Part 2 (HL)

1. Part 2 (HL) THE d BLOCK SS CI 11.5 The d block

2. The d block: • The d block consists of three horizontal series in periods 4, 5 & 6 • 10 elements in each series • Chemistry is “different” from other elements • Special electronic configurations important • Differences within a group in the d block are less sharp than in s & p block • Similarities across a period are greater

3. d-Block Elements • Remember that something odd happens with electron configurations after argon. • At argon, the 3s and 3p levels are full, but rather than fill up the 3d levels next, the 4s level fills instead for potassium and then calcium. • You don’t see any filling of d orbitals until you get to the d-block.

4. Electronic Configuration • ScandiumSc, TitaniumTi, VanadiumV, ChromiumCr, ManganeseMn, IronFe, CobaltCo, NickelNi & Copper Cu • Across the 1st row of the d block (Sc to Zn) each element • has 1 more electron and 1 more proton • Each “additional” electron enters the 3d sub-shell • The core configuration for all the period 4 transition elements is that of Ar • 1s22s22p63s23p6

5. Most transition elements are found combined with other elements in ores. • The Iron Triad • Three elements in period 4—iron, cobalt, and nickel—have such similar properties that they are known as the iron triad.

6. Iron triad • Industrial magnets are made from an alloy of nickel, cobalt, and aluminum. • Nickel is used in batteries along with cadmium. • Iron is a necessary part of hemoglobin, the substance that transports oxygen in the blood. • Iron also is mixed with other metals and with carbon to create a variety of steels with different properties.

7. Uses of Transition Elements • Most transition metals have higher melting points than the representative elements. • The filaments of lightbulbs are made oftungsten • Tungsten has the highest melting point of any metal (3,410°C) and will not melt when a current passes through it.

8. Mercury, which has the lowest melting point of any metal (–39°C), is used in thermometers and in barometers. • Mercury is the only metal that is a liquid at room temperatures. • Like many of the heavy metals, mercury is poisonous to living beings.

9. Many other transition elements combine to form substances with brilliant colors. • Chromium’s name comes from the Greek word for color, chrome. • palladium, platinum, iron, vanadium, nickel, cobalt etc are used as catalysts

10. Electron configuration d block Ar 1s2 2s2 2p6 3s2 3p6 SS CI 11.5 The d block

11. Energy 4p 3d 4s 3p 3s 2p 2s Ar 1s2 2s2 2p6 3s2 3p6 1s

12. Energy 4p 3d 4s 3p 3s 2p 2s Sc 1s2 2s2 2p6 3s2 3p6 3d1 4s2 1s

14. Chromium and Copper • Cr and Cu don’t fit the pattern of building up the 3d sub-shell, why? • In the ground state electrons are always arranged to give lowest total energy • Electrons are negatively charged and repel each other • Lower total energy is obtained with e- half filled or fully filled in orbitals (3d) rather than if they are paired in an orbital (4s) • Energies of 3d and 4s orbitals very close together in Period 4

15. Chromium and Copper • Atom of Cr • Orbital energies such that putting one e- into each 3d and 4s orbital gives lower energy than having 2 e- in the 4s orbital. Then the atoms have half filled d orbital (stable) • Atom of Cu • Putting 2 e- into the 4s orbital would give a higher energy than filling the 3d orbitals; so filled d orbital(stable)

16. Energy 4p 3d 4s 3p 3s 2p 2s Cr 1s2 2s2 2p6 3s2 3p6 3d5 4s1 1s

17. Energy 4p 3d 4s 3p 3s 2p 2s Cu 1s2 2s2 2p6 3s2 3p6 3d10 4s1 1s

18. What is a transition metal? • Transition metals [TM’s] have characteristic properties when they are in ionic form. • e.g. coloured compounds, variable oxidation states • These are due to presence of an inner incomplete d orbitals. Electrons from both inner d orbitals and outer s orbital can be involved in compound formation. • So transition metals must have incomplete d orbital in their most stable ionic form.

19. What is a transition metal? • Not all d block elements in ionic form have incomplete d sub-shells • e.g. Zn has e.c. of [Ar]3d104s2, the Zn2+ ion ([Ar] 3d10) has no partially filled d orbitals. So Zinc is not a typical TM. • Similarly Sc forms Sc3+ which has the stable ionic structure [Ar]. Sc3+ has no 3d electrons. So Scandium is also not a transition metal.

20. What is a transition metal? • In period 4 only Ti-Cu are Transition Metals! • Note that when d block elements form ions the s electrons are lost first.

21. What are Transition Metals like? • TM’s are metals: • They are similar to each other but different from s block metals eg Na and Mg • Properties of TM’s: • Dense metals • Have high Tm and Tb • Tend to be hard and durable • Have high tensile strength • Have good mechanical properties

22. What are Transition Metals like? • Properties derive from strong metallic bonding • TM’s can release e- into the pool of mobile electrons from both outer and inner shells • Strong metallic bonds formed between the mobile pool and the +ve metal ions • Enables widespread use of TMs! • Alloys very important: inhibits slip in crystal lattice usually results in increased hardness and reduced malleability

23. Effect of Alloying on TM’s

24. TM Chemical Properties • Typical chemical properties of the TM’s are • Formation of compounds in a variety of oxidation states • Catalytic activity of the elements and their compounds • Strong tendency to form complexes • Formation of coloured compounds

25. Find solution for the following • What is the meaning of oxidation state? • Look at the diagram and state the pattern of oxidation state among the period4, d block elements. • Suggest a reason for the pattern in the above statement • From the above statement conclude why Sc and Zn are not TM

26. Variable oxidation numbers

27. Key Observations • All transition metals show 2+ & 3+ (except for Cu, the 1+) • There is a stepwise progression across the period • Mn has the max • Oxidation states of 3+ or higher show covalent character • Compounds with these higher state will act as oxidizers (K2Cr2O7)

28. Explanation of the variable oxidation number • Calcium only shows the +2 oxidation state in ionic compounds. • Iron can show +2 or +3. • This is because the successive ionization energies of transition metals (d fillers) are much closer together than those of other metals.

29. Oxidation States of TM’s • No of OS’s shown by an element increases from Sc to Mn • In each of these elements highest OS is equal to no. of 3d and 4s e- • After Mn decrease in no. of OS’s shown by an element • Highest OS shown becomes lower and less stable • Seems increasing nuclear charge binds 3d e- more strongly, hence harder to remove

30. Stability of OS’s • General trends (continued) • Relative stability of +2 state with respect to +3 state increases across the series • For compounds early in the series, +2 state highly reducing • E.g. V2+(aq) & Cr2+(aq) strong reducing agents • Later in series +2 stable, +3 state highly oxidising • E.g. Co3+ is a strong oxidising agent, Ni3+ & Cu3+ do not exist in aqueous solution.

31. Complex ion formation • A complex ion has a metal ion at its centre with a number of other molecules or ions surrounding it. • The surrounding molecules or ions are called Ligands. • Ligands are attached to the central ion by co-ordinate (dative covalent) bonds. 

32. Complex ions • So ligands should have a lone pair of electrons. • Transition metal ions can form complex ions due to unfilled s,p or d orbitals. • The ligands form dative bonds with the unfilled orbitals

33. The nature of Ligands • Simple ligands include water, ammonia and chloride ions. • What all these have got in common is active lone pairs of electrons in the outer energy level. • These are used to form co-ordinate bonds with the metal ion.

34. Fe(H2O)63+ • Iron has the electronic structure 1s22s22p63s23p63d64s2 • Example: Formation of: • When it forms an Fe3+ ion it loses the 4s electrons and one of the 3d electrons to leave 1s22s22p63s23p63d5 • Looking at this as electrons-in-boxes, at the bonding level:

35. Now, be careful! The single electrons in the 3d level are NOT involved in the bonding in any way. • Instead, the ion uses 6 orbitals from the 4s, 4p and 4d levels to accept lone pairs from the water molecules. • Before they are used, the orbitals are re-organised (hybridised) to produce 6 orbitals of equal energy.

36. Co-ordination number • Co-ordination number of the metal ion in a complex ion is the number of bonds(ligands) attached to it. • Because the iron is forming 6 bonds, the co-ordination number of the iron is 6.

37. CuCl42- • This is a simple example of the formation of a complex ion with a negative charge. • Copper has the electronic structure • 1s22s22p63s23p63d104s1 • When it forms a Cu2+ ion it loses the 4s electron and one of the 3d electrons to leave • 1s22s22p63s23p63d9

38. CuCl42- • To bond the four chloride ions as ligands, the empty 4s and 4p orbitals are used (in a hybridised form) to accept a lone pair of electrons from each chloride ion. • Because chloride ions are bigger than water molecules, you can't fit 6 of them around the central ion - that's why you only use 4.

39. CuCl42- • The ion carries 2 negative charges overall. • That comes from a combination of the 2 positive charges on the copper ion and the 4 negative charges from the 4 chloride ions. • In this case, the co-ordination number of the copper is, of course, 4.

40. Other complex ion examples • Most complex ions have either 6 ligands arranged octahedrally around the central atom (water and ammonia ligands) or 4 ligands arranged tetrahedrally (chloride ion ligands) or some with 2 ligands arranged linear. (any ligands)

42. Complex ions • Complex ions can be positive(cation) or negative(anion) ions. • Both are soluble in water and pass electricity. • Some complex ions are neutral because the charge on TM and ligands cancel.(insoluble) • Example: [Pt(NH3)2Cl2] • Chloride ions are strongly bonded to central ion, aq. AgNO3 does not give ppt.

43. Catalytic Activity • TM’s and their compounds are effective and important catalysts • Industrially and biologically!! • Once again, • availability of 3d and 4s e- • ability to change OS • among factors which make TM’s such good catalysts

44. Heterogeneous Catalysis • Catalyst in different phase from reactants • Usually means solid TM catalyst with reactants in liquid or gas phases • TM’s can • use the 3d and 4s e- of atoms on metal surface to from weak bonds to the reactants. • Once reaction has occurred on TM surface, these bonds can break to release products • Important example is hydrogenation of alkenes using Ni or Pt catalyst

45. Heterogeneous Catalysis

46. Homogeneous Catalysis • Catalyst in same phase as reactants • Usually means reaction takes place in aqueous phase • Usually involves • TM ion forming intermediate compound with one or more of the reactants • Intermediate then breaks down to form products

47. Questions need explanation • What is the meaning of complementary colour? • What kind of interaction takes place when ligands join to a TM ion? • Why d orbitals spit in to two groups? • What happens to the electrons for the formation of colour? • What is the difference in the octahedral and tetrahedral complex colour formation at orbital level? SS CI 11.5 The d block

48. Coloured compounds • The diagram shows an approximation to the spectrum of visible light.

49. Why is copper(II) sulphate solution blue? • If white light (ordinary sunlight, for example) passes through copper(II) sulphate solution, some wavelengths in the light are absorbed by the solution. • Copper(II) ions in solution absorb light in the red region of the spectrum • Light comes the other side will have all the colours in it except for the red. We see this mixture of wavelengths as pale blue.

50. The colour you would see after absorption of colours is the complementary colours of the absorbed colours. • Colours directly opposite each other on the colour wheel are said to be complementary colours.