number * Atom structure
* Electron
* Electron arrangement (examples)
* Electron shell rules
* ions * Isotopes
* Mass
(nucleon) number
* Neutron
* Neutron
number
* Nuclide symbol notation
* Periodic Table (and electron structure)
* Periodic Table (its general structure)
* Proton
* stable/unstable
electron arrangements
Email
query?comment
1. The
Structure of Atoms
- fundamental particles
Atoms are the smallest particles of matter whose
properties we study in Chemistry. However from experiments done in the late
19th and early 20th century it was deduced that atoms
are
made up of three fundamental sub-atomic particles, protons, neutrons and
electrons, which are listed below with their
relative masses and electrical charges.

A
Portrait of an Atom
- an image of what you can't see!
- The diagram below gives some idea on the
structure of an atom, it also includes some important definitions and
notation used to describe atomic structure.
- 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
Appendix 1).
- 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.
- 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
particular element element.
- The mass number (A)
is also known as the nucleon number, that
is the number of particles in the nucleus of a particular isotope
(notes on isotopes - definition and examples).
-
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, 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.
- Other more
'practical' diagrams than the one below are shown in
Appendix
2.
- In the example below for lithium-7,
the nuclide notation states that
- before the chemical symbol of
the element
- the top left number =
nucleon/mass number = 7
- and the bottom left number =
proton/atomic number = 3


2. ISOTOPES
- Isotopes are atoms
of the same element with different numbers of neutrons.
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 dentical
chemically.
- There are small physical differences between
the isotopes e.g. the heavier isotope has a greater density or boiling point.
- However, because they have the same
number of protons (proton/atomic number) isotopes of a particular element have the same electronic structure and identical
chemistry.
Examples are illustrated below.
- 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
stable and NOT radioactive i.e. most of the atoms that make up you and the
world around you!
,
and
are the three isotopes of hydrogen with mass numbers of 1, 2 and 3, with 0,
1 and 2 neutrons respectively, but all have 1 proton. Hydrogen-1 is the most
common, there is a trace of hydrogen-2 naturally but hydrogen-3 is very
unstable and is used in atomic fusion weapons.
and
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, it picks up two electrons and becomes the
stable atoms of the gas helium.
and
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.
- The relative atomic mass of an element is the
average mass of all the isotopes present compared to 1/12th of the mass of
carbon-12 atom (12C = 12.00000 i.e. the standard). See
also calculations
page section
- DO NOT CONFUSE ISOTOPES and
ALLOTROPES - see Appendix 3.

3.
The Electronic Structure of Atoms
- rules to be learned
(electron
configuration, electron structure of atoms - arrangement in shells or energy levels)
-
The electrons are arranged in energy levels or shells around the nucleus and with increasing 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 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 (see the Periodic Table diagrams further down).
-
The 1st shell has a maximum of 2 electrons
-
The 2nd shell has a maximum of 8 electrons
-
The 3rd shell has a maximum of 8 electrons
-
The 19th and 20th electrons go into the 4th shell. (required limit of GCSE knowledge)
-
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).
Examples: diagram, symbol or name of element (Atomic Number = number of electrons in a neutral atom), shorthand electron arrangement.
Filling 1st shell, electron level 1
2 elements only,
Period 1 of the Periodic Table
Filling 2nd shell, electron level 2
to to
3 of the 8 elements of Period 2
Filling 3rd shell, electron level 3
to
3
of the 8 elements of Period 3
The first 2 elements of the 4th shell
to Kr [2.8.18.8],
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!)
Advanced level
chemistry - s, p, d and f electron figurations
explained

4.
Which electron arrangements are stable
and which are not?
-
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 [2], 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 - most
of chemistry is about what outer electrons do or don't!
-
[2],[2,8]
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.
-
The most reactive metals
have just one outer electron.
-
These are the Group
1 Alkali Metals, lithium [2,1], sodium [2,8,1],
potassium [2,8,8,1]
-
With one outer shell
electron, they have one more electron than a stable Noble Gas
electron structure.
-
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
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 ==>
sodium ion, Na ==> Na+ is [2.8.1] ==> [2.8] electronically.
-
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
atoms react!
-
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 ==>
chloride ion, Cl ==> Cl- is [2.8.7] ==> [2.8.8]
electronically,
-
the positive
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
the ion.

5. The Periodic Table and
Electronic Structure
- more patterns!
Selected Elements of the Periodic Table are shown below
with
atomic number and chemical symbol.

-
The elements are laid out in order of
Atomic Number.
-
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
Noble Gases),
was known well before the electronic structure of atoms was understood.
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. BUT once
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.
-
For example:
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. Argon, in terms
of its physical, chemical and electronic properties is clearly a Noble Gas in
Group 8 (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 e.g. Group 1 The Alkali Metals (Li, Na, K etc.), Group 7 The Halogens (F, Cl, Br, I etc.) and Group 8 or 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
Periodic Table
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).



APPENDIX
1. The Alpha Particle Scattering Experiment
When alpha
particle beams are fired on very thin layers of metals (e.g. very fine
gold leaf) the results were found to be rather surprising to
scientists of the early 20th century
By using a 360o
charged particle detection system it was found that
-
most
particles passed through un-deflected
-
a small
proportion were deflected slightly
-
and
about 1 in 20,000 were 'bounced' back through an angle of
over 90o
From a detailed
mathematical analysis of the scattering results, the only 'model'
which could account for the pattern was an atom of ...
-
mainly
empty space (why most alpha particles passed through),
-
a
positive centre causing deflection (like charges repel),
alpha particles are positively charged and so are the 'later
to be discovered' protons in the nucleus,
-
a tiny
dense centre of similar or greater charge or mass to an
alpha particle (which we now call the nucleus),
-
in other
words, an atom is well represented by the
diagram
near the top of this page.
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 superceded by the model
described in the previous
picture.
It was the only
model that could explain the scattering of the high speed
alpha particles by a small dense and positive atomic centre.
Later
experiments showed that the out bits could be knocked off
atoms and these had a very tiny mass and a negative charge,
in other words the electron!
See
section 4.
Radioactivity Notes
page on other experiments with mixed particle beams and
their separation.

Appendix 2. Atomic structure
diagrams - some variations!
 
Appendix
3. Allotropes - don't
confuse with isotopes!
-
Isotopes are atoms
of the same element with different masses due to different numbers of
neutrons in the nucleus. Same protons and electrons.
e.g. atomic
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
'stable' O2
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. -
Allotropes are
defined as different forms of the same element in the same physical state.
-
They are usually
chemically similar but always physically different in some way e.g.
-
O2
(oxygen, dioxygen) and O3 (ozone, trioxygen) are both gases
but have different densities, boiling points etc.
-
Graphite,
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 molecules). They have different solubilities and melting points.
It doesn't matter which isotopes make up the structure of any of an
elements allotropes described above,
so to summarise by
example!,
oxygen-16, 17 or 18 are isotopes of oxygen
with different nuclear structures due to different numbers of neutrons,
and O2
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 4 The Mass Spectrometer
(for advanced level students only)
-
The
mass spectrometer is an instrument by which you
can separate
ionised/charged (+) particles of
different mass and determine the amounts of each particle in a
mixture.
-
The technique is called
mass spectroscopy or
mass spectrometry
('mass-spec' and 'MS' in shorthand!).
-
The substance to be analysed is introduced/injected into a
high vacuum
(extremely low pressure) tube system (at K
left diagram) where the particles are
ionised
by colliding with beam of high speed electrons (at
Q
in left diagram).
-
The resulting (+) ions are
accelerated
down a tube (from + to - plates,
P
in left diagram) and then through a powerful magnetic field.
-
The charged or ionised particles are
deflected
by this powerful magnetic field (R
in left diagram).
-
How much they are deflected depends on the particle mass
and the speed of the particle and the strength of a magnetic field i.e.
lighter particles of lower mass (and momentum) are deflected more than
heavier particles of bigger mass (see right diagram below) for a given
set of conditions.
-
By
varying the strength of the magnetic field,
it is possible to bring into focus onto an
ion detector
(N
in left diagram) at the end of the tube (effectively an electrical event
is detected), every possible mass in turn and
a measure the strength of the ion current,
which is
a measure of how much of that ion
has been formed from the sample under analysis.
-
A simplified diagram
of a mass spectrometer tube system is shown below (left) with further explanation as
to what is going on and further diagram to show the relative paths of
light to heavy ions for a given strength of magnetic field.

-
Key to diagram
and more detail of each component's function.
-
K
= sample injection
point, it must be a gas, so a liquid/solid must be vaporised at the
injection point.
-
Q
= high voltage (high +/- p.d.) electron gun which fires a beam of high
speed/energy electrons from a heated 'metal element' into the vaporised sample under analysis and causes
ionization of the atoms (or molecules) forming positive ions
(mainly monopositive in charge).
-
P
= are negative plates which accelerate the positive ions
down the tube (there are positive plates at the start of the tube). A
moving beam of charged particles creates a magnetic field around itself,
and this 'ion beam' magnetic field interacts with the magnetic field at
R.
-
R = the magnetic field that causes
deflection of ions, this is can be varied to change the extent
of deflection for a given mass and to focus a beam of it down onto the detector.
Hence, by programming the mass spectrometer to 'sweep' through all
likely particle masses, in terms of the right hand diagram, you can increase the
strength of the magnetic field to bring into focus onto the ion
detector monopositive ions of increasing mass.
-
N
= an ion detection
system which essentially generates a tiny electrical current when the
ions hit it. The strengths of the 'electronic' signals from the various ion
peaks are sent to a computer for analysis, computation and display. They
tell you the particle masses present and their relative abundance (see
the mass spectrum diagram for the element strontium below).
-
The resulting record of the
ion peaks is called the
mass spectrum
or
mass spectra. The highest peak
is called the base peak
and is often given the relative and arbitrary value of
100, particularly in the mass
spectra of organic compounds).
-
For elements
you get a series of signals or ion peaks for each isotope present and
the ratio of peak heights gives you the relative proportion of each
isotope in the element so that you can calculate the relative atomic
mass of an element. This 'simple' spectra of
mononuclear ions like Sr+
is only true for non-molecular elements like metals (see
mass spectrum of strontium diagram below) or noble gases, but for molecular elements like
nitrogen or the halogens things are not so simple (see chlorine example below).
-
For larger e.g. organic molecules, things can be very
complex indeed, as molecules fragment and many different ions can be formed.
-
Chlorine is a good example of a
molecular element
whose mass spectra can be a bit tricky when first encountered ...
-
Chlorine
consists of two principal isotopes,
chlorine-37 (25% is
37Cl) and
chlorine-35 (75% is 35Cl).
-
BUT, chlorine consists of Cl2
diatomic molecules, which, on ionisation, can split into chlorine atoms.
-
The result is a series of 5
different mass peaks from the various isotopic atomic or molecular ion possibilities...
-
[37Cl37Cl]+
m/z = 74,
[37Cl35Cl]+
m/z = 72,
[35Cl35Cl]+
m/z=70,
[37Cl]+
m/z=37,
[35Cl]+
m/z=35
-
m/z
means the relative mass of the ion over its charge, which for our
purposes the charge is taken +1 (little z) and the mass (little m) is the relative
atomic/formula mass of the particle.
-
Examples of the
calculation of the relative atomic mass of an element using % of
isotopes is given in Part
1 of GCSE-AS (basic) calculations, an example of calculating
relative atomic mass from a mass spectrum is given below for the
metallic element strontium.
-
The mass spectra of organic compounds
can be very complex as the molecules fragment under electron
bombardment, but the resulting
mass spectra can used
to identify compounds
from their 'finger-print' pattern of ion peaks of different mass and particular proportions
for a given set of experimental conditions.
-
The largest m/z value gives the
molecular mass of a molecule,
i.e. the ion of largest mass, prior to fragmentation, is formed when the
original whole and neutral molecule, loses one electron e.g. for ethane
it would be due to the formation of [C2H6]+,
m/z = 30 and is called the molecular ion peak.
-
Example of a
relative atomic mass calculation based on the mass spectrum of the
element

-
The relative atomic mass of an
element, Ar, is the weighted average mass of the isotopes
present, compared to 1/12th of
the relative mass of the carbon-12 isotope. [ 12C is given
the relative mass value of 12.0000 ]
-
Quite often the highest peak is
arbitrarily given the relative value of 100, as in this case, but the peak lines
might well indicate % abundance of isotopes.
relative
peak height = relative abundance
as measured from the ion current detector signal.
-
The mass spectrum
shows strontium consists of four isotopes, 84Sr (peak height
= 0.68), 86Sr (peak height = 12.0),87Sr (peak
height = 8.47) and 88Sr (peak height = 100.0)
-
The sum of the
heights = 0.68 + 12.0 + 8.47 + 100.0 = 121.15
-
So we can now calculate
the weighted average mass of ALL the isotopes.
-
Therefore Ar
= {(0.68 x 84) + (12.0 x 86) + (8.47 x 87) + (100.0 x 88)}/121.15 =
87.7
-
The book value is
87.62, BUT this calculation does NOT take into account the very accurate
relative atomic masses based on the carbon-12 scale, it merely uses the
mass numbers, which are always integer. The very accurate isotopic
masses are usually a tiny fraction different from a whole number.
-
Modern mass
spectrometers are exceedingly accurate and very sophisticated
instruments and can measure mass to at least 4 decimal places. They can
readily
distinguish between N2 and CO molecules, both with an Mr
of 28.

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guidance, private tuition, school retakes revision. Whether you are a
teacher/tutor teaching, a student studying, using the pages as self-study guides
for your science-chemistry studies etc. etc. I hope the site supports your endeavour.
15-12-07 © Dr W P Brown
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