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 Is divided into groups and periods: Groups: Vertical Periods: Horizontal

Periods 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).

Groups 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)

Group 1: Alkali metals. 1 valence electronNa + Group 2: Alkaline earth metals. 2 valence e - Mg 2+ Group 3: 3 valence electronsAl 3+ Group 4: 4 valence electronsC 4+ /C 4- Group 5: 5 valence electronsN 3- Group 6: 6 valence electronsO 2- Group 7: Halogens. 7 valence electronsCl - Group 0: Noble gases. 8/0 valence electronsHe 0 Transition metals (normally M + - M 3+ ) Valence electrons- the electrons in the outermost shells (the electrons in the outermost s/p-orbitals)

3.2 Physical properties Periodic trends Data is listed in Chemistry Data Booklet. (not allowed for paper 1)

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 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 Decrease in size – Loss of a whole outer shell (Na  Na + ) – Less electron-electron repulsion – Higher effective charge is experienced

Negative Ions Increase in size – Increase of electron-electron repulsion – Lesser effective charge is experienced

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 bonds between metal atoms- STRONG Van der Waals bonds between molecules- WEAK

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

Decreasing m.p. down the group. Unusual behaviour. Normally the m.p increase down a group. Melting point- across group 1

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. Reaction with water: 2 Na (s) + 2 H 2 O (l)  2 Na + (aq) + 2 OH - (aq) + H 2 (g) Reaction with halogens: 2 K (s) + Cl 2 (g)  2 KCl (s)

Halogens Diatomic molecules. F 2(g), Cl 2(g), Br 2(l), I 2(s) Van der Waal’s forces between molecules. Increasing with molar mass. Coloured: F 2, Cl 2 green-yellow, Br 2 brown, I 2 brown-purple Reactive: 7 valence electrons, gain 1 => X -

Halogens, cont. Reactivity decreases down the group. Further from nucleus => weaker attraction Can give salts with metals. Usually water soluble. 2 K (s) + Cl 2 (g)  2 KCl (s) potassium chloride 2 Ag (s) + Cl 2 (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, cont. Reaction between Halogens, X 2, and Halides, X -,: Going down: Decrease of Electronegativity and Oxidation power. Oxidation power: the ability to oxidise I - to I 2 and Br - to Br 2 in the example below Cl Br -  2 Cl - + Br 2 Cl I -  2 Cl - + I 2 Cl 2 can oxidise Br - and I - Br I -  2 Br - + I 2 Br 2 can only oxidise I -

Metal oxides -across period 3 The oxides of these elements will form basic solutions (pH>7) Na 2 O + H 2 O  2 NaOH (aq) Sodium oxide Sodium hydroxide MgO + H 2 O  Mg(OH) 2 (aq) Magnesium oxide Magnesium hydroxide

Non-metal oxides -across period 3 Most oxides of these elements will form acidic solutions (pH<7) P 4 O H 2 O  4 H 3 PO 4 (aq) Phosphoric(V) acid SO 3 + H 2 O  H 2 SO 4 Sulphuric acid

Amphoteric oxides: some elements forming oxides that can be either a base or an acid- for example Al 2 O 3 Oxides with higher state of oxidation is more acidic than the corresponding compound with lower state of oxidation.