3.1.1

Periodicity

(a)(i)→periodic table: list of all known elements arranged in order of increasing atomic number [proton number] from 1 to 118

(ii)→periodicity: periods [rows] showing repeating trends in physical and chemical properties [periodic trends]

→periods and electrons in periodic table: elements with one shell are placed in 1st row [H and He], elements with 2 shells in 2nd row and so on

(iii)→elements in the same group [columns] shows similar chemical properties →H and He are allocated a group based on similarities in physical/chemical properties as have unusual e- config; He is group 0 and H its own group

→groups and electrons in periodic table: elements of the same groups have the same amount of valence electrons

Periodic trend in e- config and ionisation energy

(b)(i)→periodic trend in e- config:

  • across period 2: always 2s subshell [2e-], then 2p subshell [6e-]

  • across period 3: 3s subshell [2e-], then 3p subshell [6e-]

(ii)→classification of elements in s,p,d blocks: [image]

  • s block elemts only have s e- in outer shell

  • p block elements have at least 1 p e- in outer shell

  • d block elements have at least 1 d block and 1 s block outershell

(c)(i)→first ionisation energy: removal of 1 mol of electrons from 1 mol of gaseous atoms e.g. Na (g) → Na+ (g) + e-  [First ionisation energy

= +496 kJ mol-1]

→ionisation energy across a period increases:

  • increased nuclear charge as more protons so more attraction holding onto outer e-

  • shielding remains constant so no effect

  • decreased atomic radius so outer electrons are more attracted to nucleus

  • this means outer e- held more tighly by nucleus so harder to remove

→ionisation energy down a group decreases:

  • increase in nuclear charge as higher no of protons but

  • increased shielding as more shells between outer e- and nucleus so less attraction

  • increased atomic radius so outer e- further away from nucleus

  • outer e- held loosley so easier to remove

→Exceptions to first ionisation energy rules:

  1. Beryllium and Boron: slight decrease from Be to B as the outer e- in B is in 2p subshell, which is further away from nucleus than 2s subshell

  2. Nitrogen and Oxygen: slight decrease from N to O as paired e- in O 2p subshell repel each other making it easier to remove

(ii)→successive ionisation energies of an element: increases as removing electron from +ve ion more difficult than neutral atom; as more e- removed the attractive forces increase due to decreased shielding and increase in proton e- ratio

Periodic trend in structure and melting point

(d)(i)→metallic bonding: strong electrostatic attraction between +ve cations and delocalised e-

(ii)→metallic lattive structure: metal atoms tighly packed in lattice structures held together by +ve ions and sea of delocalised e-

(e)→giant covalent lattices: large number of atoms are bonded by covalent bonds

→diamond: a giant covalent lattice of C atoms; each C atom bonded w 4 others in a tetrahedral arrangment, resulting with strong bonds in all directions [as is hardest substance known] [image]

→graphite: each C atom bonded to 3 others in hexagonal layers w bond angle 120degrees; spare electrons are delocalised between layers; atoms in same layer held by strong covalent bonds but layers held by weak intermolecular bonds allowing sliding [image]

→graphene: single layer of C atoms bonded in repeated pattern of hexagons [image]

→silicon oxide: each Si atom shared by 4 O and each O shared by 2 Si; adopts same shape as diamond [image]

(f)→metallic lattices properties: high mp/bp as lots of energy required to overcome strong electrostatic forces of attraction between +ve ions and sea of e-; no solubility; conducts electricity in l and s states as e- free to move around and carry charge

→giant covalent lattices properties: high mp/bp as large number of strong covelent bonds require lots of energy to break; graphite soft fue to weak intermolecular forces between carbon layers but rest are hard as difficult to break the 3d network of covalent bonds; insoluable; do not conduct electricity except graphite/graphene due to delocalised e-

(g)→melting points across period 2 and 3:

  • increases from group 1 to 4 as g1/2/3 have metallic bonding requiring large amounts of energy due to increased attraction between more electrons and +ve ions

  • group 4 has a giant covalent structure which needs lots of energy to overcome

  • sharp decrease from group 14 to 15 as have simple molular structures with weak london forces requiring little energy to overcome