Topic 13 Periodicity HL

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

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)

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.

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).

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.

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

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

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

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 Al3+ + 3 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

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

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

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-)

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

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

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

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-

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

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

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.

http://www.chemguide.co.uk/inorganic/complexions/colour.html

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

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.

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

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.[

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.

Co in vitamin B12 Pd and Pd in catalytic converters