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GCSE-KS4-IGCSE-AS Science-CHEMISTRY revision-information notes on

 Atomic Structure, Isotopes, Periodic Table and Electronic Structure of Atoms

Atomic Structure page CONTENTS

1. The Structure of Atoms - fundamental particles
2. Isotopes
3. The Electronic Structure of Atoms - rules to be learned
4. Which electron arrangements are stable and which are not?
5. The Periodic Table and Electronic Structure - more patterns!
Appendix 1. The Alpha Particle Scattering Experiment
Appendix 2. Atomic structure diagrams - some variations!
Appendix 3. Allotropes - don't confuse with isotopes!
Appendix 4 The Mass Spectrometer (for advanced level students)

Atoms are the smallest particles of matter whose properties we study in Chemistry. 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 sub-atomic particles, protons, neutrons and electrons, which are listed below with their relative masses and electrical charges.

A Portrait of an Atom - an image of what you can't see!

• The diagram below gives some idea on the structure of an atom, it also includes some important definitions and notation used to describe atomic structure.
• Protons and neutrons are the 'nucleons' or 'sub-atomic' particles present in the minute positive nucleus and the negative electrons are held by the positive protons in 'orbits' called energy levels or shells.
• Some important evidence for this 'picture' is obtained from alpha particle scattering experiments (see Appendix 1).
• The atomic/proton number (Z) is the number of protons in the nucleus and is also known as the proton number of the particular element.
• It is the proton/atomic number (Z) that determines the number of electrons an element has, its specific electron structure and therefore the specific identity of a particular element in terms of its physical and chemical properties. It cannot be overemphasised that it is the electronic structure that determines the chemical character of an element, hence the proton/atomic number determines everything about a particular element element.
• The mass number (A) is also known as the nucleon number, that is the number of particles in the nucleus of a particular isotope (notes on isotopes - definition and examples).
• The neutron number (N) = mass number - proton/atomic number
• In a neutral atom the number of protons (+) equals the number of electrons (-), that is the number of positive charges is equal to the number of negative charges.
  • If not, the atom has an overall surplus electrical charge and is then called an ion e.g.
    • the positive sodium ion Na⁺ (11 protons, 10 electrons, excess positive protons),
    • or the negative chloride ion Cl^- (17 protons, 18 electrons, excess negative electrons,
    • for more details and examples see ionic bonding notes.
• Other more 'practical' diagrams than the one below are shown in Appendix 2.
• In the example below for lithium-7, the nuclide notation states that
  • before the chemical symbol of the element
  • the top left number  = nucleon/mass number = 7
  • and the bottom left number = proton/atomic number = 3

• Isotopes are atoms of the same element with different numbers of neutrons. This gives each isotope of a particular element a different mass or nucleon number, but, being the same element they have the same atomic or proton number and are dentical chemically.
• There are small physical differences between the isotopes e.g. the heavier isotope has a greater density or boiling point.
• However, because they have the same number of protons (proton/atomic number) isotopes of a particular element have the same electronic structure and identical chemistry. Examples are illustrated below.
• Do NOT assume the word isotope means the atom it is radioactive, this depends on the stability of the nucleus i.e. unstable atoms (radioactive) might be referred to as radioisotopes. Many isotopes are stable and NOT radioactive i.e. most of the atoms that make up you and the world around you!
• , and are the three isotopes of hydrogen with mass numbers of 1, 2 and 3, with 0, 1 and 2 neutrons respectively, but all have 1 proton. Hydrogen-1 is the most common, there is a trace of hydrogen-2 naturally but hydrogen-3 is very unstable and is used in atomic fusion weapons.
•  and are the two isotopes of helium with mass numbers of 3 and 4, with 1 and 2 neutrons respectively but both have 2 protons. Helium-3 is formed in the Sun by the initial nuclear fusion process. Helium-4 is also formed in the Sun and as a product of radioactive alpha decay of an unstable nucleus. An alpha particle is a helium nucleus, it picks up two electrons and becomes the stable atoms of the gas helium.
•  and are the two isotopes of sodium with mass numbers of 23 and 24, with 12 and 13 neutrons respectively but both have 11 protons. Sodium-23 is quite stable e.g. in common salt (NaCl, sodium chloride) but sodium-24 is a radio-isotope and is a gamma emitter used in medicine as a radioactive tracer e.g. to examine organs and the blood system.
• The relative atomic mass of an element is the average mass of all the isotopes present compared to 1/12th of the mass of carbon-12 atom (¹²C = 12.00000 i.e. the standard). See also calculations page section 
• DO NOT CONFUSE ISOTOPES and ALLOTROPES - see Appendix 3.

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

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

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

• There are rules about the maximum number of electrons allowed in each shell and you have to be able to work out the arrangements for the first 20 elements (see the Periodic Table diagrams further down).

  • The 1^st shell has a maximum of 2 electrons

  • The 2nd shell has a maximum of 8 electrons

  • The 3rd shell has a maximum of 8 electrons

    • This is only up to atomic number 20, after that it is for Advanced Level Chemistry Students - electron configurations/arrangements!

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

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

Examples: diagram, symbol or name of element (Atomic Number = number of electrons in a neutral atom), shorthand electron arrangement.

• When an atom has its outer level full to the maximum number of electrons allowed, the atom is particularly stable electronically and very unreactive.

  • This is the situation with the Noble Gases: He is [2], neon is [2,8] and argon is [2,8,8] etc.

  • There atoms are the most reluctant to lose, share or gain electrons in any sort of chemical interaction because they are so electronically stable - most of chemistry is about what outer electrons do or don't!

  • [2],[2,8] and [2,8,8] etc. are known as the 'stable Noble Gas arrangements', and the atoms of other elements try to attain this sort of electron structure when reacting to become more stable.

• The most reactive metals have just one outer electron.

  • These are the Group 1 Alkali Metals, lithium [2,1], sodium [2,8,1], potassium [2,8,8,1]

  • With one outer shell electron, they have one more electron than a stable Noble Gas electron structure.

  • So, they readily lose the outer electron when they chemically react to try to form (if possible) one of the stable Noble Gas electron arrangements - which is why atoms react in the first place!

  • When group 1 metal atoms lose an electron they form a positive ion because the positive proton number doesn't change, but with one negative electron lost, there is a surplus of one + charge e.g.

    • sodium atom ==> sodium ion, Na ==> Na⁺ is [2.8.1] ==> [2.8] electronically.

    • Ions are atoms or group of atoms which carry an overall electrical charge i.e. not electrically neutral.

• The most reactive non-metals are just one electron short of a full outer shell.

  • These are the Group 7 Halogens, namely fluorine [2,7], chlorine [2,8,7] etc.

  • These atoms are one electron short of a stable full outer shell and seek an 8th outer electron to become electronically stable - yet again, this is why atoms react!

  • They readily gain an outer electron, when they chemically react, to form one of the stable Noble Gas electron arrangements either by sharing electrons (in a covalent bond) or by electron transfer forming a singly charged negative ion (ionic bonding) e.g.

    • chlorine atom ==> chloride ion, Cl ==> Cl^- is [2.8.7] ==> [2.8.8] electronically,

    • the positive proton number of Cl doesn't change but the chloride ion carries one extra negative electron to give the surplus charge of a single - on the ion.

• The elements are laid out in order of Atomic Number.

• It is important to realise that the 'chemical structure' of the periodic Table (shown above), that is the chemical similarity of vertical groups 'like' elements (apart from the Noble Gases), was known well before the electronic structure of atoms was understood. However, it wasn't understood why they behaved in the same way chemically e.g. similar compound formulae and reactions etc. nor was it understood at first why Noble Gases were so unreactive towards other elements. BUT once the electronic structure of atoms was understood, 'electronic' theories could then be applied to explain the chemical similarity of elements in a vertical Group of the Periodic Table.

• Originally they were laid out in order of ' relative atomic mass' (the old term was 'atomic weight'). This is not correct for some elements now that we know their detailed atomic structure in terms of protons, neutrons and electrons, and of course, their chemical and physical properties.

• For example: Argon (at. no. 18, electrons 2,8,8) has a relative atomic mass of  40. Potassium (at. no. 19, electrons 2,8,8,1) has a relative atomic mass of 39. Argon, in terms of its physical, chemical and electronic properties is clearly a Noble Gas in Group 8 (0). Likewise, potassium is clearly an Alkali Metal in Group 1.

• Hydrogen, 1, H, does not readily fit into any group

• A Group is a vertical column of chemically and physically similar elements e.g. Group 1 The Alkali Metals (Li, Na, K etc.), Group 7 The Halogens (F, Cl, Br, I etc.) and Group 8 or 0 The Noble Gases (He, Ne, Ar etc.). The group number equals the number of electrons in the outer shell (e.g. chlorine's electron arrangement is 2.8.7, the second element down Group 7 on period 3).

• A Period is a horizontal row of elements with a variety of properties (left to right goes from metallic to non-metallic elements. All the elements use the same number of electron shells which equals the period number (e.g. sodium's electron arrangement 2.8.1, the first element in Period 3).

• The ten elements Sc to Zn are called the Transition Metals Series and form part of a period between Group 2 and Group 3 from Period 4 onwards.

• Below are the electron arrangements for elements 1 to 20 set out in Periodic Table format (Hydrogen and The Transition metals etc. have been omitted). When you move down to the next period you start to fill in the next shell according to the maximum electrons in a shell rule (see previous section).

• 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 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 similar outer electron structure.

• More GCSE notes on the Periodic Table and the electronic explanations of chemical bonding-formulae

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

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APPENDIX 1. The Alpha Particle Scattering Experiment

• The Rutherford and Marsden scattering experiment.

• When alpha particle beams are fired on very thin layers of metals (e.g. very fine gold leaf) the results were found to be rather surprising to scientists of the early 20th century

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

  • most particles passed through un-deflected

  • a small proportion were deflected slightly

  • and about 1 in 20,000 were 'bounced' back through an angle of over 90^o

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

  • mainly empty space (why most alpha particles passed through),

  • a positive centre causing deflection (like charges repel), alpha particles are positively charged and so are the 'later to be discovered' protons in the nucleus,

  • a tiny dense centre of similar or greater charge or mass to an alpha particle (which we now call the nucleus),

  • in other words, an atom is well represented by the diagram near the top of this page.

• Earlier theories of atomic structure, e.g. the 'plum pudding' model in which 'protons' and 'electrons' were scattered or arranged evenly across the atom, were superceded by the model described in the previous picture.

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

• Later experiments showed that the out bits could be knocked off atoms and these had a very tiny mass and a negative charge, in other words the electron!

• See section 4. Radioactivity Notes page on other experiments with mixed particle beams and their separation.

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Appendix 2. Atomic structure diagrams - some variations!

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

• Isotopes are atoms of the same element with different masses due to different numbers of neutrons in the nucleus. Same protons and electrons. e.g. atomic number 6 = 6 protons = carbon, but there can be 6, 7 or 8 neutrons giving isotopes of carbon-12, 13 or 14.

• Oxygen atoms usually form 'stable' O₂ oxygen molecules (also called dioxygen), BUT they can form an unstable molecule O₃ ozone (also called trioxygen). The mass of the oxygen atoms in each of the molecules is mainly 16 (99.8%), and about 0.2% of two other stable isotopes of masses 17 and 18. Whatever isotope or isotopes make up the molecule, it doesn't affect the molecular structure or the respective chemistry of the O₂ or O₃ molecules.

• However, what sometimes confuses the issue is the fact that oxygen O₂ and ozone O₃ are examples of allotropes.

• Allotropes are defined as different forms of the same element in the same physical state.

• They are usually chemically similar but always physically different in some way e.g.

  1. O₂ (oxygen, dioxygen) and O₃ (ozone, trioxygen) are both gases but have different densities, boiling points etc.

  2. 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)

  3. Rhombic and monoclinic sulphur have different geometrical crystal structures, that is different ways of packing the sulphur atoms (which are actually both made up of different packing arrangements of S₈ molecules). They have different solubilities and melting points.

• It doesn't matter which isotopes make up the structure of any of an elements allotropes described above, so to summarise by example!,

  oxygen-16, 17 or 18 are isotopes of oxygen with different nuclear structures due to different numbers of neutrons,

  and O₂ and O₃ are different molecular structures of the same element in the same physical state and are called allotropes irrespective of the isotopes that make up the molecules.

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Appendix 4 The Mass Spectrometer (for advanced level students only)

• The mass spectrometer is an instrument by which you can separate ionised/charged (+) particles of different mass and determine the amounts of each particle in a mixture.

• The technique is called mass spectroscopy or mass spectrometry ('mass-spec' and 'MS' in shorthand!).

• The substance to be analysed is introduced/injected into a high vacuum (extremely low pressure) tube system (at K left diagram) where the particles are ionised by colliding with beam of high speed electrons (at Q in left diagram).

  • Note: If the sample is not already a gas, then substance must be vapourised, i.e. the material must be in the gaseous state to be analysed in a mass spectrometer.

• The resulting (+) ions are accelerated down a tube (from + to - plates, P in left diagram) and then through a powerful magnetic field.

• The charged or ionised particles are deflected by this powerful magnetic field (R in left diagram).

• How much they are deflected depends on the particle mass and the speed of the particle and the strength of a magnetic field i.e. lighter particles of lower mass (and momentum) are deflected more than heavier particles of bigger mass (see right diagram below) for a given set of conditions.

• By varying the strength of the magnetic field, it is possible to bring into focus onto an ion detector (N in left diagram) at the end of the tube (effectively an electrical event is detected), every possible mass in turn and a measure the strength of the ion current, which is a measure of how much of that ion has been formed from the sample under analysis.

• A simplified diagram of a mass spectrometer tube system is shown below (left) with further explanation as to what is going on and further diagram to show the relative paths of light to heavy ions for a given strength of magnetic field.

• Key to diagram and more detail of each component's function.

  • K = sample injection point, it must be a gas, so a liquid/solid must be vaporised at the injection point.

  • Q = high voltage (high +/- p.d.) electron gun which fires a beam of high speed/energy electrons from a heated 'metal element' into the vaporised sample under analysis and causes ionization of the atoms (or molecules) forming positive ions (mainly monopositive in charge).

    • The collision of high KE electrons with atoms or molecules causes another electron to be knocked off the particle leaving a negative deficit i.e. a positively charged particle is formed e.g.

      • M_(g) + e^- ==> M⁺_(g) + 2e^-, usually written as just

      • M_(g) ==> M⁺_(g) + e^- (M might represent e.g. a metal atom or a molecule)

      • The ions formed should be written as [M]⁺, a notation that is handy if you are dealing with ionised molecule fragments with an overall single positive charge e.g. [CH₃]⁺ is seen in the mass spectrum of methane gas, CH₄.

      • The low pressure (~vacuum) is needed to prevent the ions from colliding with air particles which would stop them reaching the ion detector system.

  • P = are negative plates which accelerate the positive ions down the tube (there are positive plates at the start of the tube). A moving beam of charged particles creates a magnetic field around itself, and this 'ion beam' magnetic field interacts with the magnetic field at R.

  • R = the magnetic field that causes deflection of ions, this is can be varied to change the extent of deflection for a given mass and to focus a beam of it down onto the detector. Hence, by programming the mass spectrometer to 'sweep' through all likely particle masses, in terms of the right hand diagram, you can increase the strength of the magnetic field to bring into focus onto the ion detector monopositive ions of increasing mass.

  • N = an ion detection system which essentially generates a tiny electrical current when the ions hit it. The strengths of the 'electronic' signals from the various ion peaks are sent to a computer for analysis, computation and display. They tell you the particle masses present and their relative abundance (see the mass spectrum diagram for the element strontium below).

  • The resulting record of the ion peaks is called the mass spectrum or mass spectra. The highest peak is called the base peak and is often given the relative and arbitrary value of 100, particularly in the mass spectra of organic compounds).

  • For elements you get a series of signals or ion peaks for each isotope present and the ratio of peak heights gives you the relative proportion of each isotope in the element so that you can calculate the relative atomic mass of an element. This 'simple' spectra of mononuclear ions like Sr⁺ is only true for non-molecular elements like metals (see mass spectrum of strontium diagram below) or noble gases, but for molecular elements like nitrogen or the halogens things are not so simple (see chlorine example below).

  • For larger e.g. organic molecules, things can be very complex indeed, as molecules fragment and many different ions can be formed.

• Chlorine is a good example of a molecular element whose mass spectra can be a bit tricky when first encountered ...

  • Chlorine consists of two principal isotopes, chlorine-37 (25% is ³⁷Cl) and chlorine-35 (75% is ³⁵Cl).

  • BUT, chlorine consists of Cl₂ diatomic molecules, which, on ionisation, can split into chlorine atoms.

  • The result is a series of 5 different mass peaks from the various isotopic atomic or molecular ion possibilities...

  • [³⁷Cl³⁷Cl]⁺ m/z = 74, [³⁷Cl³⁵Cl]⁺ m/z = 72, [³⁵Cl³⁵Cl]⁺ m/z=70, [³⁷Cl]⁺ m/z=37, [³⁵Cl]⁺ m/z=35

  • m/z means the relative mass of the ion over its charge, which for our purposes the charge is taken +1 (little z) and the mass (little m) is the relative atomic/formula mass of the particle.

  • Examples of the calculation of the relative atomic mass of an element using % of isotopes is given in Part 1 of GCSE-AS (basic) calculations, an example of calculating relative atomic mass from a mass spectrum is given below for the metallic element strontium.

• The mass spectra of organic compounds can be very complex as the molecules fragment under electron bombardment, but the resulting mass spectra can used to identify compounds from their 'finger-print' pattern of ion peaks of different mass and particular proportions for a given set of experimental conditions.

  • The largest m/z value gives the molecular mass of a molecule, i.e. the ion of largest mass, prior to fragmentation, is formed when the original whole and neutral molecule, loses one electron e.g. for ethane it would be due to the formation of [C₂H₆]⁺, m/z = 30 and is called the molecular ion peak.

• Example of a relative atomic mass calculation based on the mass spectrum of the element

• The relative atomic mass of an element, A_r, is the weighted average mass of the isotopes present, compared to 1/12th of the relative mass of the carbon-12 isotope. [ ¹²C is given the relative mass value of 12.0000 ]

• Quite often the highest peak is arbitrarily given the relative value of 100, as in this case, but the peak lines might well indicate % abundance of isotopes. relative peak height = relative abundance as measured from the ion current detector signal.

• The mass spectrum shows strontium consists of four isotopes, ⁸⁴Sr (peak height = 0.68), ⁸⁶Sr (peak height = 12.0),⁸⁷Sr (peak height = 8.47) and ⁸⁸Sr (peak height = 100.0)

• The sum of the heights = 0.68 + 12.0 + 8.47 + 100.0 = 121.15

• So we can now calculate the weighted average mass of ALL the isotopes.

• Therefore A_r = {(0.68 x 84) + (12.0 x 86) + (8.47 x 87) + (100.0 x 88)}/121.15 =  87.7

• The book value is 87.62, BUT this calculation does NOT take into account the very accurate relative atomic masses based on the carbon-12 scale, it merely uses the mass numbers, which are always integer. The very accurate isotopic masses are usually a tiny fraction different from a whole number.

• Modern mass spectrometers are exceedingly accurate and very sophisticated instruments and can measure mass to at least 4 decimal places. They can readily distinguish between N₂ and CO molecules, both with an M_r of 28.

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