1. SCH4U – UNIT 1 STRUCTURE AND PROPERTIES
CHAPTER 4 – CHEMICAL BONDING

2. Activity
With a partner discuss everything you remember about chemical bonding
Eg. Types of bonds? why? What happens?

3. 4.1 Types of Chemical Bonds
What are the two main types of chemical bonds?
Ionic: chemical bond between oppositely charged ions
Electrostatic attraction
Covalent: a chemical bond in which atoms share bonding electrons
Bonding Electron Pair: electron pair that is involved in bonding

4. Bond type depends on the attraction for electrons of the atoms involved
i.e. electronegativity

5. Ionic Compounds

6. How do these work?
Metal + Non-Metal  Metal+ + Non-Metal- Low IE High IE Isoelectronic with noble gases Low EA High EA Opposites attract in no particular direction, considered non-directional Ions cling together in clusters known as crystals

7. Get a lattice structure
Lattice energy: energy change when one mole of an ionic substance is formed from its gaseous ions
Depends on:
Charge on the ions
Size of the ions

8. Ionic Compounds and Bonding
Properties – WHY?
Do not conduct electric current in the solid state
Conduct electric current in the liquid state
When soluble in water, form good electrolyte
Relatively high MP and B
Brittle, easily broken under stress

9. Covalent Bonds Balance of attractive and repulsive forces
What are the forces acting here?

12. Octet Rule
Atoms share electrons so that they are surrounded by 8 electorns
# bonds = 8 - # valence electorn
Example: Carbon, Oxygen, Nitrogen
Two covalent bonds = double bond
Three covalent bonds = triple bond

13. Lewis Structures
Atoms and ions are stable if they have a full valence shell
Electrons are most stable when they are pair
Atoms form chemical bonds to achieve full valence shells of electrons
Full valence shell may occur by an exchange or by sharing electrons
Sharing – covalent; exchange - ionic

16. Polar Covalent Bonds
When electrons are shared unevenly in a covalent bond
Example: HF, H2O

17. Coordinate Covalent Bonds
Both electrons are contributed by one atom
Example:
NH4+
H3O+ CO
N2O
NHO3

18. Resonance Structures
Single bonds are longer than double bonds, which are longer than triple bonds
Example: SO3
Resonance Structure: Electron pair is shared over three bond evenly
Delocalized electrons

19. Less than 8
BeH2
BCl3

20. More than 8 Octet rule only applies to the first two periods
After that, can have expanded octets
Example:
PF5
BrF5
SiF63-

21. Practice - Worksheet H2 F2
OF2
O2F2

22. Valence Bond Theory and Quantum Mechanics
Covalent bonds occur when orbitals overlap and two electrons occupy the same region of space
Example: H2

23. HF What are the electron configurations for H and F?
How would the orbitals interact

24. H2O What are the electron configurations for H and O?
How would the orbitals interact

25. Problem
We know from experiments in atomic structure that the bond angle in H2O is 104.5°… not 90° as predicted by valence bond theory
True for CH4 (109.5°) and NH3 (107.5°) – VBT always predicts bond angles of 90°
So, we need a better theory…

26. Hybrization Two problems still exist from Lewis Bonding Theory
Carbon atoms form 4 EQUAL C-H bonds in CH4 (or any other molecule)
Not predicted due to electron configuration of C
Recall: s orbitals have lower energy than p orbitals, therefore the bond length would be different
Existence of double and triple bonds

27. Hybridization of Carbon Orbitals
An s electron gets promoted to the empty p-orbital
This stabilizes the p- and s- orbitals and gives them all the same energy;
Half-filled subshells
Called sp3 orbitals (HYBRID ORBITALS)
Each sp3 orbital lies at 109.5°

29. Additional Hybrid Orbitals – sp - LINEAR

30. Additional Hybrid Orbitals – sp2 - PLANAR

31. Additional Hybrid Orbitals – sp3 - Tetrahedral

32. Double and Triple Bonds
Two types of orbital overlap exist
What we have seen so far is one type
Sigma bonds: σ-bonds
End-on-end overlap of orbitals
Pi bonds:π-bonds
Sideways overlap of orbitals

33. Sigma Bonds
Occur in single bonds and account for the FIRST bond in a double or triple bond
Examples:

34. Pi Bonds
Occur when p-orbitals not on the bonding axis (py or pz) overlap with each other

35. Making Double Bonds Example: C2H4 Draw a Lewis Structure
What occurs with the C atoms hybridization?

36. For double bonds, there must be one σ-bond from overlapping hybrid orbitals and one π-bond from overlapping py or pz orbitals
Come from sp2 hybridized orbitals and result in trigonal planar structures

37. Making Triple Bonds
Triple bonds have one σ-bond and two π-bonds; come from sp-hybridized orbitals, and result in linear structures
Central atom has two un-hybridized p-orbitals

39. Practice
Explain the structure of the following molecules using electron configurations, orbital hybridization and VBT.
C2Cl4
C2Cl2
CO2

40. VSEPR Theory - Valence Shell Electron Pair Repulsion Theory

41. Work through VSEPR Chart
Fun times with molecular structure…

42. Practice Problems
Use Lewis Theory and VSEPR Theory to predict the structure of the following molecules:
Homework/Practice - Worksheet

43. Polar Molecules
Polar molecules are molecules where the electron charge is not distributed evenly

44. Electronegativity and Polar Covalent Bonds
Ionic Bond: ΔEN = >1.7
Electron transfer
Polar Covalent Bond: ΔEN =
Electrons shared unevenly
Pure Covalent Bond: ΔEN =
Electrons shared evenly
Remember: Think of electrons as electron probabilities, electron cloud density is greater around one atom or another, therefore one gets a slight negative, the other slight positive charge

45. Think of the scale as a continuum

46. Polar Molecules Cannot exist if there are no polar bonds!
Bond dipole: electronegativity difference of two atoms represented by an arrow pointing from the positive to the negative end (lower to higher EN)
Non-polar molecule: either perfectly symmetrical so the bond dipoles cancel out, or when no polar bonds exist
Polar molecule: occur when bond dipoles do not cancelout

47. Example:
Determine the polarity of the following molecules
H2O, CCl4, NH3, PCl5
Practice:
CH3Cl, BeCl2, SiO2, BrF4
CHF2Cl
CH3NH2

48. Intermolecular Forces
Forces that exist between molecules
Three types:
Dipole-Dipole
Hydrogen Bonds
London Dispersion
In order to determine the Intermolecular Forces (IMF), you need to first determine the polarity of the molecule
Much weaker than covalent bonds

49. Dipole-Dipole Forces Occur in polar molecules
The slightly negative end on one molecule is attracted to the slightly positive end on another molecule
Strength depends on the size of the dipole

50. London Dispersion Forces
Simultaneous attraction of the electrons in one molecule to the nuclei in the surrounding molecules
Increase as the number of electrons and protons in a molecule increase
Exist in ALL molecules
Weakest Force

51. Hydrogen Bonds
Attraction between H on one molecule and O, N, or F on another molecule
Strongest of the intermolecular forces
Found in H2O, NH3, and HF, or whenever there is a –OH, -NH2 in a molecule

52. Predicting Strength of IMF
Use pol

53. Predicting Boiling Points
Boiling points increase as IMF strength increases
Arrange the following molecules in order of increasing boiling points
SiH4, SnH4, GeH4, CH4,
C3H8, C2H4, C4H10
CH4, CCl3H, CBr3H

54. Practice
Determine the intermolecular forces that exist in each molecule
CCl4, C5H12, CH3CH2OH
Which molecules would have the strongest IMF
C2H5OH, C2H6, C2H5Cl
Explain you answer

55. Structure and Properties of Solids
Different types of solids result depending on the type of bonding in the solid
These solids have different properties

56. Ionic Crystals Crystal Lattice
Properties result from the lattice structure
Brittle, high melting/boiling point, conduct electricity when dissolved in water or in liquid form, hard

57. Molecular Crystals
Arrangement of neutral molecules held together by weak intermolecular forces
Properties vary depending on the strength of the IMF

58. Covalent Network Solids
Array of covalently bonded atoms, structure is held together by covalent bonds
High MP/BP
Example: Silicate (SiO)

59. Carbon Network Solids
Diamond, Graphite, Carbon Nanotubes, Buckminster Fullerenes
Explain the difference in properties between graphite and diamond.

60. Metallic Crystals
Lots of electrons, but low ionization energy means they are loosely held
Lots of empty valence orbitals with similar energy, therefore electrons are free to move around
Strong, non-directional bonding

61. Properties of Metals Property Explanation Shiny, Silvery Flexible
Electrical Conductivity
Hard Solids
Crystalline

62. End of UNIT 1 – STRUCTURE AND PROPERTIES!!
Your unit test is on: October 28
Review package handed out