STPM Chemistry Form 6 Notes – Terminology and Concepts: The Periodic Table (Part 4)

by BerryBerryTeacher

in Berry Reference (Notes)

This is Part 4 of the “STPM Chemistry Form 6 notes on The Periodic Table” from Berry Berry Easy which is without doubt, one of the hardest part of this periodic table chapters. Here, you’ll be learning all about the physical properties of vaporisation, its electrical properties in terms of electrical conductivity and the cumbersome (but still easy) ionisation energy. At first glance, it might seemed a lot to take for a short sub-chapter, but in reality, it IS a lot to study for a short subchapter. So do try to understand everything on this post before proceeding to other subtopics.

Berry Readers will most likely memorise this than trying to really understand this. Berry Berry Teacher’s advise is to understand the concepts rather than memorising it since there are too much to be memorised for this subchapter. If you really cannot understand this, try mnemonic methods instead. (For those who cannot understand, try dropping a line at the comments, we would be interested to know what you don’t understand) But really, try to understand as this is one topic which is very popular in exams.

STPM Chemistry Form 6 Notes – Terminology and Concepts: The Periodic Table (Part 4)

A) Boiling Point, Melting Point and Enthalpy of Vaporisation

Enthalpy of vaporisation – the heat energy required to convert 1 mol of a liquid to its vapour at the boiling point of the liquid.

Example: Period 2 (Li, Be, B, C, N, O, F, Ne) & Period 3 (Na, Mg, Al, Si, P, S, Cl, Ar)

Across the period: the element become less metallic.

Li, Be, Na, Mg and Al are metals (metal lattice):

  • B.P, M.P and enthalpy of vaporisation increase and the atoms are held together by strong metallic bond.
  • Increasing the number of valence electrons cause the strength of the metallic bond increase.

B, C (graphite) and Si are metalloids & C (diamond) is non-metal (giant covalent molecule):

  • B.P, M.P and enthalpy of vaporisation are very high.
  • The atoms are held together by strong covalent bonds which form giant covalent structure (crystal lattice structure) in a 3-D structure.
  • All the covalent bonds are needed to be broken before the solid melts.

N, O, F, P, S, Cl are non-metallic elements (simple molecular structure):

  • B.P, M.P and enthalpy of vaporisation are relatively low that involves only the breaking of weak Van der Waals forces.
  • N2, O2, F2, P4, S8, Cl2 consist of small and discrete molecules.
  • The covalent bonds within the molecules are very strong, but the Van der Waals forces of attraction between the molecules are very weak.

Ne and Ar are non-metallic (monoatomic structure):

  • B.P, M.P and enthalpy of vaporisation are very low.
  • Noble gases are uncombined atoms and have very weak Van der Waals forces of attraction between the atoms.

B) Electrical Conductivity

  • All metals (Li, Be, Na, Mg and Al) are good conductors either in the solid or molten state.
  • Metals have delocalised electrons which will move freely across the metal in the solid lattice structure when an electrical potential or voltage is applied.
  • Non-metals (C diamond, N, O, F, Ne, P, S, Cl, Ar) are non-conductors.
  • All the valence electrons in non-metals are used to form covalent bonds between atoms and there are no mobile electrons in the structure. Ne and Ar are noble gases and have the stable octet electronic configuration and do not have any mobile electrons.
  • Metalloids (C graphite, Si) are semi-conductor.
  • Conductivity of metalloid increases with the increasing of temperature.

C) Ionisation Energy

Ionisation energy of an element – the amount of energy required to pull one electron off an atom.

Down the group in the periodic table, ionisation energy decreases because of the screening effect / shielding effect (electrons in low-energy levels repel electrons in higher-energy levels away from the nucleus)

First ionisation energy of an element – the minimum energy required to remove 1 mol of electrons from 1 mol of atoms in the gaseous state.

M(g) –> M+(g) + e, ΔH = first ionisation energy

Second ionisation energy of an element – the minimum energy required to remove 1 mol of electrons from 1 mol of unipositive ion in the gaseous state.

M+(g) –> M2+(g) + e, ΔH = second ionisation energy

Third ionisation energy of an element

M2+(g) –> M3+(g) + e, ΔH = third ionisation energy

Fourth ionisation energy of an element

M3+(g) –> M4+(g) + e, ΔH = fourth ionisation energy

i) Factors affecting ionisation energy

  • Distance of the outer electrons from the nucleus (atomic size)
  • Size of the nuclear charge (nuclear charge)
  • Screening effect of the electrons in the inner shells (screening effect)

ii) Ionisation energy across a period 2 and period 3

  • The first ionisation energy increase with increasing proton numbers for the elements (atomic size decreases, nuclear charge increases and the screening effect remains the same).
  • First ionisation of Be (period 2) and Mg (period 3) is higher than expected because the first electron to be removed is from a fully filled s orbital.
  • Be: 1s2 2s2
  • Mg: 1s2 2s2 2p6 3s2
  • First ionisation of N (period 2) and P (period 3) is is higher than expected because the first electron to be removed is from a half filled p orbital.
  • N: 1s2 2s2 2p3
  • P: 1s2 2s2 2p6 3s2 3p3

So there you go. Not so difficult as you though, isn’t it? Try and revise all the way from the first post on this chapter about the periodic table for STPM Chemistry Form 6.

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