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Topic 3 Periodicity . SL + HL. 3.1 The periodic table of the elements. The elements are arranged in order of increasing atomic number , reading from left to right starting in the top left corner. Names and symbols are given in Chemistry Data Booklet

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3 1 the periodic table of the elements
3.1 The periodic table of the elements

The elements are arranged in order of increasing atomic number, reading from left to right starting in the top left corner.

Names and symbols are given in Chemistry Data Booklet

  • Is divided into groups and periods:

Groups: Vertical

Periods: Horizontal

  • Going across a horizontal row.
  • From reactive metals to reactive non-metals and the Noble gases.
  • Gives the number of Energy level (Shell) that contains electrons (K, L, M, N).
  • Numbered 1 to 7, 0 from left to right

(there are also some other numbering systems e.g. 1A2A3B-12B3A-8A).

  • Elements with similar chemical characteristics in the groups.
  • Groups 1, 2, 3-7 gives the number of valence electrons (electrons in the outer shell)

Valence electrons- the electrons in the outermostshells

(the electrons in the outermost s/p-orbitals)

Group 1: Alkali metals. 1 valence electron Na+

Group 2: Alkaline earth metals. 2 valence e- Mg2+

Group 3: 3 valence electrons Al3+

Group 4: 4 valence electrons C4+/C4-

Group 5: 5 valence electrons N3-

Group 6: 6 valence electrons O2-

Group 7: Halogens. 7 valence electrons Cl-

Group 0: Noble gases. 8/0 valence electrons He0

Transition metals (normally M+ - M3+)

3 2 physical properties
3.2 Physical properties
  • Periodic trends
  • Data is listed in Chemistry Data Booklet.

(not allowed for paper 1)

going down a group
Going down a group
  • More energy levels (shells) filled with electrons

=> Size of atoms increasing.

  • Valence electrons in a higher level =>
  • Decreasing Ionisation energy (the energy it takes to remove an electron from an atom) down the group and decreasing
  • Decreasing electronegativity (how strongly the atom attracts the electrons in chemical bonds.)
going across a period
Going across a period
  • Number of protons/charge increase in the nucleus.
  • Electrons (on same energy level) more strongly attracted to the nucleus.
  • Atomic radii decrease
  • Ionisation energy increase (but not smooth, see topic 2)
  • Electronegativity increase
positive ions
Positive Ions
  • Decrease in size
    • Loss of a whole outer shell (Na  Na+)
    • Less electron-electron repulsion
    • Higher effective charge is experienced

negative ions
Negative Ions
  • Increase in size
    • Increase of electron-electron repulsion
    • Lesser effective charge is experienced
melting points boiling points and density
Melting points(boiling points and density)
  • Depends to a considerable extent to the nature of the bonding between particles of the element.

Bonding types: metals-metallic bonds (strong)

Non-metals- molecules with covalent bonds, van der Waals forces between (weak)


Metallic bondsbetween metal atoms- STRONG

Van der Waals bondsbetweenmolecules- WEAK

melting point across period 2
Melting point- across period 2
  • Increase to start with: increasing number of valence electrons Li- Be => stronger metallic bonds in Be
  • Peaks at carbon, C: Giant covalently bonded structure
  • Drop in m.p. at N-Ne: Covalent bonding and van der Waal’s forces between the molecules
melting point across group 1
Melting point- across group 1
  • Decreasing m.p. down the group. Unusual behaviour. Normally the m.p increase down a group


3 3 chemical properties alkali metals
3.3 Chemical properties- Alkali metals
  • One atom in valence electronic shell => easily lost to form a Noble gas electron configuration => reactive metals
  • More reactive down the group. Decrease in Ionisation energy.
alkali metals cont
Alkali metals cont.
  • Reaction with water:

2 Na (s) + 2 H2O (l)  2 Na+ (aq) + 2 OH-(aq) + H2(g)

  • Reaction with halogens:

2 K (s) + Cl2 (g) 2 KCl (s)

  • Diatomic molecules. F2(g), Cl2(g), Br2(l), I2(s)
  • Van der Waal’s forces between molecules. Increasing with molar mass.
  • Coloured: F2, Cl2 green-yellow,
  • Br2 brown, I2 brown-purple
  • Reactive: 7 valence electrons, gain 1 => X-
halogens cont
Halogens, cont.
  • Reactivity decreases down the group. Further from nucleus => weaker attraction
  • Can give salts with metals. Usually water soluble.

2 K (s) + Cl2 (g)  2 KCl (s) potassium chloride

2Ag (s) + Cl2 (g)  2 AgCl (s) silver chloride

  • Silver halides insoluble. (Test for halides)

Ag+ (aq) + Cl- (aq)  2 AgCl (s) white precipitate

Ag+ (aq) + Br- (aq)  2 AgBr (s) yellow precipitate

halogens cont1
Halogens, cont.
  • Reaction between Halogens, X2, and Halides, X-,: Going down: Decrease of Electronegativity and Oxidation power.

Oxidation power: the ability to oxidise I- to I2 and Br- to Br2 in the example below

Cl2 + 2 Br- 2 Cl- + Br2

Cl2 + 2 I- 2 Cl- + I2 Cl2 can oxidise Br- and I-

Br2 + 2 I- 2 Br- + I2 Br2can only oxidise I-

metal oxides across period 3
Metal oxides -across period 3
  • The oxides of these elements will form basic solutions (pH>7)

Na2O + H2O 2 NaOH (aq)

Sodium oxide Sodium hydroxide

MgO+ H2O  Mg(OH)2(aq)

Magnesium oxide Magnesium hydroxide

non metal oxides across period 3
Non-metal oxides -across period 3
  • Most oxides of these elements will form acidic solutions (pH<7)

P4O10+ 6 H2O  4 H3PO4(aq)

Phosphoric(V) acid

SO3+ H2O H2SO4



Amphoteric oxides: some elements forming oxides that can be either a base or an acid- for example Al2O3

  • Oxides with higher state of oxidation is more acidic than the corresponding compound with lower state of oxidation.