On Tap for Today— BONDS

1. On Tap for Today—BONDS • Not the pieces of paper issued by the Gvmnt • But because atoms graduate to molecules…

2. Chapter 7 “Pre-Review” • This chapter has the greatest long term impact on success in chemistry (particularly Organic). • Organic topics covered, bonding, nomenclature structures… • First we cover all types of bonds, focus on covalent and polar covalent • It really pays to thoroughly understand this

3. End of Public Service Anoncmnt • Overview of Chemical Bonds (types, Lewis, octets)—extension of what’s known • Ionic bonding—strongest of all chemical bonds—based almost entirely on EC’s • Covalent bonds (this is the organic) • Polar covalent/electronegativity/etc • Lewis structures (biggest lesson to learn)

4. More on Lewis Structures • Strategy for writing them (HINT, as much ART as it is SCIENCE) • Skeletons, resonance, central atoms, # of bonds, formal charge, all in due time • Exceptions (there are always exceptions) • Bond length/strength data

5. Lewis structures continued • Also covered term “octet” rule—which states atoms tend to gain or lose electrons in order to adopt a noble gas configuration • Remember, noble gases have ns2np6 • Start with Lewis symbols (dots)

6. Covalent Bonding—Lewis Structutres • Covalent bonds—atoms share electrons • Number of dots = # of valence electrons • Number of unpaired electrons = # bonds • C = 4 • N / B = 3 • O = 2 • F = 1 Already paired Unshared electron

7. Determine the number of valence electrons • Draw a skeletal structure (takes practice…use table below) • Distribute excess electrons—make octet • Excess electrons go to central atom, deficiency—multiple • Calculate formal charge of each atom (table below helps)

8. Number of valence electrons • The reason I’ve stressed Group/Valence electron trends starts to make sense now. We need to know… • HOW MANY ELECTRONS does a molecule have? • Each atom brings valence electrons to the table, so to speak • Anions ADD electrons (more negative charge) • Cations subtract electrons (less e’s) • I KNOW THIS SOUNDS DUMB…BUT ADD THEM UP • ALWAYS!! It’s important!! If you miss a lone pair, you mess up geometry (later)…so just GET USED TO THIS STEP!!

9. That was step 1 • Now that you know the # of bonds an atom typically forms—step two is drawing the bare bones structure (skeletal) • This is more art than science. BUT, here are some tips • H is always a terminal atom (forms only 1 bond) NEVER > 1!!! • Central atoms normally lowest electronegativity • Terminal atoms normally higher electronegativity, F, ALWAYS • Symmetrical/compact structures better

10. Electronegativity Trends I HHAATTEE memorization, but you should remember the ‘big’ four elements: N, O, F, Cl.

11. Formal Charge—Quick and Dirty • Books give a formula for calculating formal charge, but those are pretty clunky…something simpler • Count all bonds (as 1) and lone pair electrons (1 for each electron) • Btw…thiocyanate? Replace the O with an S O C N # valence e’s (for element) 6 4 5 Number e’s (see counting) 6 4 6 Subtract 2 from 1 0 0 -1 Formal Charge resides on Nitrogen Here, formal charge is on the oxygen. Are these resonance forms?

12. Adjacent atoms in a structure should not carry formal charges of the same sign Formal Charge Usually, the most plausible Lewis structure is one with no formal charges When formal charges are required, they should be as small as possible and negative formal charges should appear on the most electronegative atoms

13. EOS Resonance: Delocalized Bonding Resonance theory states that whenever a molecule or ion can be represented by two or more plausible Lewis structures that differ only in the distribution of electrons, the true structure is a composite, or hybrid, of them Resonance structures

14. Molecules that Don’t Followthe Octet Rule Molecules with an odd number of valence electrons have at least one of them unpaired and are called free radicals

15. Molecules that Don’t Followthe Octet Rule Some molecules have incomplete octets. These are usually compounds of Be, B, and Al, generally have some unusual bonding characteristics, and are often quite reactive EOS

16. Molecules that Don’t Followthe Octet Rule Some compounds have expanded valence shells, which means that the central atom has more than eight electrons around it EOS

17. Molecules that Don’t Followthe Octet Rule An expanded valence shell may also need to accommodate lone-pair electrons as well as bonding pairs EOS

18. Valence Shell Electron Pair Reupulsion • Fancy way of saying electrons repel each other. • VSEPR predicts shape based on maximizing distance • Number of Electron Groups dictates the shape a molecule adopts • 2 Electron Groups—linear • 3 Electron Groups—Trigonal planar • 4 EG’s (See Amy or Ernie) • 5 EG’s trigonal bipyramidal • 6 OH (six bonds? OCTA hedral? What gives?

19. Helps to define electron group EG’s are ANY group of electrons, bonding, lone etc. Single bond Double bond Triple bond Lone pair of electrons Radical (one e) Remember, it maximizes the repulsion between ALL electrons EGG’s ARE important though… HOWEVER, we more often care about the shape of the MOLECULE; bonding pairs only (lone/radicals be damned) Molecular Group Geometry describes the orientation of JUST the bonding pairs More accurately stated…the bonding GROUPS Maddeningly enough, some molecules have LONE pairs, so what to do? Electron Groups “defined”

20. Shapes of simple structures • Two and Three EG’s: • Little variation in the shape • Two is always linear • Three has two options • Bent, or trigonal planar

21. 4-Coordinate…Organic • Not really going to drag this out much, you ALL should know this pretty well • Methane, ammonia and water, the unholy trinity? • 4:0 tetrahedral • 3:1 trigonal pyramidal! • 2:2 bent (I hate ‘angular’) • Note bond angles!!

22. For the uber lazy…the summary panel

23. Promotion—followed by hybridztn • Remember that promotion is energy intensive, takes energy…BUT…not as much energy saved when you can form additional bonds • Making bonds RELEASES energy, breaking takes E

24. Consistent with what we’ve learned? • Yes…remember that orbitals are mathematical ‘equations’—probability of finding an electron • When you combine several orbitals, the eq changes

25. This helps explain methane • Also important to remember that the number of atomic orbitals INTO a hybrid scheme MUST equal the number of hybrid orbitals OUT of that scheme • Previous example means we’re dealing with an sp3 orbital (one s and three p orbitals).

26. Not all hybrids are bonding orbitals • In each case…methane, ammonia, and water are all sp3 hybridized, but lone pairs of electrons occupy some of these orbitals

27. What about other schemes • Yep, can do those too. What about an sp2 scheme (one s and two p orbitals)? • 3 in, 3 out, all at 120° (just like Electron group geometry) • One orbital isn’t

28. sp Hybridization in Be … with two unused p orbitals. Two AOs combine to form … … two hybrid AOs …

29. Let’s relate Electron Cloud/hybrid geo! • Note that the EGG and the hybrid orbital geometry are the same (which is why it’s often necessary to assign EGG in the first place).

30. Hybrid Orbitals andMultiple Covalent Bonds • Covalent bonds formed by the end-to-end overlap of orbitals are called sigma (s) bonds. • All single bonds are sigma bonds. • A bond formed by parallel, or side-by-side, orbital overlap is called a pi (p) bond. • A double bond is made up of one sigma bond and one pi bond. • A triple bond is made up of one sigma bond and two pi bonds.

31. VB Theory for Ethylene, C2H4 The hybridization and bonding scheme is described by listing each bond and its overlap. Sigma bonds, and pi bonds

32. Formic acid, HCOOH, is the simplest carboxylic acid. • Predict a plausible molecular geometry for this molecule. • Propose a hybridization scheme for the central atoms that is consistent with that geometry. • Sketch a bonding scheme for the molecule.