2 electrons 2s 2 2p 6. 8 electrons (octet rule) 3s 2 3p 6 3d 10

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1 Main Group and Transition Metal Chemistry: Reading: Moore chapter 22, sections 22.1, 22.6 Questions for Review and Thought: 14, 16, 24, 26, 30, 34, 36, 42, 48, 50, 58, 60. Key ncepts and Skills: definition of bidentate, monodentate, coordination number, lanthanide contraction, chelating ligands, coordinate covalent bond, hexadentate ligand. Be able to write electron configurations for transition metals and their ions; explain trends in the sizes of transition metal atomic radii; explain coordinate covalent bonding in coordination compounds and complexes; understand the different types of isomerism. Lecture Topics: I. Review of the properties of the main group elements (Chapter 7) The Aufbau principle allows one to build up the ground-state electronic configurations of atoms in the periodic table. 1s 2 2 electrons 2s 2 2p 6 8 electrons (octet rule) 3s 2 3p 6 3d electrons 4s 2 4p 6 4d 10 4f 14 5s 2 5p 6 5d 10 5f 14 5g 18 6s 2 6p 6 6d 10 6f 14 6g 18. Principal quantum number n (1-7)- integer assigned to each of the main electron energy levels in an atom. Number increases as we move down the periodic table Angular quantum number, l (l=n-1) refers to the s (l=0), p (l=1), d (l=2), f (l=3) subshells within a shell. Each subshell has (l.0 -l) orbitals. Note: the transition series is filling d orbitals (d block elements); the alkali and alkaline earth metals (groups 1A and 2A, s block elements) are filling s orbitals; groups 3A-8A are filling p orbitals (p block elements); the lanthanides and actinides are filling f orbitals (f-clock elements) There are 1 s orbital, 3 p orbitals, 5 d orbitals and 7 f orbitals. Know the orbital shapes. Spin quantum number: +1/2 and 1/2: each orbital within a subshell can contain up to 2 electrons; if two electrons fill an orbital they must have opposite spins (up and down) (Pauli exclusion principle) Hund s Rule: electrons pair only after each orbital in a subshell is occupied by a single electron. Example: Nitrogen: [He] 2s 2 2p 3 vs. Oxygen: [He] 2s 2 2p 4 What is the ground state electronic configuration of Si, Mg 2+, F -, S 2-, Xe, As, Pb, Ra? Check your knowledge: Si: [Ne] 3s 2 3p 2 What do the unfilled p orbitals look like? Mg 2+ : [Ne] 3s 0 F - : [He] 2s 2 2p 5 What do the unfilled p orbitals look like? S 2- : [Ne] 3s 2 3p 4 What do the unfilled p orbitals look like? Xe: [Kr] 5s 2 (4d 10 )5p 6 As: [Ar] 4s 2 (3d 10 ) 4p 3 What do the unfilled p orbitals look like? Pb: [Xe] 6s 2 (4f 14 ) (5d 10 ) 6p 2 What do the unfilled p orbital look like? Ra: [Rn] 7s 2

2 II. Atomic/ionic radii Atomic radius decreases across a row and increases down the periodic table. Why? Going down the periodic table the principal quantum number increases, and we know that higher quantum numbers correspond to higher energy levels (shells) which are further from the nucleus. Going across the periodic table, as electrons are added to a shell (filling the subshells), each electron avoids other electrons (due to electron-electron repulsion) and thus fails to effectively shield the nuclear charge from other added electrons; each added electron thus feels a greater net nuclear charge (Z eff ), causing the atom to contract and radii to decrease. Note: electrons in p, d, f orbitals are further from the nucleus than electrons in s orbitals and thus are less effective at shielding the nuclear charge mpare radii:al 3+, B 3+, He, Ne, Li +, Be 2+. Electronegativity, a measure of the ability of an atom in a molecule to attract bonding electrons to itself, increases across a period until the noble gases which have very small electronegativities. Why should this be so? Why would you expect electronegativity to decrease down a group in the periodic table? III. Chemical reactivity Group 1A and 2 A metals form ionic compounds since they are relatively electropositive, giving up electrons to obtain a stable noble gas core (octet) Na + Na Mg + 2Br MgBr 2 [Ne]3s 1 [Ne]3s 2 3p 5 [Ne] [Ar] [Ne]3s 2 [Ar]4s 2 4p 5 [Ne][Kr] 2 Hybridization is the mixing of s and p orbitals to maximize the number of unpaired electrons on an atom available for sharing. Sp, sp 2, and sp 3 orbitals are slightly higher in energy than s orbitals and lower in energy than p orbitals. nsider: BF 3,, CH 4,, H 2 O, HF BF 3 : 2s 2 2p 1 sp 2 p 0 ; boron has three sp 2 orbitals each with one unpaired electron (available for sigma bonding with an electron from each F atom) and one empty p orbital. CH 4 : 2s 2 2p 2 sp 3 ; carbon has four equivalent sp 3 orbitals each with one unpaired electron (available for sigma bonding with an electron from each H atom). : 2s 2 2p 3 sp 3 ; nitrogen has four equivalent sp 3 orbitals, three with one unpaired electron (available for sigma bonding with an electron from each H atom) and one with a lone pair. H 2 O: 2s 2 2p 4 sp 3 ; oxygen has four equivalent sp 3 orbitals, two with one unpaired electron (available for sigma bonding with an electron from each H atom) and two with lone pairs. HF: 2s 2 2p 5 sp 3 ; fluorine has four equivalent sp 3 orbitals, one with an unpaired electron (available for sigma bonding with an electron from an H atom) and three with lone pairs. Geometry: sp 3 tetrahedral; sp 2 trigonal planar; sp linear. Hybrid (sp) orbitals participate in sigma (σ) bonding; p-orbitals participate in pi (π) bonding IV. Transition Metals Electron configurations Filling order: 4s 2 3d 10 4p 6 ; energies of 4s and 3d levels are quite close despite the different quantum numbers. Ordinarily 4s fills before 3d.

3 K: [Ar]4s 1 Ca: [Ar] 4s 2 Sc: [Ar] 4s 2 3d 1 Ti: [Ar] 4s 2 3d 2 V: [Ar] 4s 2 3d 3 Cr: [Ar] 4s 1 3d 5 Mn: [Ar] 4s 2 3d 5 Fe: [Ar] 4s 2 3d 6 : [Ar] 4s 2 3d 7 Ni: [Ar] 4s 2 3d 8 Cu: [Ar] 4s 1 3d 10 note: spin pairing in the 4s orbital is a higher energy situation than having all spins parallel note: upon filling up with electrons, 3d subshell is lower in energy than 4s Zn: [Ar] 4s 2 3d 10 Thus, Cu, Ag, Au have filled d 10 orbitals, just like Zn, Cd, and Hg. For Cr, two half-filled subshells (4s 1 3d 5 ) is more stable than 4s 2 3d 4 When a transition metal loses its electrons to form an ion, the s electrons are lost before the d electrons; example: Ti 3+ [Ar]4s 0 3d 1 ; Cr 2+ [Ar] 4s 0 3d 4 ; Cu +1 [Ar] 4s 0 3d 10 b. How do we know that s electrons are lost first? Magnetic measurements confirm paramagnetism. Substances that have unpaired spins are paramagnetic and are attracted into a magnetic field because of alignment of some unpaired spins with the magnetic field. Diamagnetic substances have all electrons spin-paired so their magnetic fields effectively cancel one another; they therefore experience negliglible attraction to a magnetic field. The extent of attraction into a magnetic field, measured by the apparent mass of the sample in the field, is an indication of the number of unpaired electrons. nsider: Fe 2+ : [Ar] 4s 0 3d 6 (four unpaired spins) vs. Fe 3+ : [Ar] 4s 0 3d 5 (5 unpaired spins) Alternatively, removing d electrons first: Fe 2+ : [Ar] 4s 2 3d 4 (four unpaired spins) vs. Fe 3+ :4s 2 3d 3 (3 unpaired spins). It is found that Fe 3+ has a greater apparent mass in a magnetic field than Fe 2+ c. Atomic radii of the transition elements Atomic radii decrease across a period due to the increasing effective nuclear charge felt by the valence electrons (Z eff, see above). However, down a group an interesting trend is observed: Ti<Zr =~ Hf; V<Nb =~ Ta; Cr<Mo=~ W. Why? Beginning at Lanthanum, the 4f subshell starts to fill. In the lanthanide series of elements, effective nuclear charge builds up, causing a decrease in atomic radii because all additional electrons go into 4f orbitals, which are so far from the nucleus that they do not screen other valence electrons from the nuclear charge. The increased nuclear charge pulls the valence electrons closer to the nucleus, offsetting the size increase on going from quantum number 5 to 6. This phenomenon is known as the Lanthanide contraction d. Oxidation states Multiple oxidation states are possible for most of the transition elements: remove s electrons first, then paired d electrons, then unpaired d electrons. The maximum oxidation state for the first five elements of a transition series is the sum of the ns and nd electrons of the uncharged atom: Ti: [Ar] 4s 2 3d 2 4+, 3+, 2+ V: [Ar] 4s 2 3d 3 5+, 4+, 3+, 2+; Cr: [Ar] 4s 1 3d 5 6+, 5+, 4+, 3+, 2+ Mn: [Ar] 4s 2 3d 5 7+, 6+, 5+, 4+, 3+, 2+

4 Note: compounds in high oxidation states are good oxidants and form covalent complex ions: example: MnO 4 -, Cr 2 O 7 2- mpounds in low oxidation states form ionic compounds. Examples: Fe 2, Fe 3, Mn 2 Electron gain and electron loss: Fe 2+ Fe 3+ + e - ; Fe 3+ + e - Fe 2+ V. Bonding in transition metal complexes. The most common form of bonding in transition metal complexes is coordinate covalent bonding, in which one atom contributes both electrons for the shared pair in a bond. Typical example: F 3 B NH 3 complex. The lone pair on nitrogen donates to the mepty p orbital on boron. Transition metals have incompletely filled d shells that can accept lone pairs. The molecules bonded to the central metal atom are ligands; ligands may be charged (CN -, - ) or uncharged (H 2 O, NH 3, (CH 3 ) 3 P) nsider: [Ni(OH 2 ) 6 ] 2+ and [Ni() 4 ] 2-, both are complex ions- formed when several molecules or ions are connected to a central metal ion or atom by coordinate covalent bonds. mpounds in which complex ions are combined with oppositely charged ions to form neutral compounds are called coordination compounds: [Ni(H 2 O) 6 ] The number of coordinate covalent bonds between the ligands and the central metal ion is the coordination number (CN) of the metal ion CN=2 examples: [Ag(NH 3 ) 2 ] +, [Au 2 ] - CN=4 examples: [Ni 4 ] 2-, [Pt(NH 3 ) 4 ] 2+ CN=6 examples: [Fe(H 2 O) 6 ] 2+, [(NH 3 ) 6 ] 3+ Ligands that can form only one coordinate covalent bond to a metal are monodentate ligands. Examples: H 2 O, NH 3, H 2 S, HO -, -, NC -, CO Some compounds can form two or more coordinate covalent bonds to the same metal ion because they possess two or more atoms with lone-pairs separated by several intervening atoms. These are polydentate ligands. Examples: bidentate ligands- ethylenediamine: H 2 NCH 2 CH 2 NH 2 2- Carbonate ion: CO 3 Oxalate ion: - - O 2 CCO 2 A Hexadentate ligand can form a claw-like complex or chelate with a metal ion, occupying all possible ligand sites (for an octahedral complex) with lone pairs from a single ligand molecule. Example: EDTA ( - O 2 CCH 2 ) 2 N CH 2 CH 2 N(CH 2 CO 2- ) 2 EDTA complexes metal ions so strongly that it is used to treat lead and mercury poisoning. VI. Geometry of coordination compounds (see appendix below) The geometry of coordination compounds is dictated by the arrangement of electrondonor atoms of the ligands attached to the central metal ion. nsider the situations with monodentate ligands: CN=2 linear geometry (bond angle 180 ) : Ag + NH 3 ; also [Cu 2 ] -, [Au(CN) 2 ] - CN=4 either tetrahedral (bond angles 109 ) or square planar (bond angles 90 ) Tetrahedral: [Zn(NH 3 ) 4 ] 2+ Square planar: [Pt (NH 3 ) 4 ] 2+ difference: sp 3 vs. sp 2 d hybr. CN=6 octahedral (bond-angles are 90 and 180 ): [(NH 3 ) 6 ] 3+, [Fe(H 2 O) 6 ] 2+ (sp 3 d 2 )

5 VII. Isomerism nstitutional isomers are molecules with the same molecular formula but different arrangements of atoms 3 different types : Linkage isomers: a ligand can bond to a metal atom using two different electron-donating atoms: (H 2 O) 5 :SCN: vs. (H 2 O) 5 :NCS: Geometric isomerism: positional isomerism that occurs both in square planar and octahedral complexes: cis vs. trans isomers: Examples: square planar- Pt(NH 3 ) 2 2 NH 3 Pt Pt NH 3 cis trans + Ocatahedral: (NH 3 ) 4 2 NH 3 NH 3 NH 3 and (NH 3 ) 3 3 NH 3 N H 3 NH 3 trans cis cis cis Optical isomerism: enantiomeric compounds are mirror images. Enantiomers are nonsuperimposable molecules which have mostly identical properties but can be distinguished by enzymes containing chiral receptor sires or binding pockets. Optical isomerism is not possible in square planar or linear geometries but is possible in tetrahedral and octahedral geometries: OH 2 OH 2 NH 3 OH 2 H 2 O NH 3 mirror

6 Additonal problems: 1. Will methylamine be a mono- or bidentate ligand? With which of its atoms will it bind to a metal ion? 2. Show how glycinate ion (H 2 N CH 2 COO - ) can act as a bidentate ligand. Wbhich atoms in the glycinate ion will bind to a metal ion? 3. Draw the structures of all possible isomers for the following complexes. Indicate which isomers are enantiomer pairs. (a) (NH 3 ) 2 Br Pt(II) (square planar) (b) [(H 2 O) 2 (CN) 3 (III)] - (octahedral) (c) [( - O 2 CCO 2 - ) 3 V(III)] 3- (octahedral) 4. Iron (III) forms octahedral complexes. Sketch the structures of all the distinct isomers of [Fe(en) 2 2 ] +, indicating which pairs of structures are mirror images of each other. en= ethylenediamine = H 2 NCH 2 CH 2 NH 2 5. nsider the ions Mn 3+, Fe 3+, 3+, Ni 3+. Which ion in its ground state has the greatest number of unpaired electrons? 6. Explain why the atomic radius of Rh is essentially equal to that of Ir. 7. How many isomers would you expect for the complex [Ni(NH 3 ) 3 3 ] - if the coordination geometry is octahedral? Draw the isomers.

7 Appendix: hybridization in transition metal complexes and ligand geometry about a metal. nsider several cases: 1.) linear geometry for [Cu 2 ] -, [Au(CN) 2 ] -, Ag(NH 3 ) 2 + : All of these metals are +1 charge. The electron configuration on the metal in each case is 4s 0 3d 10 4p 0. Since all d-orbitals are filled, to make empty orbitals available for coordinate covalent bonding, hybridization must take place between s and p orbitals only. To accept two ligands, you need only hybridize one s and one p orbital to make two sp hybrid orbitals which are empty and capable of accepting lone pairs from ligands. sp hybridization is always indicative of a linear geometry! Note there are also 5 completely filled (unhybridized) d orbitals on the metal center. 2). Tetrahedral geometry: [Zn(NH 3 ) 4 ] 2+ Zn 2+ has an electron configuration 4s 0 3d 10 4p 0. To accept lone pairs from four ligands, zinc must open up four empty hybrid orbitals. The only way to do this, since all of the d orbitals are filled, is to hybridize one s orbital and all three p- orbitals to make four equivalent empty sp 3 orbitals. sp 3 hybridization is always indicative of tetrahedral geometry! Note there are also 5 completely filled (unhybridized) d orbitals on the metal center. 3.) Square planar geometry: [Pt (NH 3 ) 4 ] 2+ Pt 2+ has an electron configuration 4s 0 3d 8 4p 0. To accept lone pairs from four ligands, platinum must open up four empty hybrid orbitals. hybridization is easier for orbitals which are closer in energy, and certainly 4s and 3d orbitals are close in energy. Since there are 8 3d electrons, spin-pairing all of these electrons in four d orbitals leaves one empty d orbital which can hybridize with one s orbital and 2 p orbitals, giving a total of four empty sp 2 d hybrid orbitals. sp 2 d hybridization is always indicative of a square-planar geometry. Note there are also 4 completely filled (unhybridized) d orbitals on the metal center. 4.) Trigonal bipyramidal: Cu 5 2-, Fe(CO) 5 Cu 3+ has an electron configuration 4s 0 3d 8 ; Fe has an electron configuration 4s 2 3d 6. To accept five lone pairs from five ligands, each metal must make available 5 empty hybrid orbitals. For copper, the eight d electrons can be spin-paired in 4 d orbitals, leaving one d orbital available for hybridization with one s orbital an all three p orbitals, giving 5 empty sp 3 d hybrid orbitals. For iron, 2 4s electrons can easily be promoted to a 3d orbital, and now the 8 d electrons can spin-pair in 4 d orbitals, leaving one d orbital available for hybridization with the now empty s orbital and all three p orbitals to make five empty sp 3 d hybrid orbitals. sp 3 d hybridization is always indicative of trigonal bipyramid geometry! Note there are also 4 completely filled (unhybridized) d orbitals on each metal center.

8 5.) Octahedral: [(NH 3 ) 6 ] 3+, [Fe(H 2 O) 6 ] 2+ Fe 3+ has an electron configuration of 4s 0 3d 6 4p 0 ; 3+ has an electron configuration of 4s 0 3d 6 4p 0 In both cases, 6 empty orbitals must be available to accept 6 lone pairs from the ligands. The 6 d electrons can be spin-paired in three d orbitals, leaving two empty d orbitals available for hybridization. mbining one s, two d, and three p orbitals results in six equivalent, empty sp 3 d 2 hybrid orbitals. sp 3 d 2 hybridization is always indicative of octahedral geometry! Note there are also 3 completely filled (unhybridized) d orbitals on each metal center. 6.) Dodecahedral: Mo(CN) 8 4- Mo 4+ has an electron configuration of 4s 0 3d 2 4p 0. Mo must open up 8 empty orbitals to accept 8 lone-pairs from the ligands. The 2 d-electrons can be spinpaired in a single d orbital, leaving 4 empty d orbitals available for hybridization. mbining one s, 4 d, and 3 p orbitals gives 8 equivalent, empty sp 3 d 4 hybrid orbitals. sp 3 d 4 hybridization is always indicative of dodecahedral geometry! Note there is also 1 completely filled (unhybridized) d orbital on the metal center

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