1. 6/26/2015Counting Electrons1 Electron Count Oxidation State Coordination Number Basic tools for understanding structure and reactivity. Doing them should be “automatic”. Not always unambiguous  don’t just follow the rules, understand them!

2. 6/26/2015Counting Electrons2 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 Every orbital wants to be “used", i.e. contribute to binding an electron pair. Therefore, every element wants to be surrounded by 2/8/18 electrons. The strength of the preference for electron-precise structures depends on the position of the element in the periodic table. The basis of counting electrons

3. 6/26/2015Counting Electrons3 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  -donor (nucleophile) and susceptible to electrophilic attack. The basis of counting electrons

4. 6/26/2015Counting Electrons4 H 2 Every H has 2 e. OK CH 4 H has 2 e, C 8. OK NH 3 N has 8 e. Nucleophile! OK Use a localized (valence-bond) model to count electrons

5. 6/26/2015Counting Electrons5 C 2 H 4 C has 8 e. OK singlet CH 2 C has only 6 e, and an empty p z orbital: extremely reactive ("singlet carbene"). Unstable. Sensitive to nucleophiles and electrophiles. triplet CH 2 C has only 6 e, is a "biradical" and extremely reactive ("triplet carbene"), but not especially for nucleophiles or electrophiles.

6. 6/26/2015Counting Electrons6 CH 3 + C has only 6 e, and an empty p z orbital: extremely reactive. Unstable. Sensitive to nucleophiles. CH 3 - C has 8 e, but a lone pair. Sensitive to electrophiles. Cl - Cl has 8 e, 4 lone pairs. OK Somewhat sensitive to electrophiles.

7. 6/26/2015Counting Electrons7 BH 3 B has only 6 e, not stable as monomer, forms B 2 H 6 : B 2 H 6 B has 8 e, all H's 2 (including the bridging H!). 2-electron-3-center bonds! OK AlCl 3 Al has only 6 e, not stable as monomer, forms Al 2 Cl 6 : Al 2 Cl 6 Al has 8 e, all Cl's too (including the bridging Cl!). Regular 2-electron-2-center bonds! OK

8. 6/26/2015Counting Electrons8 2 MeAlCl 2  Me 2 Al 2 Cl 4 2-electron-3-center bonds are a stopgap! H 3 B·NH 3 N-B: donor-acceptor bond (nucleophile NH 3 has attacked electrophile BH 3 ). Organometallic chemists are "sloppy" and write. Writing or would be more correct (although the latter does not reflect the “real” charge distribution).

9. 6/26/2015Counting Electrons9 PCl 5 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 PCl 3 en Cl 2. HF 2 - Write as FH·F -, mainly ion-dipole interaction.

10. 6/26/2015Counting Electrons10 1.Number of valence electrons (from periodic table) 2.Correct for charge, if any (only if it belongs to that atom!) 3.Count 1 e for every covalent bond to another atom 4.Count 2 e for every dative bond from another atom 5.Add How do you count ?

11. 6/26/2015Counting Electrons11 B=3 -=1 4  H=4 tot=8 OK Examples: counting electrons C=4 1  =O=2 2  H=2 tot=8 OK Pd=10 -=1 3  Cl=3 1  NH 3 =2 tot=16 could have additional 2 e (Pd-Cl  -bond?) Ru=8 2  Cl=2 2  PMe 3 =4 1  CH 2 =2 tot=16 could have additional 2 e

12. 6/26/2015Counting Electrons12 Counting is not always trivial Pd=10 2  -=2 3  Cl=3 1  CH 2 =1 tot=16 could have additional 2 e

13. 6/26/2015Counting Electrons13 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. Remember, when counting:

14. 6/26/2015Counting Electrons14 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. Oxidation States

15. 6/26/2015Counting Electrons15 1.Start with the formal charge on the metal 2.Ignore dative bonds 3.Ignore bonds between atoms of the same element (this one is a bit silly) 4.Assign every covalent electron pair to the most electronegative element in the bond: this produces + and – charges (usually + at the metal) 5.Add Calculating the formal oxidation state

16. 6/26/2015Counting Electrons16 charge C=0 4  C-Cl: C + -Cl - =+4 tot=+4 Examples: oxidation states charge C=0 2  C-Cl: C + -Cl - =+2 1  C=O: C 2+ -O 2- =+2 tot=+4 charge Al=-1 4  Al-Cl: Al + -Cl - =+4 tot=+3 charge Mn=-1 4  Mn=O: Mn 2+ -O 2- =+8 tot=+7 charge Pd=-2 4  Pd-Cl: Pd + -Cl - =+4 tot=+2

17. 6/26/2015Counting Electrons17 charge C=0 3  C-Cl: C + -Cl - =+3 tot=+3 trivalent carbon ? Examples: oxidation states charge Mg=0 4  Mg-Me: Mg + -Me - =+4 tot=+4 impossible, Mg has only 2 valence electrons! charge Pt=-2 3  Pt-Cl: Pt + -Cl - =+3 tot=+1 univalent Pt ?

18. 6/26/2015Counting Electrons18 Oxidation states are formal. However, they do give an indication whether a structure or composition is reasonable (apart from the M-M complication). The significance of an oxidation state ?

19. 6/26/2015Counting Electrons19 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  0 for metals –usually +n for very electropositive metals –usually 0-3 for 1 st -row transition metals of groups 6-11, often higher for 2 nd and 3 rd row –electronegative ligands (F,O) stabilize higher oxidation states,  - acceptor ligands (CO) stabilize lower oxidation states –oxidation states usually change from m to m-2, m or m+2 in reactions Acceptable oxidation states

20. 6/26/2015Counting Electrons20 Simply the number of atoms directly bonded to the atom you are interested in, regardless of bond orders etc. CH 4 :4 C 2 H 4 :3 C 2 H 2 :2 AlCl 4 - :4 Me 4 Zn 2- :4 OsO 4 :4 Coordination number B 2 H 6 :4 (B) 1 (terminal H) 2 (bridging H)

21. 6/26/2015Counting Electrons21 For complexes with  -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: Coordination Number C.N. 4

22. 6/26/2015Counting Electrons22 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. Coordination Number

23. 6/26/2015Counting Electrons23 C.N."Normal" geometry 2linear or bent 3planar trigonal, pyramidal, "T-shaped" 4square planar, tetrahedral 5square pyramid, trigonal bipyramid 6octahedron Coordination number and coordination geometry

24. 6/26/2015Counting Electrons24 Could WH 6 (PMe 3 ) 3 be ? Count W: 18 VE (OK), oxidation state 6 (OK), coordination number 9 (very high). Possible. Protonation gives WH 7 (PMe 3 ) 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! Illustration: protonation of WH 6 (PMe 3 ) 3

25. 6/26/2015Counting Electrons25 Give electron count and oxidation state for the following compounds. Draw conclusions about their (in)stability. Me 2 MgPd(PMe 3 ) 4 MeReO 3 ZnCl 4 Pd(PMe 3 ) 3 OsO 3 (NPh) ZrCl 4 ZnMe 4 2- OsO 4 (pyridine) Co(CO) 4 - Mn(CO) 5 - Cr(CO) 6 V(CO) 6 - V(CO) 6 Zr(CO) 6 4+ PdCl(PMe 3 ) 3 RhCl 2 (PMe 3 ) 2 Ni(PMe 3 )Cl 4 Ni(PMe 3 )Cl 3 Ni(PMe 3 ) 2 Cl 2 Exercises