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Atomic Structure, Isotopes, Electronic Structure of Atoms

Doc Brown's Chemistry KS4 science GCSE/IGCSE/AS/O Level  Revision Notes

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 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 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 notes. 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.

Index for Atomic Structure page

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 KS4 Science GCSE/IGCSE/A Level science–chemistry students

Relative Atomic Mass is dealt with on a separate calculation page

KEYWORDS–phrases for this page: allotropes * Alpha particle scattering * Atomic (proton) number  * Atom structureElectronElectron arrangement (examples)Electron shell rules * ionsIsotopesMass (nucleon) number * Mass spectrometer (advanced students only)NeutronNeutron numberNuclide symbol notationPeriodic Table (and electron structure)Periodic Table (its general structure)Proton * stable/unstable electron arrangements

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 (separate page for advanced level students!)

1. The Structure of Atoms – three fundamental 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 are 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 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

Sub–atomic particle Relative mass Electric charge Comments
Proton 1 +1 (+ positive) In the nucleus, a nucleon
Neutron 1 0 (zero) In the nucleus, a nucleon
Electron 1/1850 or 0.00055 –1 (– negative) 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, so, a pretty small mass BUT they occupy most of the space of atom!!!


What can we say about 'A Portrait of an Atom'? – an image of what you can't see!

(c) doc b

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, it is not convenient to give a brief diagrammatic picture of the composition of an atom like the style of the diagram below. The central nucleus of protons and neutrons (most of mass) 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. Bohr theorised the negative electrons can only exist in certain specific energy levels (shells) held in place by the positive nucleus. 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 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.


(c) doc b

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,
      • 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 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.
      • A neutral atom carries no overall charge because the number of positive protons equals the number of negative electrons, and this information is given by the atomic/proton number.
  • 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 negative ion.
      • An ion, by definition, cannot be neutral.
  • Some important evidence for this 'picture' is obtained from alpha particle scattering experiments (see Appendix 1).
  • The atomic 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 particular element
  • 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).
  • 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 or deficiency electrical charge, and is the resulting electrically charged particle 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.
  • In the example above 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
    • Similarly for ...
    • , atom of hydrogen–1, symbol H, mass 1, just 1 proton and NO neutrons (only atom with no neutrons)
    • , atom of helium–4, symbol He, mass 42 protons, 4 – 2 = 2 neutrons
    • , atom of sodium, symbol Na, mass 23, 11 protons, 23 – 11 = 12 neutrons
  • The electrons are arranged in specific energy levels according to a set of rules (dealt with in section 3).

  • 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 atomic theory.

  • 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 = nucleons)

    • 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 atom (isotope iron–56), mass 56, 26 protons, 30 neutrons (56 – 26), 26 electrons

      • Cobalt atom (isotope cobalt–59), mass 59, 27 protons, 32 neutrons (59 – 27), 27 electrons

      • Californium atom (isotope californium–246), mass 246, 98 protons, 148 neutrons (246 – 98), 98 electrons

      • So, at this point we had better explain, slightly belatedly, what isotopes are!

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  • 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.
    • They are also chemically identical, because they have the same number of electrons, hence the same electron structure.
    • Study the diagrams of the isotopes of carbon further down the page.
    • Relative Isotopic Mass is dealt with on a separate calculation page

  • The phrase 'heavier' or 'lighter' isotope means 'bigger' or 'smaller' mass number for a particular element.
  • 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 chemistry.
  • 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!
  • (c) doc bhydrogen–1, (c) doc bhydrogen–2, and (c) doc bhydrogen–3 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.
  • (c) doc b and (c) doc b 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 a stable atoms of the gas helium. For more details see Radioactivity Revision Notes Part 4
  • (c) doc b and (c) doc b 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 in the nucleus and 11 surrounding electrons. 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.
  • and 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 neutrons respectively (and both have 17 surrounding electrons).
  • and are the two nuclide symbols for the two most common and stable isotopes of the element bromine. They both have 35 protons in the nucleus and 79–35 = 44 neutrons and 81–35 = 46 neutrons respectively. By coincidence, there are almost exactly 50% of each isotope present in naturally occurring bromine.
  • The three known isotopes of carbon
    • isotope nuclide symbol protons neutrons electrons % abundance
      carbon–12 126C 6 6 6 98.9%, stable
      carbon–13 136C 6 7 6 1.1%, stable
      carbon–14 146C 6 8 6 trace, unstable radioactive
    • The table of information on the three isotopes of carbon is illustrated by the diagrams above it.
    • Now is an appropriate point to introduce the concept and definition of relative atomic mass (Ar), which is required for very accurate quantitative chemistry calculations.
    • 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 amu 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!
      • (c) doc b See also chemical calculations on how to calculate relative atomic mass
      • I've put this calculation on its own page because there is plenty on atomic structure already on this page!
      • Anything on this page relevant to the calculation of RAM is repeated on the page.
  • Knowledge of isotopes is important in modern science.
    • 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.

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3. The Electronic Structure of Atomsrules 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.

    • 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 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 (electrons 3–10)

    • The 3rd shell also has a maximum of 8 electrons (electrons 11–18)

    • 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 (2 elements)

      • the electron arrangement is written out simply as 1 or 2

    • Period 2 – elements 3 to 10 (8 elements)

      • electron arrangements of 2.1 to 2.8 (since 1st shell is full with 2 electrons i.e. the first number)

    • Period 3 – elements 11 to 18 (8 elements)

      • denoted by 2.8.1 to 2.8.8 (1st,2nd full shells with 2,8 electrons)

    • Period 4 – first two elements 19 to 20

      • written out as and (1st,2nd,3rd full shells with 2,8,8 electrons)

      • 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 diagram below

    • (c) doc b

    • 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.

      • Some examples 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 periodic table.

      • 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 advanced electron configuration system upto at least Z=36 using s, p, d and f orbital notation BUT for advanced level chemistry students only!

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 (c) doc b (c) doc b2 elements only, Period 1 of the Periodic Table

Filling 2nd shell, electron level 2 (c) doc b to (c) doc b to (c) doc b 3 of the 8 elements of Period 2

Filling 3rd shell, electron level 3 (c) doc b to (c) doc b (c) doc b  3 of the 8 elements of Period 3

The first 2 elements of the 4th shell (c) doc b (c) doc b to Kr [], 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. = etc. upto Z = 20

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4. Which electron arrangements are stable and which are not?

Both atoms and ions are considered




  • 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.

    • For all elements most of their 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.

    • 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], 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 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 ==> sodium ion

      • 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 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

      • in fundamental particles [17p + 17e] ==> [17p + 18e]

      • 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.

  • EXTRA NOTE ON 'ATOMIC' NOTATION – representation of isotopes of ions

  • Nuclide notation and ions (interpretation required for advanced level students only)

    • sodium–24 isotope ion, 11 protons, 13 neutrons, 10 electrons (one electron lost to form a positive ion)

    • sodium–23 isotope ion, 11, protons, 12 neutrons, 10 electrons (one electron lost to form a positive ion)

    • isotope sulfur–32 in the form of the sulfide ion, 16 protons, 16 neutrons, 18 electrons (two electrons gained to form the double charged negative ion)

  • For more on electron structure and chemical changes and compound formation see ...

  • and for more on metal and non–metal reactivity see

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5. The Periodic Table and Electronic Structure – more patterns!

Selected Elements of the Periodic Table are shown below with atomic number and chemical symbol.


(c) doc b

  • 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 Noble Gases), 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.

    • 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. 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 e.g.

    • 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 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).

  • NOTE: In the most modern periodic table notation Groups 3–7 and 0 are numbered Groups 3 to 18.

(c) doc b

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APPENDIX 1. The history of the atom concept and the Alpha Particle Scattering Experiment

(some of the ideas described here go above GCSE/IGCSE level!)

The Greeks 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 Dalton proposed his atomic theory that all matter was made up of tiny hard particles/spheres called atoms and the 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 indivisible spheres'!). 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.

J J Thompson Around 1897 proposed his 'plum pudding' theory based on the growing evidence that atoms were themselves composed of even small more fundamental particles and the mass and charge of the proton and electron i.e. atoms were not hard indivisible spheres. His experiments had shown that atoms contained small negatively charged particle called electrons. From 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 idea of positive particles balancing the negative particles was correct but the relative size and nature of the nucleus were not.

Ernest Rutherford, 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). 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.

(c) doc bThe Rutherford and Geiger–Marsden scattering experiments (1902–1910).

When alpha 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 ...

3. most particles passed through un–deflected (as if there was nothing there!)

2. a small proportion were deflected slightly (so there was something there!)

1. about 1 in 20,000 were 'bounced' back through an angle of over 90o, in other words were reflected backwards, a totally unexpected result. So, whatever was there was substantial, positively charged to cause the repulsion 'bounce', BUT not very big!

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 (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),

... a tiny dense centre of similar or greater charge or mass to an alpha particle (which we now call the nucleus), and this is the modern picture of the 'nuclear atom'.

So an atom is quite well represented by the Bohr model of the atom diagram near the top of this page were we started!  Bohr's suggestion that the negative electrons can only exist in certain specific energy levels (shells) held in place by the positive nucleus complimented the Rutherford model of the atom to gives a reasonably complete picture of an atom (at least for this academic level!).

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 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.

Later 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 electron!

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. 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 by Chadwick 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.

See section 2. Radioactivity Notes page on other experiments with mixed particle beams and their separation.

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Appendix 2. Atomic structure diagrams – some variations!

e.g. for the element lithium 73Li consisting of three protons and four neutrons

(c) doc b

Appendix 3. Allotropes – don't confuse with isotopes!


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. 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.

The different physical allotropic forms arise from different arrangements of the atoms and molecules of the element and in the case of solids, different crystalline allotropes.

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 ring molecules). They have different solubilities and melting points. There is also a 3rd unstable allotrope of sulfur called plastic sulphur made by pouring boiling molten sulphur into cold water which forms a black plastic material consisting of chains of sulphur atoms –S–S–S–S–S– etc..

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 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.

other associated web pages

GCSE/IGCSE Foundation Atomic Structure multiple choice QUIZ

GCSE/IGCSE Higher Atomic Structure multiple choice QUIZ

(c) doc b GCSE/IGCSE Atomic Structure Crossword Puzzle * ANSWERS!

(c) doc b GCSE/IGCSE multi–word gap–fill worksheet on atomic structure

and definitely NOT GCSE/IGCSE pages on atomic structure

Advanced Level Chemistry notes on electronic structure – s, p, d orbitals etc.

A Level notes on electron configurations of elements & the periodic table

A Level Notes on mass spectrometers, mass spectrometry and relative atomic mass

top indexKEYWORDS–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 * 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

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