1. Topic 13 Periodicity HL

2. Ionic or covalent bonding?
H-Cl
Na+ Cl-
Cl-Cl

3. 13.1 Trends across third period; Chlorides
When you go  the number of valence electrons increase => increase the number of valence electrons to form bonds.
NaCl, MgCl2, AlCl3 (Al2Cl6(g)), SiCl4, PCl5 (PCl3 exist), (sulphur chlorides not required), (Cl2)

4. Chlorides of metals (NaCl, MgCl2, AlCl3 )
Ionically bonded crystalline solids with high melting points.
Dissolves in water without a chemical reaction to its ions:
NaCl (s) → Na+ (aq) + Cl- (aq)
Conduct electricity in melted or in aqueous solution.

5. Chlorides of non-metals (SiCl4, PCl5 )
Molecular covalent structure.
Weak forces between molecules => low melting and boiling points.
Don’t conduct electricity (no ions and no mobile charges).

6. Reacts with water: Hydrolysis
PCl3 + 3 H2O  H3PO3 + 3 HCl Acidic solution (Phosphoric(III) acid, oxyacid of the element) H3PO3 + H2O H3O+ + H2PO3- The oxyacid may also dissociate into acidic oxoniumions.

7. In water the chlorides will conduct electricity; Cl- (aq).
Chlorine, Cl2, if seen as Chlorine chloride, behaves in the same way:
React with water in a hydrolysis reaction
Cl2 + H2O  HCl + HClO
Aluminium chloride reacts as a non-metal chloride due to small size and high charge. It’s very reactive with water:
AlCl3 + H2O  Al2O3 + 6 HCl

8. Oxides- across period 3 Trend: From basic to acidic character
Base Acid
Na2O, MgO, Al2O3, SiO2, P4O10, SO3 (SO2), Cl2O7 (Cl2O, Cl2O3, Cl2O5)
Ionic Giant Covalent
structure

9. Left side- oxides are basic
Na2O + H2O  2 Na+ + 2 OH-
Magnesium hydroxide only weakly dissociated because of low solubility.
Reacts with acids (basic oxides):
MgO(s) + 2 H+  Mg2+ + H2O

10. In the centre- oxides are amphoteric
Both aluminium and silicon oxides are almost insoluble
Aluminium oxides have amphoteric properties; reacts with both base and acid
Al2O3(s) + 6 H+  2 Al H2O
Al2O3(s) + 2 OH- + 3 H2O  2 Al(OH)4-(aq)
Silicon dioxide can show weakly acidic properties; reacts with strong alkali to form silicates
Giant covalent lattices with high melting and boiling points

11. To the right in period 3
Molecular bonding: Gases, liquids or low melting points
The elements can often form 2 or more oxides with different state of oxidation.
Reacts with water to form acids.
SO3(g) + H2O  H2SO4
H2SO4 + H2O  H+ + HSO4-
  Cl2 + H2O  H+ +Cl- + HClO

12. 13.2 First row d-block elements (Sc Zn) The transition elements
An element that contain an incomplete d level of electrons in one or more oxidation states
d-orbitals starts to fill up with electrons
They have some common characteristics
(except Sc and Zn):
A variety of stable oxidation states
The ability to form ions
Coloured ions
Catalytic activity

14. Oxidation states The 4s and 3d orbitals are quite close in energy
The electrons in 4s orbitals can easily be lost
Gives stable state to the right of the d-block. To the left it’s a powerful reductant. (Ti2+ + water  Hydrogen)
Sc to Mn can loose all 4s and 3d electrons and stay stable. More to the right they become strong oxidants
Highest oxidation state usually occur as oxanions: E.g. dichromate (Cr2O72-), permanganate (MnO4-)

15. Energy 3d 3d 4s 4s
Mn2+ ion [Ar] 3d5
4s higher than 3d
Mn atom [Ar]4s23d5
4s lower than 3d

16. Common oxidation states of the d-block elementsV Cr Mn Fe Co Ni Cu +7 X +6 +5 +4 +3
(x) +2 +1

17. All transition elements can show an oxidation number of +2
You should be familiar with
Cr (+3, +6),
Mn (+4, +7)
Cu (+1,+2)

18. In solution: Ligand
Ions of d-block elements have unfilled orbital's. These unfilled orbital's can attract a pair of electrons from an other compound = ligand.
The ligand must have free (non-bonding) electron pair that they can donate to the ion.
E.g. H2O, NH3, Cl-, CN-

19. In solution: Complex ion
The ion and the ligand form a dative bond,
co-ordinate bond(covalent) bond
The Ion + ligands = complex ion

20. Examples of complex ions
Most complex ions have either six ligands arranged octahedrally around the central ion (often water or ammonia ligands) or four ligands arranged tetrahedrally (often chloride ligands)
[Cu(NH3)4]2+ (forms when an excess of ammonia is added to Cu(II)-salt)
[Ag(NH3)2]+
[Fe(H2O)6]3+
[Fe(CN)6]3-
[CuCl4]2-
Complex formation can stabilise certain oxidation states and affect the solubility of the ion

21. Complexes have often specific colours
In an isolated atom all d-orbital’s have the same energy.
The Ligands in a complex ion affect the energy in the d-orbital’s.
The orbitals split up to two groups with different energy. The energy gap is in the visible region.
When light going through a transition metal solution energy is absorb when electrons are lifted from the lower level to the higher.

24. White light (all colours) hits Copper(II) salt and red and yellow light absorbs => blue-green colour.
Sc3+ and Ti4+ : no electrons in d-orbitals => colourless
Zn2+ : filled d-orbital => colourless

25. Catalytic activity
Catalyst is a substance that speeds up a reaction without being consumed by it self. Reduce the activation energy.
Transition metals often have catalytic behaviour due to:
Ability to form complexes. Close contact.
Many oxidation states. Easy to lose or gain electrons in redox reactions.

26. Homogeneous catalyst In the same phase as the reactants
E.g. dissolved ion in water solution

27. Heterogeneous catalyst
On the surface of the metal.
E.g.
MnO2, Manganese(IV)oxide: 2 H2O2  2 H2O + O2
Ni: Alkenes + hydrogen  Alkanes
Fe: Haber process, N2 + 3 H2  2 NH3
The worldwide ammonia production in 2004 was 109 million metric tonnes.[

28. V2O5, vanadium(V)oxide: in the Contact process (manufacture sulphuric acid)
2 SO2(g) + O2(g)  2 SO3(g)
SO3 + H2O  H2SO4  
Sulphuric acid. 165 million tonnes, with an approximate value of US$8 billion. Principal uses include ore processing, fertilizer manufacturing, oil refining, wastewater processing, and chemical synthesis.

29. Co in vitamin B12
Pd and Pd in catalytic converters