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Doc Brown's GCSE/IGCSE O Level KS4 science-CHEMISTRY Revision Notes


An INTRODUCTION and OVERVIEW of the PERIODIC TABLE (modern and historic)

What is the Periodic Table? What is a 'Group'? What is a 'Period'? The basic structure of the Periodic Table is described and explained. The position of the elements in the periodic table is related to their electronic structure. A brief history of the development of the concept of the Periodic Table is also given. The typical physical and chemical characteristics of metals, semi-metals and non-metals are discussed in the context of their position in the periodic table. There are links to more detailed notes for each group or series of elements and at the end links to quizzes and worksheets based on the Periodic Table.

 See also Advanced Level Periodic Table Notes for A level students

Index of contents of this page for GCSE/IGCSE/O Level students

1. Summary of the structure of the Periodic Table

1a. The basic structure of the Periodic Table

1b. Electronic structure and the Periodic Table

1c. Electronic structure and the arrangement of elements in the Periodic Table

1d. More on patterns in the Periodic Table e.g. trends in physical and chemical properties

1e. Valency and formula patterns in the Periodic Table

2. Comparing Physical and Chemical Properties of Elements

2a. Typical Properties of Metallic Elements

2b. Typical Properties of Non-metallic Elements

2c. The Properties of Semi-metals or Metalloids

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

4. A brief Summary of some Groups & Series of elements of the Periodic Table with links to detailed notes

5. Snippets of the past and the continuing history of the Periodic Table

5a. Early classification by Antoine Lavoisier of 1789

5b. The 1829 work of Johann Döbereiner

5c. The work of John Newlands 1864

5d. Dmitri Mendeleev's Periodic Table of 1869

5e. The full modern version of the Periodic Table

6. Where do we get the elements from?

KEYWORDS-phrases for this page: Allotropes * Electron arrangements and the Periodic Table * Gp 1 Alkali Metals * Gp 7/17 Halogens * history of Periodic Table * hydrogen * Group * Gp 0/18 Noble Gases * Period * properties of metals * properties of non-metals * metalloids/semi-metals * structure of Periodic Table * Transition MetalsValency


1. Summary of the structure of the Periodic Table

1a. The basic structure of the Periodic Table

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

The thick diagonal zig-zag black line shows the main division between metals on the left of the periodic table and non-metals on the right of the periodic table. However, note that the metallic elements and non-metallic elements adjacent to this zig-zag line can show a 'mixture' of properties.

See the notes 1. to 4. in the full Periodic Table at the end of this page.

  • The idea of the Periodic Table is to arrange the elements in a way that enables chemists to understand patterns in the properties of elements,
    • AND be able to make predictions about the properties of elements and their compounds,
      • but some reminders first.
  • An ATOM is the smallest particle of a substance which can have its own characteristic properties, BUT atoms are built up of even more fundamental sub-atomic particles - the electron, proton and neutron and the structure of an atom ultimately determines its properties.
  • An ELEMENT is a pure substance made up of only one type of atom, 92 of the elements in the Periodic Table (part of which is shown above) naturally occur, from hydrogen H (element 1) to uranium U (element 92).
  • Note that each element has a unique symbol which is a single capital letter like H or U or a capital letter + small letter e.g. cobalt Co, chlorine Cl or sodium Na.
  • The majority of elements are readily divided into two types with common characteristic physical and chemical properties.
  • The elements are laid out in order of Atomic (proton) Number* (*see atomic structure page).
    • 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 0.
      • Likewise, in terms of physical and chemical properties, potassium is clearly an Alkali Metal in Group 1 of the Periodic Table.
  • Many of the similarities and differences in the properties of elements can be explained by the electronic structure of the atoms.
    • Electron configuration = electron arrangement in shells or energy levels, so watch out the varying phrases used!
  • The idea of the Periodic Table is to arrange the elements in a way that enables chemists to understand patterns in the properties of elements.
  • The main structural features of the periodic table are ...
    •  to produce columns of similar elements called Groups.
      • They are usually similar chemically and physically BUT there are often important trends in physical properties and chemical reactivity up/down a group.
      • They are similar elements because they have the same outer electron structure - same number of outer electrons.
    • When the vertical columns of elements are lined up next to each other in an appropriate manner, the resulting complete horizontal rows are called Periods and usually consist of a range of elements of different character.
      • There are important trends from left to right across a period e.g. the most important overall change is from metallic ==> non-metallic element character.
      • Certain 'horizontal blocks' of elements within a period, which have specific chemical features in common, may be known as a particular block or series e.g. from 21Sc to 30Zn are called the 1st Transition Metal Series within period 4.
    • The ideas of Group and Period are totally connected with electron structure.
      • So reading on below, try to connect electron arrangements with group number, period number and the chemical properties of the elements e.g. the formula of compounds and chemical reactions.).



1b. Electronic structure and the Periodic Table

  • Which electron arrangements are stable and which are not?
    • The maximum electrons allowed in the shells or electronic energy levels up to atomic number 20 are:

      • 1st shell 2, 2nd shell 8, 3rd shell 8, the 19th and 20th electrons go into shell 4 (this represents limit GCSE students need to know about electron arrangements - details on Atomic Structure page.

      • After element 20 Calcium, the rule changes and the 3rd shell can hold upto 18 electrons, but this knowledge is only required for advanced level students (Advanced Level Notes on Electron Configuration).

    • When an atom has its outer level full to the maximum number of electrons allowed, the atom is particularly stable electronically and very unreactive. [2],[2,8] and [2,8,8] etc. are known as the 'stable Noble Gas arrangements',

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

      • These 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 the chemistry of an element is about what the outer electrons can do, or can't, as the case maybe.

      • 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!

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

        • Group 1 metals cannot lose two electrons to form an e.g. Na2+ ion because too much energy is needed to get such a strongly held inner shell electron involved in bonding, so you can't form NaCl2.

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

      • These are the Group 7/17 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/share electrons to 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).

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

        • The chlorine cannot accept another electron to form a Cl2- ion, because its electron structure would not be that of a stable noble gas arrangement.

    • The electron arrangements of the first 20 elements are shown below.

      • This basically uses the electron arrangements of the elements to construct the modern periodic table e.g.

        • elements of the same group have the same number of outer electrons, giving them very similar chemical properties,

        • when you move down to the next period, you have added an extra shell of electrons, so the atomic radius gets bigger, reactivity patterns emerge etc.

          • and more on this in section 1c.

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

(c) doc b

Picture examples of these electron arrangements (atomic number)

Filling 1st shell, electron level 1 (c) doc b (c) doc b2 elements only in Period 1

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

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

The first 2 elements of the 4th shell (c) doc b (c) doc b to Kr [], start of Period 4

Detailed Advanced Level Chemistry Notes on the electronic structure of atoms



1c. Electronic structure and the arrangement of elements in the Periodic Table

(c) doc b

  • All substances are made up of one or more of the different types of atoms we call elements and the elements identity is solely determined by the atomic number of protons.
  • Hydrogen, 1, H, the simplest element atom, does not readily fit into any group.
  • A Group is a vertical column of like elements e.g. Group 1 The Alkali Metals (for full GCSE notes on Li, Na, K etc.), Group 7/17 The Halogens (for full GCSE notes on F, Cl, Br, I etc.) and Group 0/8/18 The Noble Gases (for full GCSE notes on He, Ne, Ar etc.).
  • For the first twenty elements, part from hydrogen (doesn't really fit in any group), helium (*), the Group number equals the number of electrons in the outer shell and the number of electron shells used equals the Period number, e.g. chlorine's electron arrangement is 2.8.7, the second element down in Group 7 on period 3.
    • So after helium, elements in the same group have the same outer electron structure.
    • Beyond element 20, apart from the Transition Metals (not shown above), the
    • (*) Although helium can't have 8 outer electrons like the rest of Group 0, its outer shell of 2 electrons is complete according to the electron shell rules, just like neon and argon etc. and therefore has the same chemical (very unreactive) and physical properties (monatomic gas molecules).
      • The Group 0 Noble Gases are never called group 8, but after helium they all have eight out electrons to fill the highest occupied electron level (outer shell).
  • The elements in a group tend to have similar physical and chemical properties because of their similar outer shell electron structure.
  • A Period is a horizontal row of elements with a variety of properties, changing from very metallic elements on the left to non-metallic elements on the  right. A period starts when the next electron goes into the next available main energy level or shell (Group 1 alkali Metals). The period ends when the main energy level is full ie reached the Group 0 Noble Gases.
  • All the elements on the same period use the same number of principal electron shells, and this equals the period number (e.g. sodium's electron arrangement 2,8,1, the first element in Period 3).
  • The first element in a period is when the next electron goes into the next available electron shell or energy level (i.e. 1 electron in the outer shell, after H it is the Group 1 Alkali Metals like sodium 2.8.1).
  • The last element in a period is when the outer shell is full resulting in a very unreactive element, the Group 0 Noble Gases e.g. argon 2.8.8. The next electron for the next element goes into the next highest level (shell) available, and so starts the next period with a group 1 element again, periodicity - a very similar element every so often - but governed by the electron rules.
  • So in terms of electrons ....
    • Period 1 is elements 1-2,  H (1) to He (2)
    • Period 2 is elements 3-10, Li (2,1) to Ne (2,8)
    • Period 3 is elements 11-18, Na (2.8.1) to Ar (2.8.8)
    • Period 4 is elements 19-36, starts with K (2,8,8,1) and Ca (2,8,8,2) and finishes with the Noble Gas Kr (2,8,18,8).
    • Note that the number of shells containing electrons is equal to the period number.
  • The similarities (e.g. same Group) or differences (e.g. across a period) of the properties of the elements can be explained by the electronic structure of the atoms.
  • From Period 4 onwards the length of a period significantly increases because it includes horizontal series of similar metals with their own characteristic physical and chemical properties e.g. Transition Metals (detailed GCSE notes on Fe, Cr, Cu etc.)

Advanced Level Electron configurations, spectroscopy, hydrogen spectrum, ionisation energies



1d. More on patterns in the Periodic Table

  • More than three-quarters of the 109 known elements are metals (elements naturally occur up to uranium 92, 93-109 are 'man-made' elements from the experiments of nuclear physicists.
    • This work will continues as heavier and heavier elements are likely to be made in nuclear reactions. They will be all metals and radioactive. BUT one theory suggests that 'super-heavy' elements of about atomic number 150? may be in a nuclear stability region and would prove most interesting to study. Chemists are trying to predict their properties now!, so it may have started with Mendeleev but it ain't finished yet!
  • Only about nineteen are definitely are non-metal but about seven more are semi-metals of mixed metallic and non-metallic physical and chemical character.
  • The metals in the periodic table are mainly found in the left hand columns (Groups 1 and 2) and in the central blocks of the transition elements.
  • There is a 'rough' diagonal division between the two principal types of element zig-zagging from B-Al in group 3 to Te-Po in Group 6 (see semi-metals section 2c.).
  • The elements in this 'band' are sometimes referred to as 'semi-metals' or 'metalloids' because of their 'mixture' of metallic and non-metallic physical or chemical character e.g. the semi-conductor silicon in group 4.
  • The 'electronic origin' of trends down a group.
    • The further the negative electrons are from the positive nucleus the less strongly they are held.
    • As you go down to the next element in a group, an extra shell of electrons is added.
    • This has several effects, e.g.
      • the atomic radius increases with the extra 'inner' shell of electrons,
      • each added shell of electrons means the outer electrons are further away from the nucleus and less strongly held,
      • the extra shell of electrons actually has a 'shielding' effect on the outer shell electrons and this further decreases the attractive force of the nucleus on the outer electrons,
      • this increase in atom size and decreased attraction of the nucleus for the outer electrons has a powerful effect on the element's reactivity i.e. it helps explain the reactivity trends, particularly of groups 1, 2 and 7 (briefly discussed below)
  • There tends to be gradual changes in physical and chemical properties down a group e.g.
    • (c) doc b==>(c) doc b==>(c) doc b==> down group 1
    • ==> ==> ==> down group 2
    • 2.2 lithium ==> 2.8.2 magnesium ==> calcium ==> down group 2
    • Down Group 1 (Alkali Metals) and Group 2 the metals get more reactive as the outer electrons are more readily lost (more details in the Group 1 Alkali Metals Notes).
    • 2.7fluorine ==> 2.8.7(c) doc bchlorine ==> down group 7
    • Down Group 7 (Halogens) the non-metals get less reactive, their colour gets darker, their melting/boiling points increase and they get more reactive up the group because the nucleus of the smaller atom attracts electrons more strongly (more details in the Group 7 Halogens Notes). You can also relate the increase in size of the atoms (as diatomic molecules F2, Cl2, Br2 etc.) to the increase in intermolecular forces causing the increase in melting/boiling points.
    • Down Group 4 you start with a definite non-metal carbon, and end up at the bottom with a the definite metal lead, so there are quite significant changes in both physical and chemical character.
  • There tends to be major changes in physical and chemical properties across a period e.g.
    • Period 2 starts with a solid low melting reactive metal lithium, in the middle there are the high melting and rather unreactive non-metals boron and carbon, next to the end is the very highly reactive non-metal gas fluorine, and the period finishes with the very unreactive gas neon. Very complicated pattern!
    • Period 4 starts with a solid low melting very reactive metal rubidium, after calcium there are ten transition metals with a wide variety of chemistry, followed by the metallic gallium, semi-metal germanium and more non-metallic arsenic/selenium, next to the end is the very reactive non-metal liquid bromine, and the period finishes with the very unreactive gas krypton. Even more complicated pattern!
    • From left to right across a period the bonding in chlorides or oxides changes from ionic (with metals e.g. Na+Cl-, Mg2+O2- to covalent (with non-metals e.g. HCl, SO2).
    • From left to right across a period the oxides change from alkaline/basic (with metals e.g. Na2O) to acidic (with non-metals e.g. SO2). More on this in Group 6/16 Oxygen and oxides.
    • Note on electron arrangements:
      • Except for boron, most non-metals have at least four electrons in the out shell.
      • Except for the noble gases, the more electrons in the outer shell the more non-metallic and the more reactive the element. The most reactive non-metals only need to share/gain one or two electrons.
      • The most reactive metals only have 1 or 2 electrons in the outer shell which tend to be easily lost to form the metal ion in reaction.
      • The most reactive metals have a low number of outer valency shell electrons (<= 3).
      • The very reactive non-metals have 5 to 7 outer valency shell electrons.
      • Elements in the 'middle' of the Periodic Table e.g. Group 4 with 4 outer electrons, show non-metal e.g. carbon or metal chemical character e.g. lead and non are very reactive elements.
      • The Noble Gas elements have full, very stable, outer valency shells.



1e. Valency and formula patterns in the Periodic Table

  • The valency, or, combining power of an element is related to the elements position in the Periodic Table.
    • For Groups 1 to 7, the group number gives the maximum valency possible and equals the number of outer electrons (well, nearly always!).
    • For many compounds, this rule works fine: e.g. for sodium valency 1 and oxygen valency 2 etc., you can deduce the following formulae for valencies of 1 to 7 across the periods for Group 1 to 7 compounds (at least up to a point!) e.g. for period 3
      • chlorides: NaCl, MgCl2, AlCl3, SiCl4, PCl5, SCl6, then Cl itself and Ar can't combine with other elements.
      • oxides: Na2O, MgO, Al2O3, SiO2, P2O5, SO3, Cl2O7 and Ar can't combine with other elements.
        • Working out formulae is an advanced idea and methods of how to do it are described on the equations and formulae page.
  • BUT things are not always so simple!
    • Na to Si no problem! great! In fact, apart from N, O, F (which have valency restrictions NOT for GCSE though!) you can usually make a reasonable prediction of the maximum valency compound for all of the elements in Groups 1 to 7 by using the group number.
    • However there are lots of other compounds where the element's valency is less than its group number and there is even a pattern of decreasing valency from Group 4 to Group 7 (as well as the pattern of increasing valency mentioned above, see table below and the decreasing pattern for hydrides which is important for GCSE level).
    • e.g. in Group 4, C forms CO (nasty!) but CO2 is more stable, Pb forms PbCl2 which is much more stable than PbCl4.
    • Xenon forms XeF8 and XO4 using its maximum valency of 8! and that got somebody a Nobel Prize in Chemistry!
    • Tabulated below are some formulae you are likely to come across in your GCSE or equivalent course in bold, but others you are unlikely to come across are included because they fit in with general formula patterns.
    • The valency of hydrogen is 1 (hydrides), oxygen 2 (most oxides) and chlorine is usually 1 (most chlorides).
    • The expected-theoretical formulae for the hydride, chloride and oxide for element X of valency 1 to 7 are given below and examples of all these formulae can be found in the Period 2-4 table further down.
      • Some combinations don't exist, but a clear periodic pattern of formulae emerges in the next two tables below.
      • At GCSE/IGCSE level you will only come across valencies of 1 to 4 and the examples in bold are like those you will become familiar with through your GCSE/IGCSE science-chemistry course.
    • max. valency of X = group number 1 2 3 4 5 6 7
      formula of hydride XH

      eg NaH


      eg MgH2


      eg NH3


      eg CH4

      - - -
      formula of chloride XCl

      eg KCl


      eg MgCl2


      eg AlCl3


      eg SiCl4


      eg PCl5

      formula of oxide X2O

      eg Li2O


      eg MgO


      eg Al2O3


      eg CO2







      • Hydrides are compounds formed by combining an element with the element hydrogen.
      • Chlorides are compounds formed by combining an element with the element chlorine.
      • Oxides are compounds formed by combining an element with the element oxygen.
    • Examples of how to work out formula from valencies or ionic charge is on the Elements, Compounds and Mixtures page and the structure, bonding and properties of many of these substances is discussed elsewhere.
    • topIn the table below you can see how the formula change horizontally from left to right with change in valency AND the vertical connection where elements in the same group form compounds of the same formula.
element, hydride, chloride and oxide Group 1 Alkali Metals Group 2 Group 3 Group 4 Group 5 Group 6 Group 7 The Halogens Group 0 Noble Gases
valency 1 2 3 4 3 or 5 2, 4, 6 often 1, can be 3,4,5,7 - complicated -
Period 2 Li
































Period 3 Na


















PCl3, PCl5

P2O3, P2O5




SO2, SO3









Advanced Level Chemistry Notes on Period 2 survey Li to Ne

Advanced Level Chemistry Notes on Period 3 survey Na to Ar

2. Comparing Physical and Chemical Properties of Elements



2a. Typical Properties of Metallic Elements

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

Physical properties of metals

  • Usually high melting points and boiling points so all solid bar one (exceptions like mercury the only liquid metal at room temperature and the Alkali Metals have untypical low melting points).
  • Often very good conductors of heat and electricity. This is due to the mobility of the free moving electrons in the structure of a metal.
  • Most have a high density  (exceptions like the Alkali Metals have untypical low densities, the first three Li, Na and K float on water before the 'fizzing'!).
  • Their appearance is always 'shiny' (usually silvery, except for copper and gold)
  • Usually quite strong materials (exceptions like the Alkali Metals which are atypically very soft, and metals like lead are relatively soft too)
  • They are easily beaten into shape (malleable) or drawn into wire (ductile) of varying strength, from very weak sodium to very strong iron or titanium).
  • Solids sonorous, they ring or resonate to produce a note when struck.

Chemical Properties of metals

  • They tend to form basic oxides that react with acids to form salts (if the oxide is soluble in water it forms an alkali of pH > 7, universal indicator blue or violet).
  • Most metals react with acids to form a salt and hydrogen. (see metal reactions: reactivity and metal-acid reaction equations with answers).
  • Metals readily form positive ions in compounds by losing electrons e.g. 
    • sodium Na - e- ==> Na+, magnesium Mg - 2e- ==> Mg2+ or aluminium Al - 3e- => Al3+ 
  • Their oxides and chlorides are usually ionic* in terms of chemical bonding.  e.g.
    • magnesium oxide, MgO or Mg2+O2-, sodium oxide Na2O or (Na+)2O2- ,
      • and aluminium oxide Al2O3 or (Al3+)2(O2-)3 
      • * At least at GCSE level, but there are some chloride exceptions at Advanced level such as FeCl3 and AlCl3.

(c) doc b Reactivity of Metals Notes and (c) doc b Metal Extraction Notes



2b. Typical Properties of Non-metallic Elements

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

Physical properties of non-metals

  • They usually have low melting points and boiling points and so can be gases, liquids or solids (exceptions like silicon, and carbon as diamond or graphite).
  • Usually poor conductors of heat and electricity (exceptions like carbon in the form of graphite).
  • Non-metals generally have a low density.
  • The appearance can be quite varied but tend to be dull if solid.
  • Often weak materials e.g. soft or brittle solids (exceptions like silicon, and carbon as diamond, which are very hard and strong)
  • When solid they are not easily beaten into shape or drawn into wire, the solids tend to be too brittle.
  • Solid non-metals are not usually sonorous, e.g. they do not usually resonate or ring with sound, like when a piece of metal is struck.

Chemical properties of non-metals

  • They tend form acidic oxides when burned in air or oxygen, these react with alkalis to form salts, if soluble in water they form acid solutions of pH <7, universal indicator yellow-orange-red
  • Non-metals do not usually react with acids e.g. to produce a salt and hydrogen like most metals do.
  • Non-metals readily form negative ions in compounds by losing electrons e.g. 
    • chlorine ==> chloride: Cl2 + 2e- ==> 2Cl- (more simply Cl + e- ==> Cl- typical of Group 7 Halogens)
    • oxygen ==> oxide: O2 + 4e- ==> 2O2- (more simply O + 2e- ==> O2- typical of Group 6 elements)
  • The oxides and chlorides, when combined with other non-metals are always covalent in terms of chemical bonding.
    • e.g. water H2O(l), methane CH4(g), sulphur dioxide SO2(g) and hydrogen chloride HCl(g)  
  • The oxides and chlorides, when combined with metals tend to be ionic in terms of chemical bonding e.g.
    • sodium chloride, NaCl or Na+Cl- , magnesium chloride MgCl2 or Mg2+(Cl-)2 ,
      • and magnesium oxide, MgO or Mg2+O2-



2c. The Properties of Semi-metals or Metalloids

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

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









BASIC IDEA: A narrow diagonal band of elements can show both metallic and non-metallic physical or chemical properties and are referred to as 'semi-metals' or 'metalloids'. Although most tend to be nearer being a metal or a non-metal, they do represent the point elements change from metal to non-metal as you move from left to right across the Periodic Table BUT please read the notes below carefully!
B C N O To me boron, B, is clearly a non-metal, showing no real metallic character and I'm not sure why it is sometimes shown as a semi-metal on some periodic tables? and is very different in character to metallic aluminium below it in the same group. Boron's oxide is acidic only, and the solid element consists of a non-conducting giant covalent structure, both classic non-metallic properties. Carbon, C, is also clearly a non-metal, its oxide is acidic and in the form of diamond, it is a non-electrical conducting 3D giant covalent structure. However, in the form of graphite, it has a layered 2D giant covalent structure that does allow electricity to conduct through the layers. (more details)
Al Si P S Physically and chemically aluminium, Al, is very much a metal, but the oxide/hydroxide reacts with both acids (metallic) and alkalis (acidic) to form salts showing dual character. Silicon is mainly non-metallic character e.g. the oxide is acidic but, although the solid element has a giant covalent structure, it shows slight electrical conducting properties (semi-conductor), especially when doped with other elements and used in computer chip technology. To me, neither are true semi-metals.
Ga Ge As Se Germanium, Ge, is considered as a true semi-metal (metalloid). Like silicon, germanium is a semi-conductor and used in electronic technology. Its oxide/hydroxide react with both acids/alkalis showing dual metal/non-metal character. Arsenic, As, is also a true metalloid with oxides/hydroxides that react both with acids/ and alkalis to form salts and the element exists in two allotropic* crystalline forms. One form is less dense, non-conducting and covalent in structure (non-metal) and the other is more dense and weakly electrical conducting (metallic) and used in transistors. Selenium, Se, is also a semi-conductor with metallic and non-metallic properties and is used in photo-electric cells (solar cells) and xerography (photocopying). (*Allotropes are different physical forms of the same element in the same physical state.)
In Sn Sb Te Arsenic, As, (like antimony in the same group), is also a true semi-metal (metalloid) with oxides/hydroxides that react both with acids/ and alkalis to form salts and the element exists in two allotropic* crystalline forms (non-metallic and metallic). Tellurium, Te, is also a semi-conductor with metallic and non-metallic properties. Both As and Te are used in electronic devices.



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

Links to the first 'experimental' editions of these new web pages are below. They are of more use to Advanced Level students studying 'Periodicity', but they are a source of useful data. There are also summaries of data for Group 1 Alkali Metals, Group 2 Alkaline Earth Metals, Group 7 Halogens, Group 0 Noble Gases and the 1st series of Transition Metals.

Elements 1-20 covering Periods 1-3 and start of Period 4

Elements 1-38 covering Periods 1-4 and start of Period 5

Elements 1-96 covering Periods 1-6 and start of Period 7


4. A brief Summary of some Groups & Series of elements of the Periodic Table

with LINKS to more detailed GCSE/IGCSE/O Level notes where necessary

only the briefest summary here, but enough to connect with ideas of how the periodic table developed and what the periodic table means as a system of classifying and arranging the elements



Group 1 Alkali Metals

  • The very reactive Group 1 The Alkali Metals  [detailed GCSE/IGCSE notes] have low density (some float on water) eg lithium, sodium and potassium etc.
  • They readily react with non-metals to form ionic compounds
  • They combine with oxygen to form oxides that dissolve in water to form alkaline hydroxide solutions.
    • eg lithium oxide Li2O or (Li+)2O2-.
  • They combine with halogens like chlorine to form salt like halide compounds that are soluble in water to give ~neutral solutions
    • e.g. sodium chloride NaCl or Na+Cl-
  • These oxides and halides ionic compounds are colourless crystals or white solids, soluble in water to give colourless solutions (usually pH 7 for their salts, pH 13-14 for the oxides because the MOH alkali formed).
  • The metals react rapidly, maybe violently, with water to form alkaline hydroxides and hydrogen gas.
  • Alkali metal atoms have one outer electron, which is readily lost to form a stable single positive ion M+.
  • Down the group, the metals get more reactive, and the melting points and boiling points decrease.
  • Detailed advanced Level Chemistry Notes on Group 1 and Group 2 Metals



Group 2 Alkaline Earth Metals

  • Group 2 are the 2nd group of metals (sometimes called "Alkaline Earth Metals") eg magnesium and calcium.
  • They are not quite so reactive as the Alkali Metals for the same period.
  • They have two outer electrons and readily lose them to form the M2+ ion.
  • This ion occurs in the ionic compounds they readily form with non-metals like the Group 7 Halogens or oxygen and sulphur from Group 6 e.g. MgCl2 or CaO.
  • Detailed advanced Level Chemistry Notes on Group 1 and Group 2 Metals



Group 3 (Group 13 in modern notation)



Group 4 (Group 14 in modern notation)



Group 5 (Group 15 in modern notation)



Group 6 (Group 16 in modern notation)

  • Group 6/16 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 X2- 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+O2- and (Na+)2S2-.
  • 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 liquified 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 Level Chemistry Notes on Group 6/16 Introduction - Oxygen & Sulfur



Group 7 The Halogens (Group 17 in modern notation)

  • The Group 7 Halogens [detailed GCSE/IGCSE notes] 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.
  • Halogens exist as molecules of pairs of atoms, X2 (diatomic molecules), form ionic salts with metals e.g. KBr or MgCl2, but form covalent molecular compounds with other non-metallic elements e.g. HCl, CBr4.
  • The halide ions, X-, are formed by halogen atoms, with 7 outer electrons, gaining 1 electron to form a stable noble gas electron structure.
  • Down the group the melting points and boiling points increase and the reactivity decreases.
  • Sodium chloride is a very important raw material from which hydrogen, chlorine and sodium hydroxide can be manufactured by electrolysis.
  • These products have many uses and are important in the manufacture of other useful compounds ranging from bleaches, hydrochloric acid and plastics etc.
  • Advanced Level Chemistry Notes on Group 7/17 The Halogens



Group 0 The Noble Gases (Group 18 in modern notation)

  • The Group 0/18 Noble Gases [detailed GCSE/IGCSE notes] are colourless non-metals with very low melting and boiling points (they are all gases at room temperature).
  • They exist as individual atoms (NOT diatomic molecules) and are very unreactive chemically due to their very stable full outer shell electron arrangement, which is very stable.
  • Apart from helium (2 electrons), all the others have 8 outer electrons giving a stable outer octet of electrons that is difficult to change so they are most reluctant to combine with other elements, hence their 'inert' or 'unreactive' nature.
  • Helium has a very low density and so is used in balloons and airships.
  • Their lack of chemical reactivity makes them useful to provide an 'inert' atmosphere to prevent oxidation e.g. argon in filament bulbs and in arc welding.
  • Advanced Level Chemistry Notes on Group 0/18 The Noble Gases



The 1st Transition Metal Series (Scandium to zinc)

  • The ten horizontal elements Sc to Zn are called the 1st series of Transition Metal Elements [GCSE/IGCSE notes] e.g. iron and copper.
  • These elements in the central blocks of the periodic table are typical metals - good conductors of heat and electricity and can be bent or hammered into shape (malleable) and they can be drawn into wire (ductile).
  • However, compared to the group 1 Alkali Metals, they have higher melting points (except mercury - a liquid at room temperature); they are harder, tougher and stronger; they are much less reactive and so do not react (corrode) as quickly with oxygen or water.
  • These properties make them useful structural materials (e.g. steel) and were things need to be good conductors e.g. copper electrical wiring or steel radiators.
  • Most transition metals form coloured compounds (e.g. blue copper salt solutions) and are used in pottery glazes, stained glass and weathered copper roofs turn green!
  • Many transition metals e.g. iron and platinum are used as catalysts. Cast iron is hard and used as man-hole covers. Steel is an alloy* based on iron and used for car bodies.
  • *alloy means a metal mixed with at least one other element.
  • see also Metal Extraction (detailed GCSE notes) and more on metal uses on the Extra Industrial Chemistry - detailed GCSE notes - use index of sub-pages.
  • Detailed Advanced Level Chemistry Notes on the 3d-block of elements and Transition Metals


5. Snippets of the past and continuing history of the Periodic Table

A more advanced history for A level students

5a. The early classification of Antoine Lavoisier of 1789

Antoine Lavoisier's 1789 classification of substances into four 'element' groups

acid-making elements gas-like elements metallic elements earthy elements
sulphur light cobalt, mercury, tin lime (calcium oxide)
phosphorus caloric (heat) copper, nickel, iron,  magnesia (magnesium oxide)
charcoal (carbon) oxygen gold, lead, silver, zinc barytes (barium sulphate)
  azote (nitrogen) manganese, tungsten argilla (aluminium oxide)
  hydrogen platina (platinum) silex (silicon dioxide)
  • The understanding that an element as a unique atomic 'building block' which could not be split into simpler substances and a compound is a chemical combination of two or more elements were not at all understood at the time of Lavoisier.
    • Scientists had no idea that atoms were made of protons, neutrons and electrons until the early part of the 20th century.
  • However, Lavoisier was the first to define an element in the correct 'chemical sense' as a substance that could not be divided into simpler substances.
  • 'light' and 'caloric' (heat), were considered 'substances' and the last 'scientific' vestige of the elements of 'earth, fire, air and water' which had there conceptual origin in the Greek civilisation of 2300-2800 years ago.

  • However, Lavoisier was correct on a few things e.g. the elements sulphur, phosphorus and carbon and correctly described their oxides as acidic e.g. dissolved in water turned litmus turns red.

  • Many metallic elements, were correctly identified though I doubt if they were pure though!

  • What he described as the 'earthy elements' are of course compounds, a chemical combination of a metal plus oxygen or sulfur (both in case of barium).

  • He didn't have very high temperature smelting technology, or a reactive metal from electrolysis (came in about 1806 onwards)' to 'separate' the elements in some way e.g. he couldn't extract a reactive metal! In other words, at this time, the wrong 'classification' was due to a lack of chemical technology as much as lack of knowledge.


5b. The 1829 work of Johann Döbereiner

  • Johann Döbereiner noted that certain elements seemed to occur as 'triads' of similar elements e.g.
    • (i) lithium, sodium and potassium
    • (ii) calcium, strontium and barium
    • (iii) chlorine, bromine and iodine
  • Döbereiner was amongst the first scientists to recognise the 'group' idea of chemically very similar elements.

  • Three groups he 'recognised' were (i) Group 1 Alkali Metals, (ii) Group 2 Alkaline Earth Metals, (iii) Group 7 Halogens.


5c. The work of John Newlands 1864

Newlands' Octaves (his 'Periodic Table' of 1864)

H Li Ga B C N O
F Na Mg Al Si P S
Cl K Ca Cr Ti Mn Fe
Co, Ni Cu Zn Y In As Se
Br Rb Sr Ce, La Zr Di, Mo Ro, Ru
Pd Ag Cd U Sn Sb Te
I Cs Ba, V Ta W Nb Au
Pt, Ir Tl Pb Th Hg Bi Cs
  • Newlands recognised that every 7 elements, the 8th seemed to be very similar to the 1st of the previous 7 when laid out in a 'periodic' manner and he was one of the first scientist to derive a 'Periodic Table' from the available knowledge.

  • e.g. his 'table' consists of almost completely genuine elements (Di was a mix of two elements), classified roughly into groups of similar elements and a real recognition of 'periodicity'

  • Newlands also recognised that the 'groups' had more than 3 elements (not just 'triads'), and was correct to mix up metals and non-metals in same group e.g. in 5th column there is carbon, silicon, tin (Sn) from what we know call Group 4. However, indium is in group 3 but Ti, Zr have a valency of 4, like Group 4 elements and do form part of vertical column in what we know call the Transition Metal series

  • Other correct 'patterns' if  not precise are recognisable in terms of the modern Periodic Table e.g. half of column 2 is Group 1, half of column 3 is Group 2, half of column 5 is Group 4, half of column 6 is Group 5, half of column 7 is Group 6. If we put his column 1 as column 7, it would seem a better match of today!

  • Although none of his vertical column groups match completely but the basic pattern of the modern periodic table  was emerging. However column's 1 and 7 do seem to be particularly mixed up compared to the modern periodic table!

  • The main criticisms of Newland's work were ...

    • Some of his groups contained elements that were very different e.g. the halogens (Cl, Br, I) in column 1 are mixed with metals like Pt, Ir.

    • The above example also shows how mixed up he had metals and non-metals (though this is the case for e.g. Group 4 of the modern periodic table.

    • Newlands, unlike Mendeleev later, didn't predict the existence of 'missing elements' yet to be discovered, so his table was not only inaccurate (but going in the right direction) but had no predictive value.


5d. Dmitri Mendeleev's Periodic Table of 1869

Mendeleev Periodic Table 1869

  • Mendeleev laid out all the known elements in order of 'atomic weight' (what we now call relative atomic mass) except for several examples like tellurium (Te) and iodine (I) whose order he reversed because chemically they seemed to be in the wrong vertical column!

    • He arranged similar elements into vertical columns (we now call groups) and switched to the next row (we now call periods) when the next similar element appeared when laid out in atomic number (left to right and then down).

    • You can see immediately his periodic table is becoming recognisable in terms of what you see in your textbook or on the internet! For 1869, pretty smart thinking!

      • In Mendeleev's periodic table the Group 1 metals (Li, Na, K, Rb) and Group 2 metals (Be, Mg, Ca, Sr) are all in place in columns 1 and 2, three of group 3 (B, Al, In), three of group 4 (C, Si, Sn) in place, 4 of group 5 (N, P, As, Sb), 4 of group 6 (O, S, Se, Te), 4 of group 7 (F, Cl, Br, I), but his group 8 (where group 0 Noble Gases are) are metals from the first and second transition metals series.

    • BUT, you also see how he left gaps, he reasoned from the patterns of physical and chemical properties of similar elements of a 'group'  that there must be one or more missing!

      • Similar properties would involve physical properties like melting point, boiling point, density, appearance and chemical properties would include formula of compounds like oxides, chlorides etc.

    • There is still some apparent mixing of metals and non-metals, but Mendeleev recognised the idea of sub-groups, some of which turned out to be the transition metals e.g. titanium in a sub-group of column four (= group 4).

  • With an increased number of known elements, groups becoming more clearly defined, and he used a double column approach which is NOT incorrect, i.e. a sort of group xA and xB classification.

    • This is due to the 'insert' of transition metals, some of whom show chemical similarities to the vertical 'groups', but this is needed to be understood for GCSE or A level!

    • The 'insertion' of the rows of transition metals in the periodic table is due to the complex, but specific ways, electron energy levels are arranged - but this is for A Level students later!

  • However, his 'presentation' was sufficiently accurate, and Mendeleev was sufficiently confident to predict missing elements and their properties * e.g. germanium (which he called eka-silicon, below Si and above Sn in Group IV and Mendeleev is deservedly called the 'father of the modern Periodic Table'.

  • Even though Mendeleev had cracked the basic idea of the periodic table, he and others in the late 19th century, still had no real knowledge of atomic structure, atoms were still thought of as some kind of indivisible tiny particles and protons, neutrons and electrons were still to be discovered decades later.


5e. The full modern version of the Periodic Table

Pd metal groups horizontal blocks of Transition Metal Series (Periods 4 to 7) metal ==> non-metal groups
Gp1 Gp2 Gp3 Gp4 Gp5 Gp6 Gp7 Gp0

1H   Note: H does not readily fit into any group which are the vertical columns

2 3Li 4Be The full Modern Periodic Table of Elements

ZSymbol, z = atomic or proton number

5B 6C 7N 8O 9F 10Ne
3 11Na 12Mg 13Al 14Si 15P 16S 17Cl 18Ar
4 19K 20Ca 21Sc 22Ti 23V 24Cr 25Mn 26Fe 27Co 28Ni 29Cu 30Zn 31Ga 32Ge 33As 34Se 35Br 36Kr
5 37Rb 38Sr 39Y 40Zr 41Nb 42Mo 43Tc 44Ru 45Rh 46Pd 47Ag 48Cd 49In 50Sn 51Sb 52Te 53I 54Xe
6 55Cs 56Ba *57-71 72Hf 73Ta 74W 75Re 76Os 77Ir 78Pt 79Au 80Hg 81Tl 82Pb 83Bi 84Po 85At 86Rn
7 87Fr 88Ra *89-103 104Rf 105Db 106Sg 107Bh 108Hs 109Mt 110Ds 111Rg 112Cn 113? 114Fl 115? 116Lv 117? 118?
Gp 1 Alkali Metals

Gp 2 Alkaline Earth Metals

Gp 7/17 Halogens

Gp 0/18 Noble Gases

Take note of the four points on the right

*57La 58Ce 59Pr 60Nd 61Pm 62Sm 63Eu 64Gd 65Tb 66Dy 67Ho 68Er 69Tm 70Yb 71Lu
*89Ac 90Th 91Pa 92U 93Np 94Pu 95Am 96Cm 97Bk 98Cf 99Es 100Fm 101Md 102No 103Lr
  1. Using 0 to denote the Group number of the Noble Gases is historic i.e. when its valency was considered zero since no compounds were known. However, from 1961 stable compounds of xenon have been synthesised exhibiting up to the maximum possible valency of 8 e.g. in XeO4.

  2. Because of the horizontal series of elements e.g. like the Sc to Zn block (10 elements), Groups 3 to 7 & 0 can also be numbered as Groups 13 to 18 to fit in with the maximum number of vertical columns of elements in periods 4 and 5 (18 elements per period).

  3. This means that 21Sc to 30Zn can be now considered as the top element in the vertical Groups 3 to 12.

  4. I'm afraid this can make things confusing, but there it is, classification is still in progress! and GCSE students can, as far as I can judge, ignore points 2 and 3.

  5. Advanced Level Periodic Table Notes

  • With are knowledge of atomic structure, the modern Periodic Table is now laid out in order of atomic (proton) number and is directly linked to the electronic structure of elements.

    • Knowledge of the proton number and electronic structure sorted out the inconsistencies which Mendeleev had recognised purely from the physical and chemical properties of the elements known at the time (~1869).

  • We now know that due to the presence of, the relative atomic mass does go 'up/down' occasionally, BUT as Mendeleev recognised, chemically Te (tellurium) is like S (sulfur) and Se (selenium) etc. and I (iodine) is like Cl (chlorine) and Br (bromine) etc. and this is now backed up by modern knowledge of electron structure.

  • We know the electronic structure of elements and can understand sub-levels and the 'rules' in electron structure (see atomic structure page) e.g. 2 electrons in shell 1 (period 1, 2 elements H to He), 8 in shell 2 (period 2, 8 elements Li to Ne), there is a sub-level which allows an extra 10 elements (the transition metals) in period 4 (18 elements, K to Kr). this also explains the sorting out of Mendeleev's A and B double columns in a group (but that's for much more advanced chemistry!). The periods are complete now that we know about Noble Gases.

  • The use and function of the Periodic Table will never cease! Newly 'man-made' elements are being synthesised.

    • In the 1940's the research team developing the materials required to produce the first atomic bombs dropped on Hiroshima and Nagasaki realised that 'trans-uranium' elements were being formed in nuclear reactions (see radioactivity-nuclear reactions page).

    • From element 93 to 112, plus 114 and 116 are now known, but sometimes just a few atoms from a cyclotron experiment and all are highly radioactive due to having very unstable nuclei.

    • BUT the structure of the bottom part of the periodic table will continue to grow and grow!

    • Physicists are hoping to eventually make some 'nuclear stable' super-heavy metallic elements around atomic number 150 onwards? 



6. Where do we get the elements from?

  • The ultimate origin of all elements is the nuclear reactions that go on when stars are formed from inter-stellar dust and gas forming a huge combined mass due to gravity, and then 'chunks' of a star 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.

  • 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 liquified 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 usually found 'native' as the yellow/silver elemental metal.

    • Relatively unreactive metals like copper and silver, can also be found in their elemental form in mineral deposits as well as 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 sulphates, but it can occur as relatively pure sulphur in yellow mineral beds of the element.

  • -



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

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

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

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

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

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

Graphite, diamond and buckminsterfullerene are all solid allotropes of the element carbon and have significantly different physical and in some ways chemical properties! (details on bonding page)

Rhombic and monoclinic sulphur have different geometrical crystal structures, that is different ways of packing the sulphur atoms (which are actually both made up of different packing arrangements of S8 ring molecules). They have different solubilities and melting points. There is also a 3rd unstable allotrope of sulfur called plastic sulphur made by pouring boiling molten sulphur into cold water which forms a black plastic material consisting of chains of sulphur atoms -S-S-S-S-S- etc..

It doesn't matter which isotopes make up the structure of any of an element's allotropes described above, so to summarise by one example ...

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

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

Links to quizzes and worksheets based on the Periodic Table

PLEASE NOTE that these LINKS are for A Level Students ONLY

Advanced Level Inorganic Chemistry Periodic Table Index * Part 1 Periodic Table history * Part 2 Electron configurations, spectroscopy, hydrogen spectrum, ionisation energies * Part 3 Period 1 survey H to He * Part 4 Period 2 survey Li to Ne * Part 5 Period 3 survey Na to Ar * Part 6 Period 4 survey K to Kr and important trends down a group * Part 7 s-block Groups 1/2 Alkali Metals/Alkaline Earth Metals * Part 8  p-block Groups 3/13 to 0/18 * Part 9 Group 7/17 The Halogens * Part 10 3d block elements & Transition Metal Series * Part 11 Group & Series data & periodicity plots * All 11 Parts have their own sub-indexes near the top of the pages

Revision KS4 Science GCSE/IGCSE/O level Chemistry Information Study Notes for revising for AQA GCSE Science, Edexcel IGCSE Chemistry & OCR 21st Century Chemistry Science, OCR Gateway Science Chemistry WJEC gcse science chemistry CCEA/CEA gcse science chemistry O Level Chemistry (revise courses equal to US grade 8, grade 9 grade 10) science chemistry courses revision guides


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