Periodicity

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Factors Influencing Ionisation Energy

  • Nuclear charge - the more protons in the nucleus of the atom, the greater the nuclear charge.
  • Distance of the electrons from the nucleus - the more shells there are, the greater the distance between the outermost electrons and the nucleaus so the easier the electrons are to remove.
  • Sheilding effect - any inner shells between the electrons and the nucleus repel outer electrons. The more repulsion there is, the less attraction and the easier the electrons are to remove.
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Trends In Ionisation Energy

Ionisation energy increases across a period:

  • Proton number increases, therefore the attractions for the outer electrons increases.
  • Electrons enter the same shell so there is no increase in shielding.
  • Outer electrons are held in more tightly (smaller radius).
  • Nuclear charge is the dominant factor.
  • Therefore electrons are harder to remove so ionisation energy increases.

Ionisation energy decreases down a group:

  • Proton number increases.
  • Electrons fill new shells and so shielding increases.
  • Outer electrons are futher from the nucleus (larger radius).
  • The increased nuclear charge is outweighed by the change in shielding and atomic radius (both increased) which are the dominant factors.
  • Therefore it is easier to remove and electron so ionisation energy decreases.
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Sub-Shell Trends In First Ionisation Energy

Graph showing trend in first ionisation energy (http://chemed.chem.purdue.edu/genchem/topicreview/bp/ch7/graphics/ch7_14.gif)

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Explaining Sub-Shell Trends In First Ionisation En

Explain why boron has a lower first ionisation energy than beryllium
The first electron in Be is removed from the 2s subshell where as in B it's removed from the 2p. 2p is higher in energy than 2s so requires less energy to remove the first electron.

Explain why oxygen has a lower first ionisation energy than nitrogen
Highest energy electrons in the 2p subshell. Oxygen has a pair of electrons in one orbital which repel so it's easier to remove an electron.

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Successive Ionisation Energies

Each time an electron is removed the ionisation energy increases (requires more energy to remove a second electron than the first one):
The electron is being removed from a charged partical (increasing the proton to electron ratio) so the same number of protons are attracting fewer electrons, thus meaning a higher ionisation energy is needed.

There is a large difference in value between some ionisation energies:
Some electrons are removed from a different shell closer to the nucleus with less shielding, so require a higher ionisation energy.

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The Periodic Table

  • Periods (rows) and groups (columns)
  • There are eight groups
  • There are four periods of transitions metal elements that slot into the table between groups two and three
  • The elements are arranged in order of atomic number

Names for the groups:

  • Group one = the alkali metals
  • Group two = the alkaline earth metals
  • Group seven = the halogens
  • Group eight = the noble gases
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Atomic Radii

As the atomic radius increases so does the ionisation energy.

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Periodic Trends In Bonding And Structure

Metallic Bonding

  • Fixed metal cations surrounded by mobile, delocalised electrons
  • These electrons mean metals can conduct as a solid and as a liquid
  • The strong metallic bonds mean that high temperatures are needed to provide the large amount of energy to overcome the electrostatic attractions between cations and electrons
  • Metals have high boiling and melting points

Giant Covalent Structures

  • Giant covalent lattices are held together by a network of strong covalent bonds (eg. diamond and silicon)
  • The strong covalent bonds need a large amount of energy to overcome them: they have high boiling and melting points
  • They are also insoluble in most solvents - the covalent bonds are too strong to be broken by interactions with solvent molecules
  • Most do not conduct electricity - all outer shell electrons are involved in covalent bonding. However, graphite and graphene conduct as they have mobile, delocalised electrons
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Melting Points And Boiling Points

  • The melting and boiling points of a particular element is dependent on the bonding present in the structure of the element
  • Metals have high boiling points because they contain metallic bonding
  • Non metals with giant covalent structures like silicon have high melting points because the covalent bonds are strong and difficult to break
  • Simple covalent molecules like phosphorus and chlorine have low melting and boiling points because they are held together by intermolecular forces which are weaker and easily broken down
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Electrical Conductivity

  • Metallic elements are good conductors of electricity
  • Their structures contain delocalised electrons which are free to move through the structure
  • Metallic character and therefore conductivity increases to a maximum in group three
  • Metallic character also increases down the group
  • Elements beyond silicon show no conduction as all the electrons are involved in bonding and no free charge carriers are present
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