1. Bonding Theory Overview
Lewis structure
Formal Charge
Bond polarity; Electrostatic Potential Map
Valence bond theory vs. Molecular orbital theory
VSEPR theory
Polarity and Solubility

2. Atomic Structure A review from General Chemistry
Protons (+1) and neutrons (neutral) reside in the nucleus
Electrons (-1) reside outside the nucleus.
High energy electrons are far away from the nucleus and others are far away
Electrons at the highest shell (far away from nucleus) are the valence electrons
Valence electrons are important as they participate bonding

3. Electrons in Atom: Atomic Orbitals
Quantum physics description of electron in Atom: Wavefunctions ψ
Wavefunction may have sign of (–), (+), or ZERO, depending on the location.
Ψ2 describes the probability of electrons (electron density) around the nucleus.
Nodal plane: Ψ2 = 0. Zero electron density.
The sign of the wavefunction will be important for orbital overlapping in bonds.

4. Shapes of Atomic Orbitals
s Orbital
p Orbitals: px , py , pz
d Orbitals

5. Covalent Bonding
A covalent bond is a PAIR of electrons shared between two atoms. For example…

6. Polar Covalent Bonds Covalent bonds are either polar or nonpolar
Nonpolar Covalent –bonded atoms share electrons evenly
Polar Covalent – One of the atoms attracts electrons more than the other
Electronegativity - how strongly an atom attracts shared electrons

7. Polar Covalent Bonds
Electrons tend to shift away from lower electronegativity atoms to higher electronegativity atoms.
The greater the difference in electronegativity, the more polar the bond.

8. Isomerism: different substances with the same formula

9. Common atoms in Organic molecules
Organic compounds are based on hydrocarbon compounds (CxHy).
Atoms that are most commonly bonded to carbon include N, O, H, S and halides (F, Cl, Br, I).
With some exceptions, each element generally forms a specific number of bonds with other atoms (HONC)

10. Lewis Structures For simple Lewis Structures…
Draw the individual atoms using dots to represent the valence electrons.
Put the atoms together so they share PAIRS of electrons to make complete octets.
HONC rule (#bonds: 1234): preference of atom as center

11. Formal Charge
FORMAL charge: Atoms with an unbalanced electrons and protons
Molecules with more formal charges tend to have higher energy, less stable, more reactive
Electrons in the molecule: Bonding electrons (BE) vs. Lone Pair electrons (LPE)
FORMAL charge = #VE – (½ #BE + #LPE)

12. Example: To Find Formal Charge
Calculate the formal charge on each atom. or
Carbon (4 single bonds): FC = 4 – ½(4x2) = 0
Hydrogen (1 single bond): FC = 1 – ½(1x2) = 0
Oxygen (1 single bond): FC = 6 – (½ x 1 x 2 + 6) = -1

13. Practice: Lewis Structures and Formal Charges
Draw Lewis structure for N2O:
two double bonds
Find formal charges for all atoms in this structure

14. Practice: Lewis Structures and Formal Charges
Draw Lewis structure for N2O:
B. Single + triple bonds
Find formal charges for all atoms in this structure

15. Valence Bond Theory
Chemical bonds form when the orbitals (wave functions of electrons in the atoms) on those atoms interact (overlap)
The kind of overlap depends on whether the orbitals align along the axis between the nuclei (sigma bond), or outside the axis (pi bond)
Tro: Chemistry: A Molecular Approach, 2/e

16. H2 according to VB Theory
The bond for a H2 molecule results from constructive interference
Where do the bonded electrons spend most of their time?

17. VB theory: Hybridized Atomic Orbitals
The bond angle in methane can not be explained by simple bonding between the 2s or 2p orbital of carbon with 1s orbital of hydrogen.

18. Hybridized Atomic Orbitals
The carbon must undergo hybridization to form 4 equal atomic orbitals
The atomic orbitals must be equal in energy to form four equal-energy symmetrical C-H bonds

19. Hybridized Atomic Orbitals
Should the shape of an sp3 orbital look more like an s or more like p orbital?

20. Hybridized Atomic Orbitals
To make CH4, the 1s atomic orbitals of four H atoms will overlap with the four sp3 hybrid atomic orbitals of C

21. Tro: Chemistry: A Molecular Approach, 2/e

22. Practice: Hybridized Atomic Orbitals
Draw a picture that shows the necessary atomic orbitals and their overlap to form ethane (C2H6).
Draw a picture that shows the necessary atomic orbitals and their overlap to form water.

23. sp2 hybridization Consider ethene (ethylene).
Each carbon in ethene must bond to three other atoms, so only three hybridized atomic orbitals are needed

24. sp2 Hybridized Atomic Orbitals
An sp2 hybridized carbon will have three equal-energy sp2 orbitals and one unhybridized p orbital

25. σ Bond The sp2 atomic orbitals overlap to form sigma (σ) bonds
Sigma bonds provide maximum HEAD-ON overlap
Free rotation along C-C single bond

26. π bond
The unhybridized p orbitals in ethene form pi (π) bonds, SIDE-BY-SIDE overlap
Rotation along C=C bond requires high energy

27. sp Hybridized Atomic Orbitals
Consider ethyne (acetylene).
Each carbon in ethyne must bond to two other atoms, so only two hybridized atomic orbitals are needed

28. sp Hybridized Atomic Orbitals
The sp atomic orbitals overlap HEAD-ON to form sigma (σ) bonds while the unhybridized p orbitals overlap SIDE-BY-SIDE to form pi bonds
Practice with SkillBuilder 1.7

29. Comparison: C-C, C=C, and CC bonds
Note the different strengths and lengths below.

30. Combination of Wavefunctions
A bond occurs when constructive atomic orbitals overlap. Overlapping orbitals is like wave interference.
Only constructive interference results in a bond

31. Molecular Orbital (MO) Theory
Atomic orbital wavefunctions overlap to form MOs that extend over the entire molecule.
MOs can be both constructive (Bonding MO, lower energy) and destructive (Antibonding MO, higher energy) interference.
The number of MOs created must be equal to the number of AOs that were used.
H2 MOs

32. MO Theory: Examples Consider TWO of the many MOs that exist for CH3Br
There are many areas of atomic orbital overlap
Notice how the MOs extend over the entire molecule
Each picture below represents ONE orbital.
bonding MO antibonding MO

33. About MO bonds
For each MO, up to TWO electrons can fit into the orbital.
In the ground state, electrons occupy some MOs and not others, depending on the relative energy level
Depending on the circumstances, we will use both MO and valence bond theory to explain phenomena

34. π bond in MO theory
The unhybridized p orbitals in ethene form pi (π) bonds, SIDE-BY-SIDE overlap of p-orbitals giving both bonding and anti-bonding MO orbitals
MO theory shows the orbitals that result.
Remember, red and blue regions are all part of the same orbital

35. Molecular Geometry
Valence shell electron pair repulsion (VSEPR theory)
Valence electrons (bonded and lone pairs) repel each other
To determine molecular geometry…
Determine the Steric number

36. Molecular Geometry
Valence shell electron pair repulsion (VSEPR theory)
Valence electrons (bonded and lone pairs) repel each other
To determine molecular geometry…
Predict the hybridization of the central atom
If the Steric number is 4, then it is sp3
If the Steric number is 3, then it is sp2
If the Steric number is 2, then it is sp

37. sp3 Geometry
For any sp3 hybridized atom, the 4 valence electron pairs will form a tetrahedral electron group geometry
Methane has 4 equal bonds, so the bond angles are equal
lone pair electrons repel stronger, smaller bond angles
More lone pair electrons causes even smaller bond angle

38. sp3 Geometry
The molecular geometry is different from the electron group geometry. HOW?

39. sp2 Geometry Calculate the Steric number for BF3
Electron pairs that are located in sp2 hybridized orbitals will form a trigonal planar electron group geometry
molecular geometry: _____________

40. sp2 Geometry
Analyze the steric number, hybridization, electron group geometry and molecular geometry for this imine?

41. sp Geometry
Analyze the Steric number, the hybridization, the electron group geometry, and the molecular geometry for the following molecules
BeH2
CO2

42. From Hybridization to Geometry

43. Molecular Polarity Electronegativity Differences cause induction
Induction (shifting of electrons WITHIN their orbitals) results in a dipole moment.
Dipole moment = (the amount of partial charge) x (the distance the δ+ and δ- are separated)

44. Molecular Polarity Polarity of some other common bonds
Percent ionic character of covalent bond depends on difference in electronegativity.

45. Molecular Polarity
For molecules with multiple polar bonds, the dipole moment is the vector sum of all of the individual bond dipoles

46. Molecular Polarity
Electrostatic potential maps are often used to give a visual depiction of polarity

47. The dipole moment for pentane = 0 D?

48. Polarity of Pentane : single bond dipole moment
: locally combined dipole moment

49. Intermolecular Forces (IMF)
Many properties such as solubility, boiling point, density, state of matter, melting point, etc. are affected by the attractions BETWEEN molecules
Neutral molecules (polar and nonpolar) are attracted to one another through…
Dipole-dipole interactions
Hydrogen bonding
Dispersion forces (a.k.a. London forces or fleeting dipole-dipole forces)

50. Dipole-Dipole
Dipole-dipole forces result when polar molecules line up their opposite charges.
Note acetone’s permanent dipole results from the difference in electronegativity between C and O
The dipole-dipole attractions BETWEEN acetone molecules affects acetone’s boiling and melting points. HOW?

51. Dipole-Dipole
Why do isobutylene and acetone have such different MP and BPs?

52. Hydrogen Bonding
Hydrogen bonds are an especially strong type of dipole-dipole attraction
Hydrogen bonds are strong because the partial + and – charges are relatively large

53. Hydrogen Bonding
Only when a hydrogen shares electrons with a highly electronegative atom (O, N, F) will it carry a large partial positive charge
The large δ+ on the H atom can attract large δ– charges on other molecules
Compounds with H atoms that are capable of forming H-bonds are called protic

54. Practice: Hydrogen Bonding in Solvent
Which of the following solvents are protic (capable of H-bonding), and which are not (aprotic)?
Acetic acid
Diethyl ether
Methylene chloride (CH2Cl2)
Dimethyl sulfoxide

55. Practice: Hydrogen Bonding and Boiling Point
among the isomers in boiling points
Dispersion force
Dipole dipole force
Hydrogen bonding

56. Hydrogen Bonding
H-bonds are among the forces that cause DNA to form a double helix and some proteins to fold into an alpha-helix

57. London Dispersion Forces
The constant random motion of the electrons in the molecule will sometimes produce an electron distribution that is NOT evenly balanced with the positive charge of the nuclei
Such uneven distribution produces a temporary dipole, which can induce a temporary dipole in a neighboring molecule

58. London Dispersion Forces
The result is a fleeting attraction between the two molecules
Such fleeting attractions are generally weak.
But like any weak attraction, if there are enough of them, they can add up to a lot

59. London Dispersion Forces
The greater the surface area of a molecule, the more temporary dipole attractions are possible
Consider the feet of Gecko. They have many flexible hairs on their feet that maximize surface contact
The resulting London dispersion forces are strong enough to support the weight of the Gecko

60. Molar Mass affects Dispersion Force
Molecules with more mass generally have higher boiling points
Boiling point: I2 > Br2 > Cl2 > F2 CH3OH < C2H5OH

61. Dispersion Forces: Shape of Molecule
Explain why more highly branched molecules generally have lower boiling points

62. Solubility The “like-dissolves-like” rule
Polar compounds generally mix well with other polar compounds
If the compounds mixing are all capable of H-bonding and/or strong dipole-dipole, then there is no reason why they shouldn’t mix
Nonpolar compounds generally mix well with other nonpolar compounds
If none of the compounds are capable of forming strong attractions, then no strong attractions would have to be broken to allow them to mix

63. Solubility
We know it is difficult to get a polar compound (like water) to mix with a nonpolar compound (like oil)
We can’t use just water to wash oil off our dirty cloths
To remove nonpolar oils, grease, and dirt, we need soap

64. How does Detergent help cleaning?
Soap molecules organize into micelles in water, which form a nonpolar interior to carry away dirt.

65. Additional Example Problems
Give all formal charges in the structures below.

66. Additional Example Problems
Analyze the geometry, polarity and types of intermolecular attractions for the following molecules.