1. Chemical Bonding
Chapters 9 & 10

2. Review Draw Lewis dot structures for
SO2 BrF41+ ClO2
Use formal charge to determine whether N or O is the more likely central atom in NOF

3. How equally do atoms share e– ?
Hydrogen and fluorine share one pair of e– in a single covalent bond
In the Lewis dot structure, they appear to share the electrons equally, but do they?

4. How equally do atoms share e– ?
When you put HF in an electric field, the molecules line up, as if the F end were negative and the H end positive.
The electrons are shared unequally.

5. How equally do atoms share e– ?
When bond e– are shared unequally, the bond is said to be a polar covalent bond.
A polar covalent bond has a dipole moment: one end is more negative than the other end

6. Electronegativity Electronegativity is the ability to attract bond e–
The higher the EN, the “greedier” the atom

7. Electronegativity
Electronegativity increases across a period and decreases down a group

8. Electronegativity and bond polarity
Bond polarity depends on the difference in EN

9. Evaluating bond type
F–F

10. Evaluating bond type
F–F
F = 4.0, F = 4.0
∆EN = 0

11. Evaluating bond type F–F F–F F = 4.0, F = 4.0 ∆EN = 0
Bond is nonpolar covalent (zero dipole moment)
F–F

12. Evaluating bond type
C = O

13. Evaluating bond type
C = O
C = 2.5, O = 3.5
∆EN = 1.0

14. Evaluating bond type C = O C=O – C = 2.5, O = 3.5 ∆EN = 1.0
Bond is polar covalent (has a dipole moment) with O end more negative
Doesn’t matter whether bond is single or double
C=O –

15. Evaluating bond type
KCl

16. Evaluating bond type
KCl
K = 0.8, Cl = 3.0
∆EN = 2.2

17. K1+ Cl1– Evaluating bond type KCl K = 0.8, Cl = 3.0 ∆EN = 2.2
Bond is ionic; e– transferred from K to Cl
K Cl1–

18. Order, length, energy Bond order = type of bond
1 = single, 2 = double, 3 = triple
Bond length = distance between bonded atoms
Higher bond order => shorter bond length
Bond energy = energy needed to break bond
Higher bond order => higher bond energy

19. Estimating ∆H from bond energy
Breaking bonds is endothermic (energy in)
Making bonds is exothermic (energy out)
∆H is overall energy change
∆H negative (exothermic) when weak bonds break & strong bonds form
∆H positive (endothermic) when strong bonds break & weak bonds form

20. CH4 + 2 O2 → CO2 + 2 H2O Observed ∆H for this reaction is –890 kJ
(values from bond energy are approximate)

21. Shapes of Molecules
Lewis dot structures do not reveal the three-dimensional shape of a molecule
Molecular shape can be predicted from the dot structure using the Valence Shell Electron Pair Repulsion (VSEPR) model

22. VSEPR The VSEPR model focuses on electron groups in the valence shell
A bond (single, double, or triple) is one electron group
A lone pair is one electron group
VSEPR proposes that electron groups will take positions around the central atom that are as far away from each other as possible, to minimize repulsions
This gives a set of 5 possible geometries

23. Terminology A = central atom
X = atom or group of atoms bonded to central atom
E = lone pair of e– on central atom
AX2
AX2E2
AX4

24. Terminology
Electron group geometry = shape made by the e– groups (bonds + lone pairs)
2 electron groups = linear
3 electron groups = trigonal planar
4 electron groups = tetrahedral
5 electron groups = trigonal bipyramidal
6 electron groups = octahedral

25. Sketching the 5 basic geometries
Linear
Trigonal planar
Tetrahedral
Trigonal bipyramidal
Octahedral

26. Terminology
Molecular geometry = shape made by joining bonded nuclei with straight lines
Each e– group geometry forms one or more molecular geometries
Example: four electron groups (tetrahedral) could be
AX4 tetrahedral
AX3E trigonal pyramidal
AX2E2 angular or bent
Bond angle = angle between adjacent bonds

27. Applying VSEPR Theory Draw a plausible Lewis structure.
Determine the number of bonds and lone pairs, and assign a VSEPR notation (AXE) to the molecule.
Establish the e– group geometry.
Determine the molecular geometry.
If there is more than one central atom, analyze each atom individually.
Sketch the electron group geometry, indicating atoms with circles and lone pairs with dots.

28. Molecular Geometry as a Function of Electron Group Geometry

29. Molecular Geometry as a Function of Electron Group Geometry

30. Beyond Lewis theory Lewis theory but
Describes what happens (dot structures)
Offers a simple idea of why (octet rule)
but
Is not correlated to modern atomic theory (orbitals)

31. Valence bond theory
Chemical bond forms when half-filled orbitals overlap at optimum balance of attraction & repulsion

32. Valence bond theory
For most molecules, molecular geometry does not match orientation of atomic orbitals

33. Hybridization of atomic orbitals
When bonds form, central atom valence orbitals combine → new set of wave functions called the hybrid set
Number of hybrid orbitals = number of atomic orbitals combined
Hybrid orbitals have different shape and orientation than original orbitals
Shapes of hybrid sets correspond to VSEPR geometries

34. sp3 hybrid set One 2s + three 2p orbitals → four sp3 orbitals
The orbitals in the sp3 set all have the same energy
Four valence e– enter the four sp3orbitals according to Hund’s Rule.

36. sp3 hybrid set C has 4 valence e– N has 5 valence e–
Each valence e– is in an sp3 orbital
Four half-filled sp3 overlap with four half-filled 1s H orbitals
N has 5 valence e–
Three half-filled sp3 overlap with three half-filled 1s H orbitals
Fourth sp3 contains lone pair

37. sp2 hybrid set
2s + 2px + 2py orbitals → three sp2 orbitals in xy plane
2pz orbital is unhybridized, remains along z axis

39. Single & double bonds in VB
The hybrid orbitals overlap head-to-head with orbitals on other atoms to form single bonds
Head-to-head overlap is called a sigma () overlap
Sigma overlaps provide the skeleton structure and VSEPR shape

40. Single & double bonds in VB
Unhybridized orbital overlaps side-to-side with orbital on another atom to form another bond
Side-to-side overlap is called a pi () overlap
Pi overlap is the second bond in a double bond

41. Bonding in H2CO

43. sp hybrid set 2s + 2px orbitals → two sp orbitals on x axis
2py + 2pz orbitals are unhybridized, remain along y and z axes

45. sp hybrid set
The sp set can form two sigma bonds (s) and two pi bonds (p)
Single bond = triple bond
Two double bonds

46. Bonding in C2H2

47. sp3d hybrid set
One 3s + three 3p + one 3d orbital → five sp3d orbitals
Valence e– must be on n ≥ 3 to form this set
No double or triple bonding

48. sp3d2 hybrid set
One 3s + three 3p + two 3d orbitals → six sp3d2 orbitals
Valence e– must be on n ≥ 3 to form this set
No double or triple bonding

49. One  + two  bonds
Two  + one  bond
All  bonds