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Topic 3 Periodicity

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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Topic 3 Periodicity

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  1. Topic 3 Periodicity SL + HL

  2. 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

  3. 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).

  4. 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)

  5. 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+)

  6. 3.2 Physical properties • Periodic trends • Data is listed in Chemistry Data Booklet. (not allowed for paper 1)

  7. 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.)

  8. 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

  9. Positive Ions • Decrease in size • Loss of a whole outer shell (Na  Na+) • Less electron-electron repulsion • Higher effective charge is experienced http://intro.chem.okstate.edu/1314F97/Chapter8/Wed1112971.html

  10. Negative Ions • Increase in size • Increase of electron-electron repulsion • Lesser effective charge is experienced

  11. 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) http://www.webelements.com/periodicity/melting_point/

  12. Metallic bondsbetween metal atoms- STRONG Van der Waals bondsbetweenmolecules- WEAK

  13. 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

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

  15. 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.

  16. 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)

  17. Halogens • 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-

  18. 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

  19. 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-

  20. 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

  21. 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 Sulphuricacid

  22. 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.

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