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

Atomic Structure, Isotopes, Electronic Structure of Atoms

Appendix 4. The Mass Spectrometer (separate page for advanced level students!)

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

Compounds are formed when two or more elements are chemically combined to form a new substance in a reaction which is not easily reversed ie its difficult to separate a compound back into its constituent elements.

However, what are elements made up of? How and why do elements bond together?

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

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

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

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

However from experiments done in the late 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.

You can think of the mass of an electron as about ¹/₂₀₀₀th of the mass of a proton or neutron

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

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

• What sub–atomic particles make up atoms? What is their mass and charge?
• The diagram above of a 'portrait of an atom' gives some idea on the structure of an atom (sometimes called the Bohr Atomic Model), it also includes some important definitions and notation used to describe atomic structure
  • The three fundamental particles you need to know are ...
  • proton: particle mass = 1, electric charge = +1, the charged particle in the nucleus
  • neutron: particle mass = 1, charge = 0, uncharged particle in the nucleus
  • electron: particle mass = 1/1850 ~1/2000, electric charge = –1, NOT in nucleus but exist in electronic energy levels around the nucleus (a sort of orbit, often described as a shell, see later).
  • The nucleus of protons and neutrons is tiny, even compared to the tiny atom!
    • So most of the volume of an atom is where the electrons are, the diameter of the nucleus of protons plus neutrons is about a ten thousandth of the diameter of an atom!
    • Since the nucleus is composed of positive protons and neutral neutrons, the nucleus itself must be positive.
• Protons and neutrons are the 'nucleons' or 'sub–atomic' particles present in the minute positive nucleus and the negative electrons are held by the positive protons in 'orbits' called energy levels or shells.
  • Although the nucleus must be positive because of the positive protons (neutrons are neutral) an individual atom is neutral because the number of electrons equals the number of protons – so the charges 'cancel out'.
  • If electrons are removed from an atom you get a positive ion and if electrons are added to an atom you get a 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/proton number (Z) is the number of protons in the nucleus and is also known as the proton number of the particular element.
  • Each element has its own atomic number, so all the atoms of a particular element have the same atomic number.
• It is the proton/atomic number (Z) that determines the number of electrons an element has, its specific electron structure and therefore the specific identity of a particular element in terms of its physical and chemical properties.
• It cannot be overemphasised that it is the electronic structure that determines the chemical character of an element, hence the proton/atomic number determines everything about a 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).
  • Therefore the mass/nucleon number = sum of the protons plus neutrons in the nucleus.
• The neutron number (N) = mass number – proton/atomic number
• In a neutral atom the number of protons (+) equals the number of electrons (–), that is the number of positive charges is equal to the number of negative charges.
  • If not, and the atom has an overall surplus electrical charge, and is then called an ion e.g.
    • the positive sodium ion Na⁺ (11 protons, 10 electrons, excess positive protons),
    • or the negative chloride ion Cl^– (17 protons, 18 electrons, excess negative electrons,
    • for more details and examples see ionic bonding notes.
• Other alternative atomic structure diagrams than the one below are described and shown in Appendix 2.
• In the example below for lithium–7, the nuclide notation states that
  • before the chemical symbol of the element Li
  • the top left number  = nucleon/mass number = 7
  • and the bottom left number = proton/atomic number = 3

• The electrons are arranged in specific energy levels according to a set of rules (dealt 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 (iron–56), 26 protons, 30 neutrons (56 – 26), 26 electrons

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

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

• Isotopes are atoms of the same element with different numbers of neutrons and therefore different masses (different nucleon/mass numbers). This gives each isotope of a particular element a different mass or nucleon number, but, being the same element they have the same atomic or proton number and are identical chemically.
• The phrase 'heavier' or 'lighter' isotope means 'bigger' or 'smaller' mass number.
• There are small physical differences between the isotopes e.g. the heavier isotope has a greater density or boiling point, the lighter the isotope the faster it diffuses.
• However, because they have the same number of protons (proton/atomic number) isotopes of a particular element have the same electronic structure and identical 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!
• , and 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.
•  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. 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. 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 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 and 81–35 = 46 neutrons respectively. By coincidence, there are almost exactly 50% of each isotope present in naturally occurring bromine.
• The isotopes of carbon

  • 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

  • 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 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 accurate enough!
    • See also chemical calculations
• 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 electrons are arranged in energy levels or shells around the nucleus and with 'orbits' on average increasing in distance from the nucleus.

• Each electron in an atom is in a particular energy level (or shell) and the electrons must occupy the lowest available energy level (or shell) available nearest the nucleus.

• When the level is full, the next electron goes into the next highest level (shell) available.

• There are rules to learn about the maximum number of electrons allowed in each shell and you have to be able to work out the arrangements for the first 20 elements (for GCSE students, upto at least 36 for Advanced level students).

  • The 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 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.
    • 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

  • 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 2.8.8.1 and 2.8.8.2 (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!

  • 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 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. upto 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.

• 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 ...

  • 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 ' 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 Group 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