Presentation is loading. Please wait.

Presentation is loading. Please wait.

20.1 The Transition Metals: A Survey 20.2 The First-Row Transition Metals 20.3 Coordination Compounds 20.4 Isomerism 20.5 Bonding in Complex Ions: The.

Similar presentations


Presentation on theme: "20.1 The Transition Metals: A Survey 20.2 The First-Row Transition Metals 20.3 Coordination Compounds 20.4 Isomerism 20.5 Bonding in Complex Ions: The."— Presentation transcript:

1 20.1 The Transition Metals: A Survey 20.2 The First-Row Transition Metals 20.3 Coordination Compounds 20.4 Isomerism 20.5 Bonding in Complex Ions: The Localized Electron Model 20.6 The Crystal Field Model 20.7 The Molecular Orbital Model 20.8 The Biological Importance of Coordination Complexes Chapter 20. Transition Metals and Coordination Chemistry Vanadium metal (center) and in solution as V 2+ (aq), V 3+ (aq), VO 2+ (aq), and VO 2 + (aq), (left to right).

2 Figure 18.1: The periodic table.

3 Figure 12.39: Special names for groups in the periodic table

4 Figure 12.39: Special names for groups in the periodic table (cont’d)

5 Figure 20.1: Transition elements on the periodic table

6 Figure 12.28: The orbitals filled for elements in various parts

7 Some Transition Metals Important to the U.S. Economy and Defense

8 20.1 A Survey of the Transition Metals → Recall the Representative Elements, Groups 1A – 8A: Chemical similarities occur within the vertical groups E.g. LiCl NaCl KCl RbCl CsCl all ionic compounds Large changes in chemistry across a given period as the number of valence electrons changes E.g. Na Mg Al Si P S Cl Ar — increasing metallic character — decreasing ionization energy NaCl MgCl 2 AlCl 3 … PCl 5 Ionic Covalent molecules → Transition Metals Similarities within a given period as well as within a given vertical group.  this huge contrast with the representative elements is due to the fact that the last electrons added to the transition metal elements are inner electrons: d electrons in d – block transition metals f electrons in the lanthanides and actinides

9 –the inner d and f electrons cannot participate in bonding as readily as the s and p electrons. Characteristics of the transition metals –typical metals metallic luster high electrical and thermal conductivities Differences in Physical Properties among the transition metals can be large –E.g. W, tungsten (mp = 3400 ℃ ) vs Hg, mercury (mp < 25 ℃ ) –hard and high strength vs soft Fe, iron and Ti Cu, Au, Ag –ready rxn w/ O 2 to form oxides vs no rxn with O 2 Cr, Ni, Co, Al, Fe Au, Ag, Pt, Pd Ionic compounds with nonmetals –Often more than one oxidation state E.g. FeCl 2 FeCl –the cations are often complex ions, species in which the transition metal ion is surrounded by a number of ligands.

10 Molecular model: The CO(NH 3 ) 6 3+ ion

11 — Ligands are molecules or ions that behave as Lewis bases, i.e. have a lone pair of electrons. Most compounds of the transition metals are colored. — the transition metal ion can absorb visible light. Many transition metal compounds are paramagnetic. — because they contain unpaired electrons  Electron Configurations (See Section 12.13) The 3d orbitals begin to fill after the 4s orbital is complete. e.g. Sc: [Ar]4s 2 3d 1 Ti: 4s 2 3d 2 Y: 4s 2 3d 3 Cr: 4s 1 3d 5 Mn: 4s 2 3d 5 : Cu: 4s 1 3d 10 Zn: 4s 2 3d 10 for most elements of the first-row transition metals 4s 2 3d n has a lower energy than 4s 1 3d n+1 except chromium and copper. The 4s and 3d orbital energies are very similar. (See Table 20.2) 4s 1 3d n+1 exceptions to 4s 2 3d n electron configuration Co 2+, Mn 2+, Cr 3+, Fe 3+, & Ni 2+

12 Table 20.2 Selected Properties of the First-Row Transition Metals

13 Electron configurations of ions of the first-row transition metals — the energy of the 3d orbitals is significantly less than that of the 4s orbital. E.g. Sc: 4s 2 3d 1 Sc 2+ : 3d 1 Ti: 4s 2 3d 2 Ti 3+ : 3d 1 Zn: 4s 2 3d 10 Zn 2+ : 3d 10 — these ions do not have 4s electrons (since the 3d orbitals are lower in energy) Oxidation States and Ionization Energies Various ions formed by losing electrons E.g. Ti → Ti 2+, Ti 3+, Ti 4+ 4s 2 3d 2 most common (See Table 20.2) — to the right of the row the higher oxidation states are not observed because the 3d orbitals become lower in energy as the nuclear charge increases, making electrons difficult to remove. e.g. Zn → Zn 2+ {Zn 3+, Zn 6 +, Zn 10 +, etc ← NOT OBSERVED} 4s 2 3d 10 observed (See Figure 20.2)

14 Figure 20.2: plots of the first (red dots) and third (blue dots) ionization energies for the first-row transition metals

15  Standard Reduction Potentials The potential of the half-reaction M (s) → M n+ + ne - characterizes the reducing ability of the metal. this is the reverse of usually tabulated half-reactions and the potentials are opposite in sign to tabulated values in Table Since by definition ξ o = 0 for: 2H + + 2e - → H 2 all the first-row transition metals, except copper, can reduce H + ions to hydrogen gas in 1M aqueous solutions of strong acids: M (s) + 2H + (aq) → H 2 (g) + M 2+ (aq)

16  The 4d and 5d Transition Series Comparison of the atomic radii of 3d, 4d, and 5d elements See Figure 20.3 — general decrease in size in going from left to right across each series — significant increase in size from 3d to 4d — 4d and 5d metals are very similar in size  this is due to the lanthanide contraction. lanthanide series: elements between lanthanum (La) and hafnium (Hf) ↑ filling of 4f orbitals which are in the interior of the atoms do not affect size of the 5d elements 4d and 5d transition metals, though not as common as 3d metals, have some very useful properties. E.g. The platinum group metals – Ru, Os, Rh, Ir, Pd and Pt – are widely used as catalysts in many industrial processes.

17 Figure 12.28: The orbitals filled for elements in various parts

18 Figure 12.31: The positions of the elements considered in Example 12.8

19 Figure 20.3: Atomic radii of the 3d, 4d, and 5d transition series.

20 Figure 20.1: Transition elements on the periodic table

21 20.2 The First-Row Transition Metals  Highlights of some properties or chemistry of the 10 3d transition metals. 1.Scandium, Sc  +3 oxidation state in compounds, e.g. ScCl 3, Sc 2 O 3, etc  most of its compounds are colorless and diamagnetic. 2.Titanium, Ti  fairly abundant (0.6% by mass of the earth’s crust)  low density + high strength + high mp (1,672 ℃ ) excellent structural material: jet engines, Boeing 747 jetliners, etc.  titanium ( Ⅳ ) oxide (or titanium dioxide), TiO 2 is the most common compound. -- white pigment used in many products: paper, paint, linoleum, plastics, cosmetics, etc. 3. Vanadium, V The most common oxidation state is +5 as in V 2 O 5 (orange, mp = 650 ℃ ) and VF 5.

22 Figure 20.4: Titanium bicycle

23 5. Manganese, Mn –The only member of the 3d metals that can exist in all oxidation states from +2 to +7. –Manganese (VII) ion: MnO 4 - permanganate ion ( a strong oxidizing agent in solution) Ferrochrome added directly to iron in steel making process common oxidation states in compounds: +2, +3, and +6 Cr 2+ (chromous ion) Cr 3+ (chromic ion) chromium (IV) oxide + conc. sulfuric acid cleaning solution 4.Chromium, Cr –Main ore is chromite (FeCr 2 O 4 ) –FeCr 2 O 4 (s) + 4 C (s) → 4 CO (g) + Fe (s) + 2Cr (s)

24 Manganese nodules on the sea floor Source: Visuals Unlimited

25 6.Iron, Fe Quite abundant (4.7% of the earth’s crust) Its chemistry mainly involve its +2 and +3 oxidation states. 7. Cobalt, Co mainly +2 and +3 states compounds with 0, +1 and +4 states are also known 8. Nickel, Ni mainly in the +2 oxidation state. aqueous solutions of nickel (II) salts contain Ni(H 2 O) 6 2+ ion, —characteristic emerald green color

26 Aqueous solution containing the Ni 2 + ion

27 9. Copper, Cu widely available in ores (sulfides, chlorides, carbonates etc.) used in electrical applications (wires, cables, etc) also used in water pipes in homes many common alloys contain copper e.g. brass, bronze, sterling silver, 18- and 14- Karat gold  chemistry involves +2 oxidation state, but also some compounds with the +1 oxidation state. 10. Zinc, Zn +2 oxidation state only used mainly for producing galvanized steel. Zinc (II) salts are colorless.

28 20.3 Coordination Compounds A coordination compound consists of a complex ion and counter ions. –it is an ionic compound, electrically neutral. –complex ion = transition metal ion + attached ligands. E.g. [Co(NH 3 ) 5 Cl]Cl 2 Co(NH 3 ) 5 Cl 2+ ← complex ion 2 Cl - ← counter ions (anions) –coordination compounds ionize in solutions (similar to simple salts) [Co(NH 3 ) 5 Cl]Cl 2 (s) Co(NH 3 ) 5 Cl 2+ (aq) + 2 Cl - (aq) Coordination Number of Metal Ions –The number of bonds formed between a metal ion and the ligands in the complex ion is termed the coordination number. –depending on the size, charge, and electron configuration of the transition metal ion, the coordination number can be from 2 to 8. H2OH2O

29 –many metal ions show more than one coordination number. –for the typical geometries for the various typical coordination numbers see Figure  Ligands –a neutral molecule or ion having a lone pair that can be used to form a bond with a metal ion. Lewis bases by definition are ligands the metal ion is a Lewis acid –a metal – ligand bond is called a coordinative covalent bond. it results from a Lewis acid – base interaction in which a ligand donates an electron pair to an empty orbital on a metal ion.

30 Figure 20.6: Ligand arrangements for coordination numbers 2, 4, and 6


Download ppt "20.1 The Transition Metals: A Survey 20.2 The First-Row Transition Metals 20.3 Coordination Compounds 20.4 Isomerism 20.5 Bonding in Complex Ions: The."

Similar presentations


Ads by Google