Chapter 7: Periodicity

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7.1 The periodic table

Atomic number: elements are arranged in order of increasing atomic number

Periods: the period number give the number of shells that an atom of the element has

Periodicity: a repeating trend in properties of the elements across each period in the periodic table

Electron configuration:

Across a period the s-subshell fills with 2 electrons and the p-subshell fills with 6 electrons

Trend: for each period the s and p-subshells are filled in the same way

Groups: each element in a group has atoms with the same number of outer shell electrons = similar electron configuration gives elements similar chemical properties

S-block: 1) alkali metals, 2) alkali earth metals, D-block: transition metals, P-block: 5) pnictogens, 6) chalcogens, 7) halogens, 0) noble gases

Blocks: s, p, d, f: depends on which subshell contains the element's outermost electrons

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7.2 Ionisation energies

  • First ionisation energy: the energy required to remove one electron from each atom in one mole of gaseous atoms to form one mole of gaseous 1+ ions
    • Second ionisation energy: the energy required to remove one electron from each 1+ ion in one mole of gaseous 1+ ions to form one mole of gaseous 2+ ions
  • Equation for the 1st IE of sodium: Na (g) ---> Na+ (g) + e-
  • Ionisation energy measures how easily an atom loses e- to form positive ions
  • The largest ionisation energy is from the shell closest to the nucleus
  • Atomic radius: larger radius = lower nuclear attraction between nucleus and outer shell electrons (because outer shell electrons are further away from nucleus)
  • Nuclear charge: greater number of protons = greater nuclear charge+
  • Shielding: greater shielding effect = lower nuclear attraction

Successive ionisation energies are greater than previous IE's because with every electron that is removed, there is a stronger nuclear attraction for the remaining electrons.

Nuclear attraction on remaining electrons increases with every electron that is removed so more IE energy is required to remove successive electrons

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7.2 Ionisation energies

Successive IE and shells

A large increase between 2 ionisation energies suggests the next electron has been removed from a different shell that: is closer to the nucleus and has less shielding

Ionisation energies allow prediction of the number of electrons in the outer shell, the group the element is in and the identity of the element

  • Number of IE's before the first large increase = the number of electrons in the outer shell
  • Each large difference between 2 IE's = the number of shells
  • Number of ionisation energies = total number of electrons (gives the identity of element)
  • General trends:
  • steady increase across a period
  • sharp decrease between the end of one period and the start of the next
  • decrease down a group
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7.2 Ionisation energies

Down a group: decreases

  • Nuclear charge increases but this is outweighed by the
  • atomic radius increase,
  • sheilding also increases which means the
  • nuclear attraction decreases and so the
  • first ionisation energy decreases

Across a period: (steady) increase

  • Nuclear charge increases so the
  • atomic radius decreases and elements have the
  • same shielding which means the
  • nuclear attraction increases so the
  • first ionisation energy increases

Subshell trend: general increase but falls at group 3 (B & Al) and group 6 (O & S)

subshell trend explanation...

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7.2 Ionisation energies

Fall at group 3 marks the start of filling the p-subshell

P-subshells have a higher energy than s-subshells 

It is easier to remove the p electron than one of the s electrons

so the first ionisation energy for the group 3 element is slightly less than the group 2 element

Fall at group 6 marks the start of electron pairing in the p-orbitals (of p-subshells)

Paired electrons repel each other so it is easier to remove a paired electron e.g. 2p4 than an unpaired electron e.g. 2p3

so the first ionisation energy for the group 6 element is slightly less than the group 5 element

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7.3 Periodic trends in bonding and structure

Metallic bonding: the electrostatic attraction between positive metal ions and delocalised electrons

Each atom in the solid metal structure has donated its outer shell electrons into a shared pool of electrons which are delocalised throughout the structure

  • Cations consist of a nucleus and the inner electron shells of the metal atoms
  • Ions are in fixed positions within the giant metallic lattice, electrons are mobile
  • Diagram: charges must balance; 8 1+ ions = 8 electrons, 8 2+ ions = 16 electrons in lattice

Electrical conductivity: conduct in solid, liquid and aqueous states because delocalised electrons are mobile and free to move through the lattice and carry a charge

Melting / boiling point (increase to group 4, then decreases):

Strong electrostatic attraction between metal cations and delocalised electrons, so high temperature needed to provide the large quantity of energy required to overcome the bonds

Solubility: insoluble; metals react with H2O (e.g. Na), don't dissolve

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7.3 Periodic trends in bonding and structure

Giant covalent structures (C and Si):

  • C uses its 4 outer shell electrons to form 4 covalent bonds with other C atoms (diamond)
  • Si uses its 4 outer shell electrons to form 4 covalent bonds with other Si atoms
    • Tetrahedral arrangment, bond angle 109.5
    • Strong covalent bonds, very stable structures, hard to break down

Melting / boiling points

  • Strong covalent bonds so high temperature needed to provide the large quantity of energy required to break the bonds = high melting / boiling point

Solubility: insoluble in all solvents because covalent bonds are too strong to break

Electrical conductivityall 4 outer shell electrons involved in the bonds so there are no delocalised electrons to move and carry a charge so non-conductors

Exceptions: graphite / graphene, form 3 bonds, spare electron is delocalised between the layers so can conduct electricity

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7.3 Periodic trends in bonding and structure

Melting points across a period (2 and 3)

Increase to group 4, sharp decrease, stays low from group 5 to group 0

The sharp decrease indicates a change from giant structures to simple molecular structures

From period 4 the trend only continues across s and p-blocks (excludes d-block)

Conductivity:

  • substance               solid              molten               aqueous          
  • ionic                            X                      Y                          Y               
  • giant covalent           X                       X                         X (insoluble)
  • metallic                       Y                       Y                         Y               
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