1a. The basic structure of the Periodic Table

Only the top portion of the periodic table is shown above (full version)

A very brief INTRODUCTION to the Periodic Table (part of which is shown above)

It is called the periodic table of elements because elements with similar properties occur at regular intervals - periodically.

Elements are the simplest substances that we can use and investigate in chemistry.

An element consist of one type of atom only.

Therefore, elements are the simplest substances that we can use and investigate in chemistry because an element cannot be split into other substances (unlike compounds).

Each element has identical atoms (except for isotopes, different numbers of neutrons) which are physically and chemically identical and each element has its own unique physical and chemical properties, but there are important patterns in properties when looking at a particular column (group) or row (period or series) of the periodic table.

Ever element has its own unique chemical symbol which is used to denote elements in the periodic table, in chemical formulae and chemical equations e.g. hydrogen is H, copper Cu, chlorine Cl or potassium K. The symbol is a single capital letter (upper case e.g. C, N, O, F, C, P etc.) or a capital letter followed by a lower case letter (e.g. Cu, Fe, Cl, Br, Li etc.).

Atoms of a particular element have identical numbers of protons (atomic number) which is often shown by the element symbol in the periodic table. This also means they all have the same number of electrons and the same electron arrangement (electron configuration). It is the unique electron arrangement of each element that determines its unique physical and chemical properties.

The elements are arranged in order of atomic number (proton number) so that very similar elements end up in vertical columns called groups AND they have the same number of outer electrons in the highest level containing electrons (except H and He) which gives them very similar chemical properties.

Over 100 elements are now known, but only 92 elements of the periodic table are found naturally on Earth as elements or (usually) compounds.

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

Picture examples of these electron arrangements (atomic number) are shown below

Note that the number of outer electrons (in the outer shell) equals the group number of 1 to 8

Group 1 alkali metals, Group 2 alkaline earth metals, group 7 halogens and group 8 noble gases - also designated group 0 because, until 1962, it was believed they couldn't form compounds with other elements i.e. to have zero valency.

For elements 3-20 you can easily work out the electron arrangement from the rules described above from the atomic number (= proton = electron number for a neutral atom. You therefore know which group the element is in from the last number (1-8)

This is not so for H and Group 0 (but effectively Group 8), noble gases were called group 0 because it was believed they could never form compounds - BUT not true since the 1960s

Examples of electron arrangement diagrams

Relate to the simplified periodic table above in terms of group AND period.

A period begins when the next electron is the first one in the next principal level/shell and the last element in a period is when the outer level/shell is full. So, apart from H & He, a period goes from a ...,1 to a ...,8 electron arrangement.

Filling 1st shell, electron level 1 2 elements only in Period 1

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], at the start of Period 4

Detailed Advanced A Level Chemistry Notes on the electronic structure of atoms

Finally, all the electronic diagrams of the first twenty elements in the context of the periodic table.

You should now fully appreciate the connection between the number of outer electrons and group numbers 1-8.

After group 7, all the noble gases have full outer electron shells, in fact 8 outer electrons after helium (just 2) which fits in with the outer shell electron patterns from group 1 to group 8. Note that Group 0 is also denoted by Group 8, which of course fits the outer electron pattern.

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1e. Valence and formula patterns in the Periodic Table

(this goes above GCSE level in places, but you are expected to work out a simple formula when group 1 and 2 elements combine with a group 6 or 7 element)

2a. Typical Properties of Metallic Elements

Physical properties of metals

Chemical Properties of metals

• These are typical electron changes for metals in groups 1, 2 and 3.
• The positive ions are formed directly from the metal atoms.
• Atoms usually react to give an electron arrangement with a full outer shell by losing, gaining or sharing electrons.
• Metallic elements on the left-hand side of the periodic table quite easily lose their few outer electrons giving them a high reactivity in forming positive ions eg groups 1 and 2 readily lose 1 or 2 electrons respectively to give ions with a full electronically stable outer shell like noble gases.
• These outer electrons are shielded by other inner electron shells and farthest from the nucleus and so are less strongly held and need less energy to ionise to give positive ions (cations) with a full outer shell of electrons.
• Its energetically easier for a metal to lose a few electrons to form a stable positive ion than gain many electrons to form unstable negative ions.
• Metals do NOT form negative ions - it is so energetically unfavourable for group 1 to 3 metals to gain 7 to 5 electrons and form a negative ion with a complete outer shell like a noble gas arrangement.
• The outer electrons of non-metals tend to be more strongly held than metals.
• Metals tend to get more reactive down a group of the periodic table (opposite of non-metals).

Reactivity of Metals Notes and Metal Extraction Notes

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2b. Typical Properties of Non–metallic Elements

Note these are TYPICAL characteristics of NON–METALS, but there are always exceptions!

Non-metals are found towards the right and top of the periodic table

(to the right of the diagonal  zig-zag line plus hydrogen and helium)

Physical properties of non–metals

• Non-metallic elements usually have low melting points and boiling points and so can be gases, liquids or solids at room temperature (exceptions like boron, silicon, and carbon as diamond or graphite - these have very stable giant covalent structures).

A tricky topic, only the basic idea should be dealt with at KS3/GCSE/IGCSE level.

There isn't a clear-cut diagonal band of semi-metals, but they exist!

Note in this table, I've included the old and more advanced numbering system for the groups (3 to 6 = 13 to 16)

3. Links to three selected Data–Graphs of selected physical properties of elements

• Group 1 elements, on the left of the periodic table, and at the start of a period, have one outer electron and so you would expect them to be very reactive metals and readily lose the outer electron to form a singly charged positive ion.

Group 3 elements

• Group 3 contains the metal Aluminium
  • See GCSE/IGCSE metal extraction notes
• and comparison with transition metals
• and titanium/steel/alloy uses GCSE notes.

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Group 4 elements

• Group 4 contains the non–metals carbon and silicon – which forms lots of compounds with hydrogen formed in oil (see GCSE Oil Products notes).
  • The structure of the allotropes of carbon and the structure and properties of silicon dioxide–silica–SiO₂ (GCSE notes on diamond, graphite and silica) are important.
• Advanced A Level Chemistry Notes on Group 4/14 Introduction – Carbon & Silicon – semi–metals e.g. Ge

Group 5 elements

• Group 5 contains the non–metal nitrogen and phosphorus – important element in natural and manmade artificial fertilisers.
  • See IGCSE/GCSE notes on ammonia and nitric acid
• Nitrogen forms 79% (⁴/₅th's) of air.

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Group 6 elements

• Group 6 are a Group of  non–metallic elements, the first 2 are O oxygen and below it S sulphur.
• They have 6 outer electrons and readily gain 2 electrons to form an X^2– ion in ionic compounds
  • e.g. they form ionic compounds with metallic elements e.g. magnesium oxide MgO and sodium sulphide Na₂S,
    • or written ionically: Mg²⁺O^2– and (Na⁺)₂S^2–.
• They form covalent small molecule compounds with other non–metallic elements e.g. H₂O or CS₂.
• The top element in the group is oxygen, a most important element.
  • Made by green plants in photosynthesis.
  • Consumed in the reverse process of respiration.
  • Pure oxygen is obtained from the fractional distillation of liquefied air, though for many industrial process, the 21% in air is quite adequate to use directly (fractional distillation is explained on another page, oxygen has a higher boiling point than nitrogen).
  • Oxygen is used in:
    • oxy–acetylene burners to produce a much hotter and intense flame for 'cutting' and welding metal,
    • oxygen 'tents' in hospitals for respiratory problems,
    • oxidant gas for burning rocket fuel.
• Oxygen combines with most other elements to form oxides of varying physical chemical character.
  • On the left and middle of the Periodic Table are the basic metal oxides which react with acids to form salts e.g. Na₂O, MgO, CuO etc. These metal oxides tend to be ionic in bonding character with high melting points. The Group 1 Alkali Metals, and to a less extent, Group 2 oxides, dissolve in water to form alkali solutions. All of them react with , and neutralise acids to form salts.
  • As you move left to right the oxides become less basic and more acidic.
  • So on the right you have the acidic oxides of the non–metals  CO₂, P₂O₅, SO₂, SO₃ etc. These tend to be covalent in bonding character with low melting/boiling points. Those of sulphur and phosphorus are very soluble in water to give acidic solutions which can be neutralised by alkalis to form salts.
  • These oxides are another example of the change from metallic element to non–metallic element chemical behaviour from left to right across the Periodic Table.
  • BUT life is never that simple in chemistry!:
    • Some oxides react with both acids and alkalis and are called amphoteric oxides. They are usually relatively insoluble and have little effect on indicators. An example is aluminium oxide dissolves in acids to form 'normal' aluminium salts like the chloride, sulphate and nitrate. However, it also dissolves in strong alkali's like sodium hydroxide solution to form 'aluminate' salts. This could be considered as 'intermediate' basic–acidic character in the Periodic Table.
    • Some oxides are neutral, tend to be of low solubility in water and have no effect on litmus, and do not react with acids or alkalis.  e.g. CO carbon monoxide (note that CO₂ carbon dioxide is weakly acidic) and NO nitrogen monoxide (note that NO₂ nitrogen dioxide is strongly acidic in water). There is no way of simply predicting this kind of behaviour from periodic table patterns!
• Sulphur is an important element used in the GCSE/IGCSE notes on the manufacture of sulphuric acid.
  • Sulphur or its compounds in oil burn to form the acidic polluting gas sulphur dioxide, one of the causes of acid rain (see Oil Product Notes).
• Advanced A Level Chemistry Notes on Group 6/16 Introduction – Oxygen & Sulfur

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Group 7 The Halogen elements

The halogen atoms and diatomic molecules

• Group 7 elements are on the far right of the periodic table with 7 outer electrons (1 short of a noble gas structure) and so you would expect them to be very reactive non-metals and form singly charged negative ions.
• The Group 7 Halogens are coloured non–metals with low melting points and boiling points eg chlorine, bromine and iodine.
• They are brittle when solid e.g. iodine and poor conductors of heat and electricity when liquid or solid.

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Names and symbols of the naturally occurring elements from atomic number 1 to 92

1 Hydrogen H ,  2 Helium He ,  3 Lithium Li ,  4 Beryllium Be ,  5 Boron B ,  6 Carbon C ,  7 Nitrogen N  ,  8 Oxygen O ,  9 Fluorine F ,  10 Neon Ne ,  11 Sodium Na ,  12 Magnesium Mg ,  13 Aluminium Al ,  14 Silicon Si ,  15 Phosphorus P ,  16 Sulphur S ,  17 Chlorine Cl ,  18 Argon Ar ,  19 Potassium K ,  20 Calcium Ca ,  21 Scandium Sc ,  22 Titanium Ti ,  23 Vanadium V ,  24 Chromium Cr ,  25 Manganese Mn ,  26 Iron Fe ,  27 Cobalt Co ,  28 Nickel Ni ,  29 Copper Cu  ,  30 Zinc Zn ,  31 Gallium Ga ,  32 Germanium Ge ,  33 Arsenic As ,  34 Selenium Se ,  35 Bromine Br ,  36 Krypton Kr ,  37 Rubidium Rb ,  38 Strontium Sr ,  39 Yttrium Y ,  40 Zirconium Zr ,  41 Niobium Nb  ,  42 Molybdenum Mo  ,  43 Technetium Tc ,  44 Ruthenium Ru ,  45 Rhodium Rh ,  46 Palladium Pd ,  47 Silver Ag ,  48 Cadmium Cd ,  49 Indium In ,  50 Tin Sn ,  51 ,  Antimony Sb ,  52 Tellurium Te ,  53 Iodine I ,  54 Xenon Xe ,  55 Caesium Cs ,  56 Barium Ba ,  57 Lanthanum La ,  58 Cerium Ce ,  59 Praseodymium Pr ,  60 Neodymium Nd ,  61 Promethium Pm ,  62 Samarium Sm ,  63 Europium Eu ,  64 Gadolinium Gd ,  65 Terbium Tb ,  66 Dysprosium Dy ,  67 Holmium Ho ,  68 Erbium Er ,  69 Thulium Tm ,  70 Ytterbium Yb ,  71 Lutetium Lu ,  72 Hafnium Hf ,  73 Tantalum Ta ,  74 Tungsten W ,  75 Rhenium Re ,  76 Osmium Os ,  77 Iridium Ir ,  78 Platinum Pt ,  79 Gold Au ,  80 Mercury Hg ,  81 Thallium Tl ,  82 Lead Pb ,  83 Bismuth Bi ,  84 Polonium Po 85 Astatine At ,  86 Radon Rn ,  87 Francium Fr ,  88 Radium Ra ,  89 Actinium Ac ,  90 Thorium Th ,  91 Protactinium Pa ,  92 Uranium U

• The ultimate origin of all elements is the nuclear reactions that go on when stars are formed from inter–stellar dust and gas in huge combined masses due to gravity, and then 'chunks' of a star spin off and cool down to form planets.

  • All the elements from atomic numbers 1–92 (H–U) naturally occur on Earth, though some are very unstable–radioactive and decay to form more nuclear stable elements.

  • For examples of nuclear fusion to produce heavier elements see

  • Nuclear fusion reactions and the formation of 'heavy elements'

  • From main sequence stars, like our own sun, converting hydrogen to helium in nuclear fusion, and also in red giants the elements from lithium (₃Li) to iron (₂₆Fe) are formed from fusing heavier nuclei, BUT ...

  • Elements heavier than iron, (cobalt to uranium, atomic numbers 27 to 92), are only formed in a supernova explosion of a red supergiant in a truly 'cosmic' scale explosion, where the temperatures are much higher than the 15 million degrees of the Sun!

  • The positive nuclei of heavier elements need enormous kinetic energies to overcome the massive repulsion forces between the positively charged nuclei, and fuse to make an even larger nucleus.

• Everything around you is made up of the elements of the periodic table, BUT most are chemically combined with other elements in the form of many naturally occurring compounds e.g.

  • hydrogen and oxygen in water, sodium and chlorine in sodium chloride ('common salt'), iron, oxygen and carbon as iron carbonate, carbon and oxygen as carbon dioxide etc. etc.!

• Therefore, most elements can only be obtained by some kind of chemical process to separate or extract an element from a compound e.g.

  • Less reactive metals are obtained by reduction of their oxides with carbon and more reactive metals are extracted by electrolysis of their chlorides or oxides (see GCSE notes on Metal Extraction)

  • Non–metals are obtained by a variety of means e.g. chlorine is obtained by electrolysis of sodium chloride solution (see Group 7 The Halogens GCSE Notes).

• However some elements never occur as compounds or they occur in their elemental form as well as in compounds e.g.

  • The Group 0 Noble Gases are so unreactive they are only present in the atmosphere as individual atoms.

    • Since air is a mixture, these gases are separated from air by a physical method of separation by distillation of liquefied air.

    • The elements oxygen and nitrogen are obtained from air at the same time, which is far more convenient than trying to get them from compounds like oxides and nitrates etc.

• Gold/platinum is are the least reactive metals and are often found 'native' as the yellow/silver elemental metal.

• Common metals like iron/aluminium are found as oxides and are extracted by carbon reduction/electrolysis.

• Relatively unreactive metals like copper and silver, can also be found in their elemental form in mineral deposits as well but usually in metal ores containing compounds like copper carbonate, copper sulphide and silver sulphide.

• The non–metal sulphur is found combined with oxygen and a metal in compounds known as sulfides, but it can occur as relatively pure sulphur in yellow mineral beds of the element.