1. Summary of the structure of the Periodic Table

1a. The basic structure of the Periodic Table

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, 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 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.
  • Most elements on the left are METALS and their typical properties are described in section 2a. e.g. elements 3 to 4 (lithium to beryllium), elements 11 to 13 (sodium to aluminium), elements 19 to 31 (potassium to gallium), elements 37 to 50 (rubidium to tin).
  • The elements on the right are NON-METALS and their typical properties are described in section 2b. e.g. elements 1 to 2 (hydrogen to helium), elements 5 to 10 (boron to neon), elements 15 to 18 (phosphorus to argon), elements 35 to 36 (bromine to krypton).
  • BUT a few elements are referred to as SEMI-METALS which have mixed metal/non-metal character and not so easy to classify, see section 2c. They occur in a diagonal band (down and L to R) e.g.  silicon (₁₄Si), germanium (₃₂Ge), arsenic (₃₃As) and tellurium (₅₂Te).
• 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 (detailed GCSE notes) 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.
• 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.
  • 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 (see below).

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 (the limit GCSE students need to know about - details on the Atomic Structure page).

  • 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 Cl^2- 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.

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

• 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.).
• Apart from hydrogen (doesn't really fit in any group), and 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.
  • * 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.
• 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 (Group 0 or 8 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. The 1st Transition Metals Series (detailed GCSE notes on Fe, Cr, Cu etc.)

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 19 are definitely are non-metal but about 7 more are semi-metals of mixed 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.
• There tends to be gradual changes in physical and chemical properties down a group e.g.
  • Down Group 1 (Alkali Metals) and Group 2 the metals get more reactive.
  • Down Group 7 (Halogens) the non-metals get less reactive, their colour gets darker, their melting/boiling points increase.
  • 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⁺O^2- 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 8, 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 chlorine valency 1 and oxygen valency 2, 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.
• 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.
  • 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! (and in scrabble too?)
  • 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 5 are given below and examples of all these formulae can be found in the Period 2-3 table further down.

  • valency of X	1	2	3	4	5	6
    formula of hydride	XH	XH2	XH3	XH4	-	-
    formula of chloride	XCl	XCl2	XCl3	XCl4	XCl5	-
    formula of oxide	X2O	XO	X2O3	XO2	X2O5	XO3

  • 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.
  • In 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.

2. Comparing Physical and Chemical Properties of Elements

2a. Typical Properties of Metallic Elements

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 [GCSE notes] 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).
• 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 reactions/equations [1] and [2] 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⁺O^2-, sodium oxide Na₂O or (Na⁺)₂O^2- ,
    • and aluminium oxide Al₂O₃ or (Al³⁺)₂(O^2-)₃ 
    • * At least at GCSE level, but there are some chloride exceptions at Advanced level such as FeCl₃ and AlCl₃.

Reactivity of Metals Notes and Metal Extraction Notes

2b. Typical Properties of Non-metallic Elements

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, see GCSE notes).
• Usually poor conductors of heat and electricity (exceptions like carbon in the form of graphite).
• N0n-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 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^- ==> 2O^2- (more simply O + 2e^- ==> O^2- 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 CH_4(g), sulphur dioxide SO_2(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⁺O^2-

2c. The Properties of Semi-metals or Metalloids

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

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 0/8 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

and six extra periodic table data links if needed

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

with links to more detailed GCSE notes where necessary

Group 1 Alkali Metals

• The very reactive Group 1 The Alkali Metals  [GCSE notes] have low density (some float on water).
• They readily react with non-metals to form ionic compounds e.g. NaCl or Na⁺Cl^-, Li₂O or (Li⁺)₂O^2-.
• These 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 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.
• AS Advanced 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").
• 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.
• AS Advanced Chemistry Notes on Group 1 and Group 2 Metals

Group 3/13

• Group 3/13 contains the metal Aluminium (see metal extraction notes and uses of metals: comparison with transition metals or titanium/steel/alloy uses notes)

Group 4/14

• Group 4/14 contains the non-metal carbon - 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.

Group 5/15

• Group 5/15 contains the non-metal nitrogen - important element in natural and manmade artificial fertilisers (see GCSE notes on ammonia and nitric acid). Nitrogen forms 79% (⁴/₅th's) of air.

Group 6/16

• 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 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 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 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).

Group 7/17 The Halogens

• The Group 7 Halogens [GCSE notes] are coloured non-metals with low melting points and boiling points.
• 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.

Group 0/18 The Noble Gases

• The Group 8/18 Noble Gases [GCSE 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 arrangements.
• 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.

The 1st Transition Metal Series

• The ten horizontal elements Sc to Zn are called the 1st series of Transition Metal Elements [GCSE 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)  pages.

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

5a. The early classification of Antoine Lavoisier of 1789

• 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.
• 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

• 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 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'

• He 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 particularly mixed up compared to the modern periodic table!

5d. Dmitri Mendeleev's Periodic Table of 1869

• Mendeleev laid out all the known elements in order of 'atomic weight' (what we know 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! Smart thinking!

• 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!

• 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'.

5e. The full modern version of the Periodic Table