Electron Count Oxidation State Coordination Number

1. Electron CountOxidation StateCoordination Number Basic tools for understandingstructure and reactivity. Doing them should be “automatic”. Not always unambiguous Þ don’t just follow the rules, understand them! Counting Electrons

2. The basis of counting electrons • Every element has a certain number of valence orbitals: 1 (1s) for H 4 (ns, 3´np) for main group elements 9 (ns, 3´np, 5´(n-1)d) for transition metals s px py pz dxy dxz dyz dx2-y2 dz2 Counting Electrons

3. The basis of counting electrons • Every orbital wants to be “used", i.e. contribute to binding an electron pair. • Therefore, every element wants to be surroundedby 2/8/18 electrons. • The strength of the preference for electron-precise structures depends on the position of the element in the periodic table. Counting Electrons

4. The basis of counting electrons • Too few electrons: An empty orbital makes the compound very electrophilic,i.e. susceptible to attack by nucleophiles. • Too many electrons: There are fewer covalent bonds than one would think (not enough orbitals available). An ionic model is required to explain part of the bonding. The "extra" bonds are relatively weak. • Metal-centered (unshared) electron pairs: Metal orbitals are fairly high in energy. A metal atom with a lone pair is a strong s-donor (nucleophile) and susceptible to electrophilic attack. Counting Electrons

5. Use a localized (valence-bond) modelto count electrons H2 Every H has 2 e. OK CH4 H has 2 e, C 8. OK NH3 N has 8 e. Nucleophile! OK Counting Electrons

6. C2H4 C has 8 e. OK singlet CH2 C has only 6 e, and an empty pz orbital: extremely reactive ("singlet carbene"). Unstable. Sensitive to nucleophiles and electrophiles. triplet CH2 C has only 6 e, is a "biradical" and extremely reactive ("triplet carbene"), but not especially for nucleophiles or electrophiles. Counting Electrons

7. CH3+ C has only 6 e, and an empty pz orbital: extremely reactive. Unstable. Sensitive to nucleophiles. CH3- C has 8 e, but a lone pair. Sensitive to electrophiles. Cl- Cl has 8 e, 4 lone pairs. OK Somewhat sensitive to electrophiles. Counting Electrons

8. BH3 B has only 6 e, not stable as monomer,forms B2H6: B2H6 B has 8 e, all H's 2 (including the bridgingH!). 2-electron-3-center bonds! OK AlCl3 Al has only 6 e, not stable as monomer,forms Al2Cl6: Al2Cl6 Al has 8 e, all Cl's too (including thebridging Cl!). Regular2-electron-2-center bonds! OK Counting Electrons

9. 2 MeAlCl2® Me2Al2Cl4 2-electron-3-center bonds are a stopgap! H3B·NH3 N-B: donor-acceptor bond (nucleophile NH3 has attacked electrophile BH3). Organometallic chemists are "sloppy" and write . Writing or would be more correct (although the latter does not reflect the “real” charge distribution). Counting Electrons

10. PCl5 P would have 10 e, but only has 4 valence orbitals, so it cannot form more than 4 “net” P-Cl bonds.You can describe the bonding using ionic structures (hyperconjugation). Easy dissociation in PCl3 en Cl2. HF2- Write as FH·F-, mainly ion-dipole interaction. Counting Electrons

11. How do you count? • Number of valence electrons(from periodic table) • Correct for charge, if any(only if it belongs to that atom!) • Count 1 e for every covalent bond to another atom • Count 2 e for every dative bond from another atom • Add Counting Electrons

12. Examples: counting electrons Pd = 10 - = 1 3´Cl = 3 1´NH3 = 2 tot = 16 could have additional 2 e(Pd-Cl p-bond?) B = 3 - = 1 4´H = 4 tot = 8 OK C = 4 1´=O = 2 2´H = 2 tot = 8 OK Ru = 8 2´Cl = 2 2´PMe3 = 4 1´CH2 = 2 tot = 16 could have additional 2 e Counting Electrons

13. Counting is not always trivial Pd = 10 2´- = 2 3´Cl = 3 1´CH2 = 1 tot = 16 could have additional 2 e Counting Electrons

14. Remember, when counting: • Odd electron counts are rare. • In reactions you nearly always go from even to even (or odd to odd), and from n to n-2, n or n+2. • Electrons don’t just “appear” or “disappear”. • The optimal count is 2/8/18 e. 16 e also occurs frequently, other counts are much more rare. Counting Electrons

15. Oxidation States Most elements have a clear preference for certain oxidation states. These are determined by (a.o.) electronegativity and the number of valence electrons: Li: nearly always +1.Has only 1 valence electron, so cannot go higher. Is very electropositive, so doesn’t want to go lower. Cl: nearly always -1.Already has 7 valence electrons, so cannot go lower.Is very electronegative, so doesn’t want to go higher. Counting Electrons

16. Calculating theformal oxidation state • Start with the formal charge on the metal • Ignore dative bonds • Ignore bonds between atoms of the same element (this one is a bit silly) • Assign every covalent electron pair to the most electronegative element in the bond: this produces + and – charges (usually + at the metal) • Add Counting Electrons

17. Examples: oxidation states charge Al = -1 4´Al-Cl: Al+-Cl- = +4 tot = +3 charge Pd = -2 4´Pd-Cl: Pd+-Cl- = +4 tot = +2 charge C = 0 4´C-Cl: C+-Cl- = +4 tot = +4 charge C = 0 2´C-Cl: C+-Cl- = +2 1´C=O: C2+-O2- = +2 tot = +4 charge Mn = -1 4´Mn=O: Mn2+-O2- = +8 tot = +7 Counting Electrons

18. Examples: oxidation states charge Pt = -2 3´Pt-Cl: Pt+-Cl- = +3 tot = +1 univalent Pt ? charge C = 0 3´C-Cl: C+-Cl- = +3 tot = +3 trivalent carbon ? charge Mg = 0 4´Mg-Me: Mg+-Me- = +4 tot = +4 impossible, Mg has only 2 valence electrons! Counting Electrons

19. The significance ofan oxidation state ? Oxidation states are formal. However, they do give an indication whether a structure or composition is reasonable (apart from the M-M complication). Counting Electrons

20. Acceptable oxidation states For group n or n+10: • never >+n or <-n (except group 11: frequently +2 of +3) • usually even for n even, odd for n odd • usually ³ 0 for metals • usually +n for very electropositive metals • usually 0-3 for 1st-row transition metals of groups 6-11, often higher for 2nd and 3rd row • electronegative ligands (F,O) stabilize higher oxidation states, p-acceptor ligands (CO) stabilize lower oxidation states • oxidation states usually change from m to m-2, m or m+2 in reactions Counting Electrons

21. Coordination number Simply the number of atoms directly bonded to the atom you are interested in, regardless of bond orders etc. CH4: 4 C2H4: 3 C2H2: 2 AlCl4-: 4 Me4Zn2-: 4 OsO4: 4 B2H6: 4 (B) 1 (terminal H) 2 (bridging H) Counting Electrons

22. Coordination Number For complexes with p-system ligands, the whole ligand is usually counted as 1: Cyclopentadienyl groups are sometimes counted as 3,because a single Cp group can replace 3 individual ligands: C.N. 4 Counting Electrons

23. Coordination Number The most common coordination numbers for organometallic compounds are: 2-6 for main group metals 4-6 for transition metals Coordination numbers >6 are relatively rare. So are very low coordination numbers (<4) together with a “too-low” electron count. Counting Electrons

24. Coordination number and coordination geometry C.N. "Normal" geometry 2 linear or bent 3 planar trigonal, pyramidal, "T-shaped" 4 square planar, tetrahedral 5 square pyramid, trigonal bipyramid 6 octahedron Counting Electrons

25. Illustration:protonation of WH6(PMe3)3 Could WH6(PMe3)3 be ? Count W: 18 VE (OK), oxidation state 6 (OK), coordination number 9 (very high). Possible. Protonation gives WH7(PMe3)3+. Could that be ? Count W: 18 VE (OK), oxidation state 8 (too high), coordination number 10 (extremely high). W+ must form 7 covalent bonds using only 5 electrons. That will not work! Counting Electrons

26. Exercises Give electron count and oxidation state for the following compounds. Draw conclusions about their (in)stability. Me2Mg Pd(PMe3)4 MeReO3 ZnCl4 Pd(PMe3)3 OsO3(NPh) ZrCl4 ZnMe42- OsO4(pyridine) Co(CO)4- Mn(CO)5- Cr(CO)6 V(CO)6- V(CO)6 Zr(CO)64+ PdCl(PMe3)3 RhCl2(PMe3)2 Ni(PMe3)Cl4 Ni(PMe3)Cl3 Ni(PMe3)2Cl2 Counting Electrons