1. OCR: Energetics, Equilibrium and Elements

2. Physical Properties  A transition metal is a d block element that has a partially filled d-subshell of electrons.  All transition metals have high melting points – above 1000 °C  This is due to strong metallic bonding as both the 3d and 4s subshell electrons are available for delocalisation

3. Electron Configurations  In the neutral transition metal atoms, the 4s subshell is at a lower energy state and so is filled first  NB – Cr and Cu are anomalous because their 4s subshell only contains 1 electron  When the elements form positive ions, the electrons are pulled closer to the nucleus.  This is particularly strong in the 3d subshell and causes 3d to become a lower energy state than 4s  The result of this is that 4s electrons are always lost first!

4. Chemical Properties  There are four distinctive properties of transition metals: 1. They are good catalysts 2. They are coloured compounds 3. They can exist in different oxidation states in compounds 4. They form complex ions

5. 1. Catalysts  They are good catalysts because of their variable oxidation states  This means they may allow an alternative reaction pathway with a lower activation energy  They can also act as solid catalysts because of their partially filled d subshells, which allow dative bonds between reactants and the transition metal surface

6. 1. Catalysts  Transition metals work as catalysts via three steps: 1. Adsorption – weak dative bonds form between reactant molecules and transition metal surfaces therefore weakening the molecular bonds of the reactant 2. Reaction – this occurs faster because the activation energy has been lowered due to the partially broken molecular bonds 3. Desorption – weak dative bonds break

7. 1. Catalysts - Examples  Catalytic converters : Pt/Pd/Rh catalyse the oxidation of CO and reduction of NO in cars  Hydrogenation of alkenes: Ni  Haber Process : finely divided Fe  Contact Process : V 2 O 5

8. 2. Coloured Compounds  Transition metals are coloured due to their partially filled d subshells  d 0 and d 10 configurations such as Sc 3+ and Cu + tend to be colourless

9. 3. Oxidation States  The 3d and 4s subshells are close in energy so using oxidising agents of increasing strength will remove an increasing number of electrons  There are two trends: 1. Sc – Mn: the maximum oxidation state corresponds to all the outer shell electrons being lost 2. Fe – Zn: the maximum oxidation state decreases due to an increased nuclear charge making it harder to remove electrons, so lower oxidation states become more stable

10. 3. Oxidation States Sc+3 Ti+1+2+3+4 V+1+2+3+4+5 Cr+1+2+3+4+5+6 Mn+1+2+3+4+5+6+7 Fe+1+2+3+4+5+6 Co+1+2+3+4+5 Ni+1+2+3+4 Cu+1+2+3 Zn+2 The highlighted oxidation states are the most common and the most stable

11. 3. Oxidation States  From the previous slide we can see that most transition metals have a stable +2 oxidation state  This is due to the relative ease of losing the two 4s electrons  Higher oxidation states do not exist as free ions, but as oxyions e.g. MnO 4 - (Mn existing as +7)  Each oxidation state of each metal has it’s own colour

12. 3. Oxidation States  Redox Reactions:  To convert from one oxidation state to another requires a reducing/oxidising agent of sufficient strength  CROSS REFERENCE: ELECTRODE POTENTIALS  Disproportionation:  A redox reaction in which the same molecule/ion is simultaneously oxidised and reduced  Many less common transition metal ions disproportionate into more common ones such as Cu +

13. 4. Complex Ions  A complex ion is formed when a central metal ion becomes attached to a number of molecules or anions called ligands  The bonds formed between ligands and the central metal ion are co-ordinate bonds  A lone pair of electrons from the ligand are donated into an empty orbital on the central metal ion  The number of co-ordinate bonds formed by the central metal ion is called the co-ordination number e.g. [Co(NH 3 ) 6 ] 3+ Central ion = Co 3+ Ligands = NH 3 Co-ordination no. = 6 Shape = octahedral Bond angle = 90 °

14. 4. Complex Ions  Complex ions may be positive or negative  It is possible to have neutral complex compounds  Ligands in a complex ion need not always be the same  NB H 2 O is a ligand

15. 4. Complex Ions  Ligand Substitution Reactions  These occur because the complex ion formed is more stable than the original ion because it has stronger dative bonds  When the stability of two complex ions is similar, there is an equilibrium; a ligand substitution reaction occurs when there is a higher concentration of the new ligand

16. 4. Complex Ions Shapes of complex ions Co-ordination no. 4 Co-ordination no. 6 Co-ordination no. 2

17. 4. Complex Ions  Polydentate ligands form two or more co-ordinate bonds with the central metal ion  Bidentate ligands have two lone pairs of electrons so can form two co-ordinate bonds with the central metal ion

18. 4. Complex Ions  Isomerism: 1. Cis-Trans; Observed in some octahedral and square planar complexes 2. Optical; most commonly observed in octahedral complexes containing bidentate ligands  Complexes with two bidentate and two monodentate ligands can show both cis-trans and optical isomerism but ONLY the cis isomer has optical isomers

19. The Stability of Complex Ions  This depends on the strength of the co-ordinate bonds formed  It can be measured in terms of the equilibrium constant for its formation from its complex with water – K stab K stab = [reactants] [products] In general, monodentate ligands are less stable than bidentate complexes which are less stable than hexadentate complexes. This is because it is harder for the central ion to break free from a polydentate ligand.

20. Vanadium, V  The stable oxidation states of vanadium are:  +2V 2+ lilac  +3V 3+ green  +4VO 2+ blue  +5VO 2 + yellow  These can be readily interconverted using suitable oxidising or reducing agents  KMnO 4 will oxidise vanadium all the way from +2 to +5  The oxidation state of a V ion can be found by titrating with KMnO 4 /H +

21. Iron, Fe  The stable oxidation states of Iron are:  +2Fe 2+ pale green  +3Fe 3+ yellow brown  In theory [Fe(H 2 O) 6 ] 3+ is violet but ligand substitution with Cl - or OH - gives the more common brown colour

22. Iron ( II )  Precipitation reaction: Fe 2+ (aq) + 2OH - (aq) Fe(OH) 2 (s) pale green solution green precipitate darkens on standing due to oxidation by oxygen in air The amount of Fe2+ present in a solution can be found by titration with acidified potassium permanganate (VII) MnO 4 + 8H + + 5Fe 2+ 5Fe 3+ + Mn 2+ + 4H 2 O purplegreenyellow colourless

23. Iron ( III )  Precipitation reaction: Fe 3+ (aq) + 3OH - (aq) Fe(OH) 3 (s) yellow-brown solution rust coloured precipitate The amount of Fe 3+ present can be found by titration. First Fe 3+ is reduced to Fe 2+ by acidified zinc. The mixture is filtered to remove excess zinc then titrated as before with acidified potassium permanganate (VII)

24. Cobalt, Co  Cobalt (II) reaction with alkali: [Co(H 2 O) 6 ] 2+ (aq) + 2OH - (aq) [Co(OH) 2 ] (s) + 6H 2 O (l)  Cobalt chloride is used as a test for water: [CoCl 4 ] 2- (aq) + 6H 2 O (l) [Co(H 2 O) 6 ] 2+ (aq) + 4Cl - (aq) pink blue

25. Copper, Cu  Copper has two stable oxidation states:  +1Cu + colourless  +2Cu 2+ blue

26. Copper ( I )  Has a full 3d subshell so usually forms colourless compounds  It is unstable in water and will immediately disproportionate: 2Cu + Cu 2+ + Cu (s)

27. Copper ( II )  White when anhydrous because there have to be ligands for the electrons in d orbitals to absorb light  Reaction with alkali (NH3): Cu 2+ + 2OH - Cu(OH) 2 (s)  With excess, blue ppt redissolves to give a deep blue/violet solution: 2H 2 O + Cu(OH) 2 + 4NH 3 [Cu(NH 3 ) 4 (H 2 O) 2 ] 2+ + 2OH -  Also forms a complex with Cl - : [Cu(H 2 O) 6 ] 2+ + 4Cl - [CuCl 4 ] + 6H 2 O deep blue violet blueyellow