C3.1 Development of the periodic table

  • Created by: Fiona S
  • Created on: 14-02-15 20:33

History of the Periodic Table

Chemists have always tried to arrange the elements into some kind of order, to try to explain their behaviour and to make predictions about how they would react with each other. Several tries were made over the years, notably:

  • Lavoisier - 1789
  • Dalton - 1808
  • Newlands - 1866
  • Mendeleev - 1869 
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In the 18th century, understanding of chemistry was very different from the present day. Some substances were classed as elements that today are known to be compounds or even not chemicals at all. The table below shows some of Lavoisier's table:

Acid making elements

Sulphur, Phosphorous, Charcoal (Carbon)

Gas-like elements

Light, Caloric (heat), Oxygen, Azote, (nitrogen) hydrogen

Metallic  elements

Cobalt Mercury, Copper Nickel, Gold Plantina (platinum), Iron Silver, Lead Tin, Manganese Tungsten ,Zinc

Earthy elements

Lime (calcium oxide), Magnesia (magnesium oxide), Barytes (barium sulphate), Arilla (aluminium oxide), Silex (silicon dioxide)

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Newlands arranged the elements in order of atomic weight and then grouped them near elements that were known to behave similarly. He found that a pattern emerged of a new row of eight elements, called the rule of octaves, - at least for the first 15 elements, after that, no pattern could be seen. Some of the elements are still similar to their modern positions.

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Mendeleev also arranged the elements in order of increasing weight, sometimes even switching the order of elements to force elements that had similar properties to be near each other. He also left gaps, assuming that these elements had yet to be discovered and making predictions about their properties - he was later proved to be correct when Germanium was discovered and behaved just as he had predicted.

Group 8/0 (noble gases) is missing because they are unreactive. Both Newlands and Mendeleev arranged their elements by increasing atomic mass.

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The Modern Periodic Table

  • Rows are called periods
  • They are numbered using English numbers
  • The period number tells us how many electron shells the atom has
  • Groups is the name given to columns
  • They are numbered using Roman Numerals
  • Elements in groups are very similar to each other
  • If an atom has one electron in it's outer energy level it will be in group I
  • If an atom is in group II it will have two electrons in it's outer shell and so on
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Group 1 - the alkali metals

  • Alkali metals react vigorously with water to produce hydroxides of the metal
  • The metal hydroxide dissolves in the water to produce an alkaline solution
  • So universal indicator turns purple and hydrogen gas is produced
  • We can prove it is hydrogen by placing a lighted splint near the alkali metal
  • The hydrogen ignites and makes a squeaky pop
  • As you go down the group the alkali metals become more reactive
  • Lithium reacts gently, sodium reacts more aggressively and potassium reacts so aggressively that it melts and burns with a purple flame

Lithium + water -> Lithium Hydroxide + hydrogen
Effervescence, slight movement, hydrogen gas produced
Universal Indicator: Blue

Sodium + water -> Sodium hydroxide + hydrogen
More effervescence, moved faster, melted to a ball, hydrogen gas produced
Universal Indicator: Purple

Potassium + water -> Potassium hydroxide + hydrogen
Purple flame, cracking sound, moves fastest, hydrogen gas produced
Universal Indicator: Purple

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Group 1 reactivity

As you go down the group each element has 1 extra full shell of electrons. This extra shell shields the outer shell electrons from the attraction of the positive nucleus. The outer shell electron is lost more easily. This is called shielding.

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Group 7 - the halogens

  1. Form coloured vapours
  2. Consist of molecules which are made up of pairs of atoms
  3. Form molecular compounds with other non-metallic elements
  4. Form ionic salts with metals in which the chloride, bromide or iodide (Halide Ion) carries a charge of -1

The Reactivity DECREASES as we go down group 7...

  • as we go down the group each element has an extra full shell of electrons
  • these full shells shield the outer shell from the positive nucleus
  • this makes it more difficult to gain a new electron

Group 1 vs. Group 7

  • losing an electron is easier the more full shells you have. Therefore the reactivity increases as you go down group 1
  • gaining an electron is easier the less full shells you have. Therefore the reactivity increases as you go up group 7
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Transition Metals


  • good thermal and electrical conductor
  • ductile
  • can be used as catalysts
  • high melting point (except for mercury)
  • malleable
  • sonorous
  • shiny

Group 1 vs. Transition Metals

Group 1

  • very reactive, soft and low melting point

Transition Metals

  • not very reactive, ductile, hard and strong, stays shiny, high melting point

But they are both metals, malleable and conduct electricity.

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