15 THE TRANSITION METALS

Transcription

1 15 THE TRANSITION METALS What is the difference between a d-block element and a transition element? Clue: Sc and Zn are not transition elements (see next slide) Write the electronic configurations of the transition elements in period 4.. Sc 1s 2 2s 2 2p 6 3s 2 3p 6 3d 1 4s 2 Ti [Ar] 3d 2 4s 2 Zn [Ar] 3d 10 4s 2 Why are the configurations of Cu and Cr as they are?

2 These elements are all d-block elements because the last electron that they received went into a d-orbital. However, the definition of a transition element depends on its ability to form one or more stable ions with a partially filled d-orbital. This property has consequences that affect the chemistry of these elements. Write the electronic configurations of the stable ions of scandium and zinc

3 Look at p108 and list the characteristics of transition metals and their compounds Are they reactive?

4 Atomic radii, ionisation energies and reactivity The d-orbital is inside the fourth shell, and so the radii of the atoms vary slightly across the 1st series Although there is a slight decrease in radius due to increased nuclear charge, the screening from the inner shells reduces the effects of this (why are there little increases for Mn and Zn?)

5 The first and second ionisation energies vary little, and the differences are not significant. Compare successive IEs of Ca and Fe The effective nuclear charge increases, causing an increase in IE. Note the large jump for Ca. With Fe, the extra ionisation energy is compensated more or less by the extra lattice enthalpy or hydration enthalpy evolved when the 3+ compound is made. The enthalpy of formations of FeCl2 or FeCl3 aren t really that different.

6 The hydration of the ions formed by transition metals and the formation of complex ions and oxoions leads to the variable valency of transition metals. This is because these processes are exothermic and compensate the successive ionisation energies.

7 The chemistry of the elements is less varied than the p-block elements in the same period due to the filling of the inner d-orbitals. The d-orbitals get more involved in bonding across each of the transitions series above but by the end of each series the increases in nuclear charge means that these electrons are held more tightly and are less available for bonding Text book pp

8 COMPLEX IONS These are central metal ion surrounded by neutral molecules or anions called ligands. Ligands always have a lone pair of electrons that they can donate to the central ion and form dative covalent bonds. The total number of dative bonds around the central metal ion is called the COORDINATION NUMBER. Often, 6 bonds are formed in an octahedral arrangement.

9 Monodentate Ligands H 2 O CN - NH 3 OH - aqua cyano ammine hydroxo Ions Fe 2+ /Fe 3+ ferrate Cu 2+ cuprate

10 The overall charge on the ion is the sum of the charges of the central ion and the ligands, and the name given to the complex ion includes the type of ligand and the oxidation state of the central transition metal ion. Eg. The Hexaamminecobalt(III) ion Name the above

11 Other possibilities Name the above

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14 A bidentate ligand Chelation: the formation of multiple dative bonds between a polydentate ligand and the central metal atom/ion Hydrazine cannot do this. Why?

15 EDTA - ethylenediaminetetraacetic acid: A polydentate ligand

16 This is a great example of a multi-dentate ligand

17 Carbon monoxide poisoning Carbon monoxide acts as a better ligand to the central Fe ion and so Ligand Exchange occurs. Text book pp /

18 Transition metal ions and colour Remember that the 3 d-orbitals are not identical:

19 When the ligands form dative covalent bonds with the central ion, the energy of all of the orbitals increases due to the greater electron density in general. However, the orbitals that coincide with the bonds have their energy increased to a greater extent. This can be represented as energy levels

20 Transition metal ions must have a partially filled d-orbital, so that the absorption of light causes an electron to jump up to the higher energy level (excited state), and then return to the ground state releasing electromagnetic radiation which happens to be in the visible range for transition metals. This explains why Zn and Sc are not transition elements, and do not form coloured compounds (although they do form complexes). Cu+ also forms colourless complexes but Cu 2+ no (it forms greens and blues). When white light passes through a solution of a transition metal ion, some wavelengths of light are absorbed and the rest are reflected. The colour wheel helps us to determine which colour will be seen A colour change is observed when: Oxidation number changes Type of ligands change The coordination number changes Text book pp

21 Types of reaction that transition metals undergo Redox reactions Acid-Base Ligand exchange Coordination number change These reactions often cause a change in colour of the metal ion complexes

22 Hydration, deprotonation (acid-base) and ligand exchange The hydrated ions can now be conceived of as the complex ion [Fe(H 2 O) 6 ] 3+ Aqua ions of transition metals are often acidic [Fe(H 2 O) 6 ] 3+ + H 2 O [Fe(H 2 O) 5 OH] 2+ + H 3 O + This is an acid-base or deprotonation reaction. How could we shift this equilibrium to the right? Eqns: The intense and distinct colours of the hydroxides formed can be used to identify the metal ion in question. The same precipitates can be observed when we add aqueous ammonia solution, due to the presence of OH-

23 Some transition metal ions in solution also show further reactions with aqueous alkali and ammonia solutions. If we add a little NaOH(aq) or NH3(aq) to a green(violet) solution of hexaquachromium(iii) ions, a green precipitate of Cr(H 2 O) 3 (OH) 3 (s) forms. In this case, the metal hydroxide formed must be amphoteric as it will react with XS NaOH(aq) and with H+(aq) ions Cr(H 2 O) 3 (OH) 3 (s) + 3H+ (aq) Cr(H 2 O) 3 (OH) 3 (s) + 3OH- (aq) We can shift these equilibria in either direction by adding acid or alkali

24 Ammonia is not a strong enough base to cause the metal hydroxides to show their amphoteric properties, but with Co 2+, Ni 2+ and Cu 2+, a soluble ammine complex is formed Eg: Cu 2+ This final reaction is an example of ligand exchange. Think about what complex ions formed would look like.

25 Water molecules are replaced when a more competitive ligand is added to the solution, eg. Aqueous Copper (II) ions and Conc HCl Similar reactions occur with Cobalt (II) ions

26 Ligand exchange can be explained in terms of stability. The formation of dative covalent bonds is favoured thermodynamically, and the ligands that can donate their lone pair more efficiently will form complexes of lower energy. Other possible factors: Charge on the ligands Size of the ligands Entropy Orbitals available for bonding Complexes are also formed by Be, Mg and Al due to their high charge density, although they are not coloured. Why? Text book pp

27 LEARN! Give the formulae and the colours for the complex ions formed in the following reactions Cu 2+ Fe 2+ Fe 3+ Co 2+ Cr 3+ Formula of Aqueous ion On adding a small amount of NaOH(aq) On adding XS of NaOH(aq) On adding a small amount of NH3(aq) On adding XS of NH3(aq) On adding conc HCl

28 Chromium As well as the ligand exchange and acid-base reactions that we have already seen with Chromium, this element also displays a range of oxidation numbers and so REDOX reactions abound What is the formula of the dichromate(vi) ion? Use your table of SEPs to show what happens if we add powdered Zn to a test tube containing this aqueous ion in acidic solution? Give ionic equations and colours

29 Oxidation of Cr(III) We can use alkaline Hydrogen Peroxide solution to oxidise Chromium(III) Hydroxide to the chromate(vi) ion. Write the ionic equations and show that the reaction is feasible using SEPs What happens if we acidify the resulting solution? Is this redox? Text book pp

30 Vanadium Vanadium is a great example of variable valency and coloured ions Oxidation Ion Colour Redox Equation E o Values Compounds State +5 VO - 3 VO + 2 YELLOW VO 2 +(aq) + 2H + (aq) + e- VO 2+ (aq) + H 2 O(l) V +4 VO 2+ BLUE VO 2+ (aq + 2H + (aq) +e - V 3+ (aq) + H 2 O(l) V +3 V 3+ GREEN V 3+ (aq) + e - V 2+ (aq) V +2 V 2+ VIOLET The +5 oxidation state is oxidising, and the +2 is reducing (it will even reduce water!!). The +4 state is the most stable in the presence of air, and the +3 in its absence.

31 Write the half equations and the overall equation for the mild reduction of V(+5) to V(+4) using sulphite ions in acidic solution. Write the half equations and the overall equation for the reduction of V (+5) by Zn Text book pp

32 TRANSITION METALS AS CATALYSTS They are affected by the reactions they catalyse, they do change form but they are not used up. An alternative route of lower energy is provided for the reaction. There are two types: HOMOGENEOUS (same state as reactants) and HETEROGENEOUS (solids, usually). Transition metals use their varied oxidation states to catalyse other reactions. Eg. 2Ce 4+ (aq) + Tl + (aq) 2Ce 3+ (aq) + Tl 3+ (aq) The reaction is slow, but can be catalysed with aqueous Mn 2+ ions. How?

33 Heterogeneous catalysts work by absorbing the reactants onto their surface (active sites), thus weakening the bonds present and favouring the reaction to form the product which is then released (desorbed) from the surface. This process depends on: the size of the lattice of the catalyst the strength of the bonds formed with the reactants and the products (must be strong and weak respectively) the use of available d-orbitals to form these bonds. Catalysts are very specific for given reactions, for the above reasons.

34 Catalytic Converters Is this a heterogeneous or homogeneous catalyst in this case? What are the reactions taking place? Describe the process What are the usual metals used to make these catalysts? Look at P132 of the textbook and outline how Vanadium(V)Oxide acts as a catalyst in the contact process Text book pp

35 Homogeneous catalysis Catalysts that are in the same phase as the reactants The reaction occurs at a very slow rate, due to the fact that the 2 ions repel each other. Both Fe(II) ions and Fe(III) ions can act as catalysts for the reaction. Try and write the equations for these reactions showing the intermediate species

36 As we have seen before, Manganese in its various oxidation states can act as an autocatalyst in redox reactions. Remember the reaction with ethanedioic acid? 2 MnO4 - (aq ) + 5 H2C2O4(aq ) + 6 H3O + (aq ) --> 2Mn 2+ (aq ) + 10 CO2(aq ) + 14 H2O Considering that the 2 ions involved in this redox reaction are both negative, try to derive equations to show how the Mn 2+ ions that are initially formed (slowly) speed up the reaction by being oxidised to Mn 3+ Text book pp

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