Why can't we trust atoms? Because they make up everything!

ATOMIC STRUCTURE

Atomic Structure, Isotopes, Electronic Structure of Atoms, Atomic Structure Experiments, ionisation - formation of simple ions (and a brief mention of allotropes in appendix 3.)

1a. Structure of Atoms – three fundamental particles

1b. 'Portrait of an atom' and particle size comparisons

1c. Atoms, sub-atomic particles, nuclide notation & diagrams

2. Isotopes – definition, nuclide notation and examples

3. Electronic Structure of Atoms – rules and the Periodic Table

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

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

Relative Atomic Mass, definition and calculation details are on a separate page

Appendix 1. History of atomic structure models, alpha particle scattering experiment

Appendix 2. Atomic structure diagrams – some variations!

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

Appendix 4. The effect of ionizing radiation on atoms - formation of positive ions

Introduction to the mass spectrometer (only for advanced A level students!)

KEYWORDS 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)

Periodic Table (its general structure) * Proton * stable/unstable electron arrangements

See also Atomic structure and radioactivity

WHAT ARE ATOMS? and WHAT DO WE MEAN BY FUNDAMENTAL PARTICLES? (sub–atomic particles)

An ATOM is the smallest particle of a substance, an element, which can have its own characteristic properties AND cannot be split into simpler substances.

However, why do we have different elements?

Is an atom the simplest particle we need to know about to understand chemistry?

In order to answer these questions we must look a bit deeper into the fundamental structure of matter, that is everything around you!

Atoms are the smallest particles of matter whose properties we study in Chemistry.

Every element or compound is comprised of atoms. All the atoms are the same in the structure of an element (ignoring isotopes - different numbers of neutrons, see later) and two or more different atoms/elements must be present in a compound.

Each element has its own chemical symbol (carbon C, oxygen O, sodium Na etc.), which with added numbers (e.g. right), can be used to indicate the composition of an atom in terms of protons, electrons and neutrons. All of this will be explained in detail below

Initially, once the concept of an atom was established, it was assumed that atoms were indestructible and not divisible into smaller particles, but merely combined in different proportions to give the range of compounds we know about e.g. Dalton's atom model.

However from experiments done in the late 19^th and early 20^th 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.

In Appendix 1. the history of the development of the atomic model is described in detail.

The table gives the relative mass and electric charge of the three sub–atomic particles known as the proton, neutron and electron

Protons and neutrons are much heavier than electrons.

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 an atom!!!

You should also realise because of the relatively small mass of the electrons most of an atom's mass is in the nucleus.

You see values of 1/1836 quoted for the relative mass of an electron, but don't worry about it, there are different ways/scales on which an electron's mass has been calculated.

The actual mass of a proton or neutrons is ~1.67 x 10^-27 kg (~1.67 x 10^-24 g)

The mass of an electron is ~9.1 x 10^-31 kg (~9.1 x 10^-28 g)

The mass of an atom varies from about 1 x 10^-20 to 1 x 10^-18 kg (1 x 10^-23 to 1 x 10^-21 g) depending on the element

The radius of the nucleus ranges from about 1 x 10^-16 to 1 x 10^-14 m (1 x 10^-7 to 1 x 10^-5 nm) depending on the element

The diameter of atoms varies from about 1 x 10^-10 to 5 x 10^-10 m (0.1 to 0.5 nm) depending on the element

Generally speaking the radius of an atom is about 10,000 times that of the nucleus!

A typical relatively small molecule would be no bigger than ~1 x 10^-10 to 1 x 10^-9 m (~0.1 to 1 nm)

1b. What can we say about 'A Portrait of an Atom'?

Images of what you normally can't see! Size comparison of atoms, molecules and cells!

The picture of 'atomic structure, illustrated below is the result of many developing 'atomic theory' backed up by successive generations of experimental results. This is the best picture we have (at least for GCSE and A Level chemistry courses!).

However this diagram, which is based on the Bohr model of atomic structure, although more realistic in terms of the real size of the nucleus compared to the atom as a whole, so it is not convenient to give a brief diagrammatic picture of the composition of an atom.

The central nucleus of protons and neutrons (most of the mass of an atom) is extremely small even compared to the size of an atom. The rest of the 'almost empty space' of an atom is occupied by the negative electrons, held by, and moving around the positive nucleus in their energy levels or 'shells'.

The electrons are also pretty tiny in mass too, compared to a proton or neutron, but the volume the electrons and their energy levels occupy, determines the size of the atom, but an atom mainly empty space with the nucleus at the centre!

Bohr theorised the negative electrons can only exist in certain specific energy levels (shells) held in place by the positive nucleus (see section on the history of development of the atomic model). These are shown in the above diagram, but fully explained later on this page.

All of these theories must, and have been, backed up by repeated and varied experiments.

As each new experiment was/is done, it must support the current theory or the theory needs to be modified to take into account new discoveries.

Some of these important experiments are described further down the page.

Even new experimental findings written up in research papers should be thoroughly peer reviewed, that is checked by scientists of at least equal academic ranking to the researchers. That's how science works!

The size of an atom compared to other 'particles'

The size of an individual atom is around 0.1 nm or 1 x 10^-10 m

and the size of the nucleus itself is only about 1 x 10^-14 m, about 1/10000th of the radius of the atom.

(nm = nanometres, 1 nm = 10^-9 m).

Most of the mass is in the centre of the atom, that is the nucleus, which has a radius of around ¹/₁₀₀₀₀th of the whole atom!

This means the radius of the nucleus is about 1 x 10^-14 m (0.00001 nm), pretty small !!, but still consists of most of the mass of an atom!!!

If you look at the table of size comparison below, its not until you get to a human hair can we see clearly something with the naked eye. The width of a human hair is approximately 10⁶ times that of an atom (a million times bigger) and 10¹⁰ times bigger than a nucleus (ten thousand million times bigger). You can of course see cells under examination with an optical microscope, but these are over 500 000 times bigger than an individual atom. You can, however, observe atoms using an electron microscope.

na means not applicable

TOP OF PAGE and sub-index

1c. Atoms, sub-atomic particles, nuclide notation and diagrams

The number of protons in the nucleus of an atom decides what element that atom is.

e.g. if the atom has 3 protons in the nucleus, it cannot be anything except lithium!

Elements consist of one type of atom only determined by the proton number (atomic number).

Some more concise and handy styles to show the atomic composition of the same lithium atom

• What sub–atomic particles make up atoms? What is their mass and charge?
• The diagram above of a 'portrait of an atom' gives some idea on the structure of an atom (sometimes called the Bohr Atomic Model), it also includes some important definitions and notation used to describe atomic structure
  • REMINDER of the three fundamental particles you need to know are ...
  • proton: particle mass = 1, electric charge = +1, the charged particle in the nucleus
  • neutron: particle mass = 1, charge = 0, uncharged particle in the nucleus
  • electron: particle mass = 1/1850 ~1/2000, electric charge = –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 empty space, BUT it is where the tiny electrons are.
    • In fact the diameter of the nucleus of protons plus neutrons is about a ten thousandth of the diameter of the whole atom!
    • Since the nucleus is composed of positive protons and neutral neutrons, the nucleus itself must be positive.
    • A neutral atom carries no overall charge because the number of positive protons equals the number of negative electrons (both singly charged + and -), and this information is given by the atomic/proton number.
      • A neutral helium atom has 2 protons and 2 electrons and a uranium atom has 92 protons and 92 electrons. 2+ balances 2- and 92+ balances 92-, easy!
      • Since a neutron is electrically neutral, the number of neutrons in the nucleus of an atom cannot affect the total positive nuclear charge of the protons or the number of negative electrons in the atom.
• 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).
  • Although the nucleus must be positive because of the positive protons (neutrons are neutral) an individual atom is neutral because the number of electrons equals the number of protons – so the charges 'cancel out'.
  • An ion particle carries an overall surplus electric charge - positive or negative - so how are they formed?
  • If electrons are removed from an atom you get a positive ion from excess positive charge, and, if electrons are added to an atom, you get a negative ion from excess negative charge.
    • An ion, by definition, cannot be neutral and the number of protons cannot equal the number of electrons.
• The atomic number (Z) is the number of protons in the nucleus and is also known as the proton number of the particular element and it is this number that decides what element a particular atom is.
  • Each element has its own atomic number, so all the atoms of a particular element have the same atomic number.
  • In a neutral atom, the number of electrons equals the atomic/proton number.
• It is the proton/atomic number (Z) that determines the number of electrons an element has, its specific electron structure and therefore the specific identity of a particular element in terms of its physical and chemical properties.
• It cannot be overemphasised that it is the electronic structure that determines the chemical character of an element, hence the proton/atomic number determines everything about a particular element.
• The mass number (A) is also known as the nucleon number, is the number of particles in the nucleus of a particular atom–isotope (notes on isotopes – definition and examples).
  • Therefore the mass number or nucleon number = sum of the protons plus neutrons in the nucleus.
  • Since the mass of a proton or neutron equals 1, the mass number equals the mass of a particular atom to the nearest whole number.

  • Relative Atomic Mass is dealt with on a separate calculation page

• The neutron number (N) = mass number – proton/atomic number
• In an individual atom the number of protons (+) equals the number of electrons (–), that is the number of positive charges is equal to the number of negative charges to make the atom neutral.
• In the example in the diagram for lithium–7, ⁷₃Li is its nuclide atomic notation ...
  • 
  • before the chemical symbol of the element Li
    • is the top left number  = nucleon number/mass number = 7 (3 protons + 4 neutrons).
    • and the bottom left number = proton number/atomic number = 3 (3 protons in nucleus).
    • therefore the number of neutrons = nucleon/mass number - proton/atomic number = 7 - 3 = 4,
    • AND it must have 3 electrons, because number of protons (+) = number of electrons (-) in a neutral atom.
  • Similarly how to interpret the following nuclear symbols ... is described for some neutral atoms ...
  • , atom of hydrogen–1, symbol H, mass 1, just 1 proton and one electron, but NO neutrons (unique, hydrogen-1 is the only atom with NO neutrons)
  • , atom of helium–4, symbol He, mass 4,  2 protons, 4 – 2 = 2 neutrons and 2 electrons
  • , atom of sodium, symbol Na, mass 23, 11 protons, 23 – 11 = 12 neutrons and 11 electrons

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

• If the proton number and electron number are different, the atom has an overall surplus or deficiency of electrical charge, resulting in an electrically charged particle called an ion e.g.
  • If an atom loses 1 or 2 electrons, the protons produce an excess of 1 or 2 units of positive charge.
    • The excess positive charge on the positive ion is written as + or 2+ etc.
  • If an atom gains 1 or 2 electrons, the extra electrons produce an excess of 1 or 2 units of negative charge.
    • The excess negative charge on the negative ion is written as - or 2- etc.
  • So, reading the symbols below, as above, BUT, now taking into account the electrical charge on the ion ..
  • the positive sodium ion Na⁺,11 protons, 10 electrons (11-1), 12 neutrons (23-11)
  • the positive magnesium ion Mg²⁺,12 protons, 10 electrons (12-2), 12 neutrons (24-12)
  • the negative chloride ion Cl^–,17 protons, 18 electrons (17+1), 20 neutrons (37-17)
  • the negative sulfide ion S^2-,16 protons, 18 electrons (16+2), 16 neutrons (32-16)
  • for more details and examples see ionic bonding notes.

• The electrons are arranged in specific energy levels according to a set of rules (dealt 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

    • Therefore from the following 'full' atomic symbols, assuming we are dealing with electrically neutral atoms, the number of sub-atomic particles for the following atoms will be as follows ...

    • 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!

• Isotopes are atoms of the same element with different numbers of neutrons in the nucleus and therefore atoms of the same element with different masses (different nucleon number or mass number).
  • Isotopes are atoms of the same atomic number but different mass numbers.
    • Some elements have just one isotope but others may have up to eight different isotopes.
    • Isotopes of an element only differ in the number of neutrons in the nucleus.
    • i.e. all the isotopic atoms of an element have the same number of protons, electrons and electronic structure - so isotopes of an element are chemically identical - same reactivity, same formula compounds etc. etc. !
    • Most elements have one or more stable isotopes, but many other isotopes are unstable, disintegrate spontaneously (nuclear decay) and are known as radioactive.
  • This gives each isotope of a particular element a different mass or nucleon number, but, being the same element they have the same atomic number or proton number, but different mass number.
  • Isotopes of a particular element are also chemically identical, because they have the same number of electrons, hence the same electron structure and will therefore behave in an identical manner to each other.
  • Study the diagrams of the isotopes of carbon further down the page.

  • Relative Isotopic Mass is dealt with on a separate calculation page

    • (and see also mass spectrometry)

• 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!
• hydrogen–1, hydrogen–2, and hydrogen–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 and 1 electron, 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 (atomic number 1).
•  and 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 and 2 electrons.
  • 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
•  and 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).
  • They both have 17 electrons surrounding the nucleus in the arrangement 2.8.7 (see section 3. on electron arrangement)
• 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.
  • Both isotopes will have 35 electrons arranged in the appropriate energy levels around the nucleus.
• Note: You can have two different isotopes with the same mass number BUT they must be of different elements!
  • e.g. is the nuclear symbol for the isotope magnesium-24 (12 protons, 12 neutrons, 12 electrons)
  • and is the nuclear symbol for the isotope sodium-24 (11 protons, 13 neutrons, 11 electrons).
  • Apart from the obviously different chemical symbol, the atomic numbers are different too!
• The three known isotopes of carbon (the electron structure is fully explained in the next section 3.)
  • 
  • 

  • 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.
  • The electronic structure is also shown and is fully explained in the next section 3.
  • Two carbon isotopes are very stable, but carbon-14 slowly decays and is used by archaeologists to radiocarbon date objects. (See archaeological radiocarbon-14 dating)
  • Now is an appropriate point to introduce the concept and definition of relative atomic mass (A_r), 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, A_r(C) = 12.011, though at this academic level 12.0 is usually accurate enough!
    • 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.

    • See also Relative Isotopic Mass

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

• Nuclide notation and ions - need to a bit more careful in interpreting nuclear symbols of ions.

  • Ions are NOT neutral, they have an overall net electrical charge caused by the atom losing or gaining electrons to give a positive or negative ion. In other words the number of protons no longer equals the number of electrons.

  • This process is called ionisation and there is never any change in the number of protons or neutrons in the nucleus, just a change in the number of orbiting electrons.

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

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

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

• 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.
    • For lots more details see the RADIOACTIVITY NOTES
• DO NOT CONFUSE ISOTOPES and ALLOTROPES – see Appendix 3.

• The Bohr model of the atom in its more elaborate form involves the maximum numbers of electrons that each shell or energy level can hold and how the shells are progressively filled with electrons from atom to another with increase in proton/atomic number.

  • The volume or space occupied by the electronic energy levels determines the radius of an atom - referred to as the atomic radius. The space occupied by an energy level or shell is known as an orbital - but this is for more advanced A level chemistry courses!

• The electrons are arranged in energy levels or shells around the nucleus and with 'orbits' on average increasing in distance from the nucleus.

  • Electrons in an atom occupy the lowest available energy levels (the innermost available shells).

  • The lowest energy levels are always filled first, you can think of the lower the shell, the nearer the nucleus, and numbered 1, 2, 3 etc. as the shell gets further from the 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, up to at least 36 for Advanced level students).

  • The 1^st 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 19^th and 20^th electrons go into the 4^th shell, (required limit of GCSE chemistry knowledge).

  • Remember the total electrons to be arranged equals the atomic/proton number for a neutral atom.

• If you know the atomic (proton) number, you know it equals the number of electrons in a neutral atom, you then apply the above rules to work out the electron arrangement (configuration).

• For elements 1 to 20 the electron arrangements/configurations are written out in the following manner:
  • Note that each number represents the number of electrons in a particular shell, dots or commas are used to separate the numbers of electrons in each shell. They are written out in order of increasing average distance from the positive nucleus which holds these negative electrons in their energy levels (shells).
  • The electron configurations or electron arrangements are summarised below with reference to the periods of the periodic table and in order of increasing atomic number up to proton number 20.
    • For more see the Periodic Table and Electron Structure notes below.

  • Period 1 – elements 1 to 2 (2 elements)

    • the electron arrangement is written out simply as 1 or 2, only the 1st shell or level involved.

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

    • have electron arrangements of 2,1 to 2,8 (since 1st shell is full with 2 electrons i.e. the first number)

      • 2,5 would mean two electrons in the 1st level (shell) and five in the 2nd level (shell).

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

    • are denoted by 2,8,1 to 2,8,8 (1st & 2nd shells full with 2 & 8 electrons, ie 2.8)

      • 2,8,3 would mean two electrons in the 1st level (shell), eight in the 2nd level (full shell) and three electrons in the 3rd outer level (shell).

  • Period 4 – first two elements 19 to 20

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

      • 2,8,8,1 would mean two electrons in the 1st level (shell), eight in the 2nd level, eight electrons in the 3rd level (shell) and one electron in the 4th outer level (shell).

    • Reminder – this is as far as GCSE students need to know, after that things get more complicated, BUT only for advanced level students!

    • For example, after element 18, the 3rd shell can hold a maximum of 18 electrons!

  • The above is summarised in the diagram below

• 

• The first 20 elements of the periodic table in terms of their electronic structure in shells and numbers.

• I've repeated the same 20 elements of the periodic table showing simple diagrams of their electronic structure and group numbers with full diagrams of the shells of electrons.

• Group 1 are the alkali metals, Group 7 are the halogens and group 0/8 are the noble gases.

• You should note that the group number of the element equals the number of electrons in the outer shell (1-8).

• The electron shell arrangements are quoted in numbers e.g. 2,4 for C (carbon) but you need to be able to draw electron diagrams showing the electronic structure of the atom.

  • In the above table, check out the atomic number, the lower subscript on the element symbol, and apply the rules, and hopefully it makes sense.

  • Some examples of electronic diagrams are given below and GCSE/IGCSE/O level students need to be able to work and draw the electronic structures of the first 20 elements.

  • You should notice that the number of shells used equals the period number of the element in the 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 up to at least Z=36 using s, p, d and f orbital notation BUT this is for advanced A 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 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!)

A few more 'snappy' examples – given atomic number, work out electron configuration (abbreviated to e.c.)

Z = 3 e.c. 2,1  or   Z = 7 e.c. = 2,5  or  Z = 14 e.c. = 2,8,4 or Z = 19 e.c. = 2,8,8,1 etc. up to Z = 20

• 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 [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.

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

  • GCSE/IGCSE/AS notes on CHEMICAL BONDING

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

  • GCSE/IGCSE notes REACTIVITY SERIES of METALS

  • GCSE/IGCSE notes Group 1 ALKALI METALS

  • GCSE/IGCSE notes GROUP 7 HALOGENS

• 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 'atomic weight' (now called relative atomic mass). This is not correct for some elements now that we know their detailed atomic structure in terms of protons, neutrons and electrons, and of course, their chemical and physical properties in more accurate and extensive detail.

• 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.

• 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 (e.g. 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 number.

• The periodicity of elements i.e. the repetition of very chemically similar elements in a group is due to the repetition of a the same outer electron structure – check out the last number from element 3 onwards.

• More GCSE/IGCSE notes on the Periodic Table

• and the electronic explanations of chemical bonding–formulae

• Advanced Level Chemistry – electron configurations/arrangements and the Periodic Table

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APPENDIX 1. The history of the atom concept, development of atomic structure models and the famous alpha particle scattering experiment

The development of the 'atomic model' is an excellent example of how new experiment evidence initiates the need to change an existing scientific model or even come up with a different theoretical model, in this case for atomic structure. You should appreciate that knowledge and understanding of atomic structure has evolved over time and as a scientific model of atomic structure, it must explain current experimental observations and pose questions for future investigations.

Any further change in an atomic structure hypothesis e.g. because of new evidence, must be re-tested out in the laboratory and the results checked by other research groups from around the world. Plus a transparent peer group review of any research paper to be published, which means scientists having their work checked by other scientists. I they don't agree, somebody has got something wrong or made false claims or just done the experiment badly! If further experimental checks don't back up a hypothesis, it must be modified or abandoned in favour of other ideas.

When a hypothesis is backed up by experimental evidence from different scientists and lots of cross-checking of results e.g. the Bohr theory of the atom, it becomes an accepted theoretical model of an atom. Even to this day the atomic structure model is still developing e.g. the hypothesis that neutrons and protons are made of quarks which are held together by gluons, but at this point I'm out of my depth! Nothing stands still in scientific theory, even with atomic structure, and the hypotheses of modern quantum physicists about the structure of the nucleus I find pretty 'whacky', actually, pretty incomprehensible!

The ancient Greeks thought that everything was made of four basic 'elements' - air, earth, fire and water. However, at the height of the great classical Greek civilisation, the Greeks Leucippus and Democritus and others ~400 BC, wondered what was the result of continually dividing a substance i.e. what was the end product or smallest bit i.e. what was left that was indivisible – the word atomic is from Greek adjective meaning 'not divisible'.

The Greeks idea was not forgotten and later revived by Boyle and Newton but with little progress.

However, in 1808 Dalton at the beginning of the 19th century, proposed his atomic theory that all matter was made up of tiny hard particles/spheres called atoms.

Dalton also proposed (correctly) the theory that different types of atoms (elements) combined together to give all the different substances of the physical world (all which of course is true, except for the 'hard solid 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.

He was incorrect by stating that atoms were indivisible, because we now know that atoms consist of electrons, protons and neutrons and that atoms can be 'taken apart' by ionisation or nuclear changes  e.g. radioactivity.

Until the discovery of the electron, atoms were thought of as hard indivisible spheres, but brilliant 'JJ' Thompson changed all that. New experimental evidence led to a new scientific model of the atom with the discovery of the electron, recognised as the first known 'sub-atomic particle', change was on the way.

J J Thomson around 1897 proposed his 'plum pudding model' theory (picture on right) based on the growing evidence that atoms were themselves composed of even smaller more fundamental particles like the electron i.e. atoms were not hard indivisible spheres, so the solid sphere had to go.

He based his model on experiments that showed that atoms contained even smaller negatively charged particle called electrons which could be removed from atoms using a vacuum tube and applying a high potential difference (voltage) to a very low pressure gas. He showed that the mass of an electron was much less than the mass of an atom and that it had a negative electric charge. Therefore the 'hard indivisible sphere' model of an atom was wrong.

From his experiments Thomson envisaged a plumb pudding atom consisting of a positively charged 'pudding' (a sort of ball of positive charge) with just enough lighter negatively charged electrons embedded in it to produce a neutral atom. Note that both the positive charges and negative charges are evenly distributed through the sphere of the atom (shown later by Rutherford and Bohr etc. to be completely wrong).

The idea of positive particles balancing the negative particles was correct but the relative size and nature of the nucleus and distribution of electrons were not, BUT it was a more advanced model.

Ernest Rutherford, assisted by Hans Geiger and Ernest Marsden (the latter two were students of Rutherford at Cambridge University) conducted alpha particle scattering experiments (1902–1910, and described in detail below). The famous alpha particle experiment in 1909 was designed to test the plum pudding theory of JJ Thomson.

By 1911, these experiments established

(i) minute nature of the nucleus even compared to the size of an atom.

(ii) the nucleus was positive and the positive charge varied from element to element.

(iii) the positive charge was concentrated in the nucleus and able to deflect other positive particles e.g. alpha particles.

Diagram of the famous Rutherford and Geiger–Marsden alpha particle scattering experiment

When positive alpha particle beams are fired on very thin layers of metals (e.g. very fine gold leaf) some rather surprising results were made by scientists of the early 20th century.

By using a 360^o charged particle detection system it was found that ...

3. most particles passed through un–deflected (as if there was nothing there!), this was expected. This was predicted from JJ Thompson's plumb pudding model, but all the alpha particles were expected to pass through or to be slightly detected (observation 2. below), though NOT big deflections.

2. a small proportion were deflected slightly (so there was something there!), this again was not unexpected.

1. about 1 in 20,000 were 'bounced' back through an angle of over 90^o, in other words were reflected backwards, a totally unexpected result and quite shocked the experimenters - not what they were expecting. This was because the JJ Thompson's plumb pudding model predicted the positive charge was spread out and not sufficiently concentrated to cause, for some alpha particles, a 180^o deflection! So, whatever was there, was substantial in mass and positive charge to cause the repulsion 'bounce' of the positive alpha particles, BUT what it was (the 'nucleus') it wasn't very big!

These results made it quite plain the JJ Thompson plumb pudding model was in some way wrong ie the positive charge was NOT spread throughout the volume of an atom, therefore a new model must be proposed to take into account the new results.

From a detailed mathematical analysis of the scattering experiment results, the only 'atomic model' which could account for the pattern was an atom consisting of ...

1. mainly empty space (which is why most alpha particles passed through undeflected), thus completely contradicting JJ Thompson's 'plum pudding' model. Other experiments showed that the electrons were orbiting in energy levels around the nucleus, but occupying virtually no significant volume in themselves as particles.

2a. a relatively minute positive centre (the nucleus) causing deflection (like charges repel, alpha particles are positively charged and so were being repelled by the 'later to be discovered' positive protons in the nucleus), we now know the nucleus is positive due to protons,

2b. a tiny dense centre of similar or greater charge or mass to an alpha particle (which we now call the nucleus), so most of the mass of an atom was in the central nucleus, we know the mass is made of protons and neutrons.

3. Most of the atom is mainly empty space with a cloud of negative electrons moving around the dense relatively massive positive nucleus.

Putting these three points together formed the basis of the modern picture of the 'nuclear atom', in other words the nuclear atomic model.

BUT, there was still a puzzle to solve - why didn't the negative electrons collapse into the nucleus?

The great physicist Niels Bohr suggested the electrons orbited in energy levels (shells) that prevented electrons from being attracted into the nucleus - but would experiments confirm this theory?

Later experiments did show that electrons are arranged in energy levels, sort orbits around the nucleus, ideas first proposed by scientists such as Bohr.

Niels Bohr adapted the nuclear model by suggesting that electrons orbit the nucleus at specific distances from the nucleus and each orbit was a specific electron energy level - a fixed electronic energies.

The theoretical calculations of Bohr agreed with experimental observations.

Later experiments led to the idea that the positive charge of any nucleus could be subdivided into a whole number of smaller particles, each particle having the same amount of positive charge.

The name proton was given to these particles.

This was a necessary extension and modification to the Rutherford model of an atom, because this model could not account for why the electrons were not attracted to the nucleus.

Bohr's suggested that the negative electrons can only exist in certain specific energy levels (shells) at fixed distances from the nucleus and held in place by the positive nucleus. This theory added too, and complimented the Rutherford model of the atom, to gives a reasonably complete picture of an atom (at least for this academic level!)

e.g. on the right the 'Bohr' electronic diagram for sodium with the (Na) representing the nucleus.

and below, a more sophisticated diagram of a lithium atom.

Bohr envisaged the electrons orbiting the nucleus in specific energy levels (or fixed shells) at specific distances from the central nucleus with nothing in between. In other words the electrons have sufficient energy to keep away from the nucleus and be confined to these specific energy levels. The negative charge of the electrons was still balanced by the positive charge (protons) of the nucleus.

So now, as far as we can tell (at GCSE/A level anyway) an atom is quite well represented by the Bohr model of the atom (picture below) which moves the Rutherford nuclear model another step forward.

  Diagrams of a 'Bohr' lithium atom

AND, most importantly, experimental results matched a theoretical mathematical model of simple 'electronic' atoms like hydrogen.

So by now, earlier theories of atomic structure, e.g. the 'plum pudding' model in which 'protons' and 'electrons' were scattered or arranged evenly across the atom, were superseded by the nuclear model of Rutherford and subsequently this was superseded by Bohr's electronic energy level model.

It was the only model that could explain the scattering of the high speed alpha particles by a small dense and positive atomic centre AND the behaviour of electrons.

Experiments had shown that the outer bits could be knocked off atoms and these had a very tiny mass and a negative charge, in other words the electron!

Further experiments showed that the nucleus (partly) consisted of positive particles with the same mass and charge as an ionised hydrogen atom, that is a proton (mass 1, charge +1).

(NOT for GCSE level) In 1913 Moseley studied the X–rays emitted by highly energised–ionised atoms and from the X–ray spectra of elements (the K alpha line, K_α) he was able to deduce the electric charge of the nucleus which we now know is equal to the atomic number of protons in the nucleus.

Moseley showed that when atoms were bombarded with cathode rays (electrons) X–rays where produced which he investigated with an X-ray spectrometer.

It was found that the square root of the highest energy emission line (called the K alpha line, K_α) gave a linear plot with the apparent atomic number Z (it wasn't known yet that this was the proton number),

Z = constant x √K_α

but the plot of √K_α against atomic weight (relative atomic mass) gave a zig–zag plot, suggesting this 'atomic number' was far more important the 'atomic weight' of an element in terms of the atom's fundamental structure.

Note:

(i) The K alpha line, K_α is due to an electronic transition of the inner most electron nearest the nucleus.

(ii) Sadly, Moseley was killed in action during the First World War at Gallipoli in 1915, a great loss to science as well as his family and friends.

(iii) We now know that Moseley's 'atomic number' is in fact the number of protons in the nucleus (atomic number = proton number).

(iv) By 1898, thanks to the German scientist Wilhelm Wien, the hydrogen ion (proton) had been identified as the simplest basic unit of positive charge. By 1925, later experiments by Rutherford and others, identified the number Z as the value of the positive charge of the nucleus and that it equated to Z protons in the nucleus - hence the atomic number = proton number = Moseley's Z value.

However, there was still the problem of why the atomic mass and atomic number where different i.e. in the case of the lighter elements, the atomic weight was often about twice the atomic number.

In 1919 Aston developed a cathode ray tube i.e. like those used by Wien and Thompson etc. into a 'mass spectrograph', which we now know as a mass spectrometer GCSE–AS atomic structure notes.

This showed that atoms of the same element had different masses but there was no experimental evidence that they had different atomic numbers (which of course they didn't). These different atoms of the same element were called isotopes.

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.

These were shown to have a relative mass of 1 (same as a proton) and were electrically neutral and quite penetrating into matter. This penetration and lack of charge had made them difficult to detect.

Prior to this, Rutherford and others had conducted experiments to show that the smallest particle in an atom was equivalent to a hydrogen atom without its electron, that is the proton.

It was not until 1932 that the nature of the neutron was finally deduced by Chadwick, and he showed that the nucleus also contained an electrically neutral particle of similar mass to a proto, and this completely explained the nature of isotopes and backed up the ideas from Moseley's work that the fundamentally important number that characterises an element is its atomic number and NOT the atomic mass (or mass number). The neutron discovery, ~20 years after the discovery of the nucleus, completed the 'modern' picture and theory of the composition of an atom in terms of the three principal sub-atomic particles - which is sufficient for the needs of us chemists!

Advanced level note on the discovery of neutrons:

Chadwick bombarded a thin metal foil of beryllium atoms (⁹₄Be) with alpha particles (⁴₂He) and this produced a highly penetrating radiation that was unaffected by electrical or magnetic fields.

⁹₄Be  +  ⁴₂He  ====>  ¹²₆C  +  ¹₀n

This 'neutral' penetrating radiation was eventually identified as a beam of neutrons!

The neutron wasn't detected directly, the neutron beam knocks protons out of atoms which could be detected.

It is very difficult to detect a neutral particle because it doesn't do anything in electric or magnetic fields.

There are no deflection observations to put into a mathematical model to work out the mass or charge of the deflected particle e.g. from the sort of experiments JJ Thompson did.

Note the carbon-12 atom produced balances the nuclear equation in terms of mass (9 + 4 = 12 + 1) and positive proton nuclear charge (4+ + 2+ = 6+).

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

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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

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Appendix 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. 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 a very reactive 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, but they behave chemically in an identical manner.

BUT 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.

Allotropes can have different physical properties (e.g. density, melting point) and different chemical properties (undergo different reactions or can differ in reactivity for the same reaction).

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Appendix 4. The effect of ionizing radiation on atoms

Ionisation by EM radiation of atoms to form a positive ion

The higher energy uv, and both X-ray and gamma radiations have enough energy to cause complete ionisation of an atom.

Ionisation has already been explained on this page, but a reminder won't go amiss!

An atom is ionised if it completely loses one or more electrons to form a positive ion.

The process is illustrated below for the formation of a monopositive ion of sodium from a sodium atom.

In terms of electrons, the change is ...

  ==  high energy uv/X-ray/gamma ray photon  ==>  ⁺    +    electron^-

This represents the ionisation of a sodium atom to form a positive sodium ion and a free electron:

Na ==> Na⁺  +  e¯  (electron configuration change of sodium from 2.8.1 ==> 2.8, as in chemistry notes!)

In this case the incoming EM radiation must have sufficient energy to promote the electron all the way up the energy levels until it is completely free of the attraction of the positive nucleus - so the atom has been ionised.

A positive ion is formed because there are now less negative electrons on the atom than positive protons - so there is a surplus of positive charge on the atom.

The more electrons 'knocked off' the bigger the positive charge on the ion.

i.e. the charge on the ion can be +, 2+, 3+ etc. by knocking off 1, 2 or 3 electrons etc. The more electrons knocked off, the bigger the positive charge on the ion.

(Note: Using X-rays, you can knock off all 92 electrons from a uranium atom, element 92, ₉₂U, to form the U⁹²⁺ ion, but this far too extreme for GCSE students!, but very exciting to contemplate!)

More on ionisation and ionising radiation in the following notes

Alpha, beta & gamma radiation - properties of 3 types of radioactive nuclear emission & symbols ,dangers of radioactive emissions - health and safety issues and ionising radiation gcse physics revision notes

and Electromagnetic radiation, sources, types, properties, uses (including medical) & dangers 

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Appendix 5. Mass Spectrometer (now on a separate page for advanced level students!)

See also Atomic structure and radioactivity

other associated web pages of notes or quizzes

FT GCSE/IGCSE Foundation Atomic Structure multiple choice QUIZ

HT GCSE/IGCSE Higher Atomic Structure multiple choice QUIZ

GCSE/IGCSE Atomic Structure Crossword Puzzle * ANSWERS!

GCSE/IGCSE multi–word gap–fill worksheet on atomic structure

Word-fill quiz "Atomic structure and elements"

2nd Word-fill quiz "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

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