1. Covalent Bonds
Chapter 6 p. 188

2. 6.1 Covalent Bonding 6.1 What is a covalent bond? 6.2 Lewis Structures
6.3 Shapes of Molecules
Molecular modeling lab

3. 6.1 Covalent Bonds Think about it………….
What does it means when we use the word bond to describe the relationship between people?

4. 6.1 Standard 2a
Students know atoms exchange electrons to form ionic bonds or,
combine to form molecules by sharing electrons to form covalent bonds……..
‘ELECTROSTATIC’
Teacher notes: cf ionic bond = transfer

5. Learning target How is a covalent bond different from an ionic bond?
What is a covalent bond?

6. 6.1 Covalent Bonds Shared pair of electrons holds the atoms together
Start 12/2
Shared pair of electrons holds the atoms together
‘diatomic’ molecule (H2)

7. 6.1 Covalent Bonds
Think about it…..
What is this diagram telling us?

8. 6.1 Covalent Bonds p. 192 (a)(c)
the bonded arrangement of atoms is lower (lower/higher) energy than the individual atoms)
Energy is released when bonds form

9. 6.1 Covalent Bonds p. 192 c  a Energy is required to break bonds
‘low energy (c)  high energy (a); ‘up-hill’
Energy is required to break bonds

10. Think about it….. Read para 3 p. 193
Study Table 1 - Bond Energies and Bond Lengths for Single Bonds
What is Table 1 telling you about covalent bonds?
Are all covalent bonds equally strong?
Are all covalent bonds equal length?
Trends/patterns?
HONS: how can this information be used to explain why energy is released when methane burns?*
CH4 + O2  CO2 + H2O + energy

12. Frayer: Electronegativity
Definition:
Ameasureoftheabilityofanatominachemicalcompoundtoattractelectrons.
A measure of the ability of an atom in a chemical compound to attract electrons. (p. 137)
Facts:
Read/Notes ‘Electronegativity’ p. 137 (first 3 paragraphs)
Start per 3,5,6

13. 6.1 Electronegativity
Scale: 0-4

14. 5.1 Electronegativity p. 194 Frayer: examples:
Fluorine has the strongest demand for electrons
Francium has the lowest demand for electrons

15. 5.1 Electronegativity p. 195
Facts: Electronegativity difference between two atoms tells us what kind of bond will form between them.

16. 5.1 Electronegativity p. 194
Calculate the EN difference between sodium and chlorine

17. 5.1 Electronegativity p. 195
What kind of bonding occurs between sodium and chlorine?

18. 6.1 Covalent Bonds: Electronegativity difference
NOTES: Ionic bonds
Electronegativity difference is large (small/medium/large?)
atoms gain/lose electrons to form ions
Start per 1

19. 6.1 Word Storm What is the word: POLAR
“The difference between the electronegativity values of hydrogen and fluorine show that H and F atoms form a polar covalent bond.”
per 4,5,6 with word storm review

20. 6.1 Word Storm: POLAR
What does the word polar mean?

21. 6.1 Word Storm: POLAR
Start per 1, 3
End per 5,6

22. 6.1 Covalent Bonds Learning Target
What is the difference between a polar covalent bond and a nonpolar covalent bond?

23. Covalent bonds: Electronegativity difference
Notes: Nonpolar covalent e.g H-H
electrons are shared equally between atoms
Electronegativity difference is small (small/medium/large?)

24. Covalent bonds: Electronegativity difference
Notes: Polar covalent bonds e.g. C-O
electrons are unevenly shared between atoms
Electronegativity difference is medium (small/medium/large?)
P.194

25. 6.1 Practice
Predict the type of chemical bonds (either ionic, polar covalent, non-polar covalent) the following compounds form: KF O2
PBr
P.194
Start per 3
YouTube: Dogs Teaching Chemistry - Chemical Bonds

26. 6.1 Dipole
Notes: Amoleculeorpartofamoleculethatcontainsbothpositivelyandnegativelychargedregions p. 195 Draw it!! PhET sim: molecule polarity

27. 6.1 Bond Dipole: Draw it!! δ = delta δ+ = partially positive
Start per 1,2
δ = delta
δ+ = partially positive

28. 6.1 Molecular Dipoles - NOTES
Individual bond dipoles add together to form an overall molecule dipole.
Vid: covalent bonding (weiner 7000)

29. 6.1 Molecular Dipoles - NOTES
Some molecules are polar and some are non-polar.
Examples:
Carbon dioxide = nonpolar molecule (b/c bond dipoles cancel)
Water = polar molecule (b/c of bent shape)

30. 6.1 Metallic Bonds: Think about it…….
What are the physical properties of metals?
Make a list
Solids (high melting points)
Hardness
Shiny (lustre)
Electrical conductivity
Malleable (can be shaped)
Ductile (can be drawn into a wire)
Start per 3,6 with activity, per 5 with notes

31. Standard
c. Analyze and interpret provided data about bulk properties of various substances to support claims about the relative strength of the interactions among particles in the substance.

32. Metallic Bonding ‘Sea of Electrons’ model

33. 6.1 Metallic Bonds: notes Electrical Conductivity
What is necessary in order to have an electric current?
Electrical charges (electrons or ions) that are free to move (current)
Youtube:metallic bonding and metallic properties

34. 6.2 Drawing and Naming Molecules
Objectives
Draw the Lewis structure of a molecule when given its chemical formula
Friday: skim rdg

35. 6.2 Lewis Structures of atoms
11/16: end per 7
Teacher notes: how electrons are arranged on page

36. 6.2 Lewis Structure
Hydrogen molecule, H2

37. 5.2 Lewis Structures Practice Lets draw the Lewis structures of Cl2
HCl
What is the type of covalent bond in each of these? (polar or nonpolar covalent)
Nonpolar covalent
Polar covalent

38. 5.2 Lewis Structures Notes:
Unshared Pair OR LONE PAIR: anonbondingpairofelectronsinthevalenceshellofanatom.
Draw it
p. 200

39. 6.2 Lewis Structures
A Lewis Structure tells us how electrons are shared in a molecule.
A Lewis Structure does NOT tell us about the shape of a molecule.

40. 5.2 Lewis Structures Practice Procedure: Lewis structure rules p. 201
H2O
More Practice
p. 202 # 1, 2
Even more Practice
p. 207 # 7 a-c
p. 217 # 30 a-e

41. 5.2 Lewis Structures: Multiple Bonds
Double Bond = 2 bonding pairs of electrons
Examples: O2
C2H4
A double bond has four electrons, so it is shorter and stronger than a single bond.

42. 5.2 Lewis Structures: Multiple Bonds
Skills Toolkit see p. 205
Practice Hint: only use multiple bonding if you have tried single bonding first
Triple Bond = 3 bonding pairs of electrons
Examples: N2
A triple bond has six electrons, so it is shorter and stronger than a single or double bond.

43. 5.2 Lewis Structures: Multiple Bonds
Skills Toolkit see p. 205
Practice Hint: only use multiple bonding if you have tried single bonding first
Practice:
p. 205 # 1,2
p a,b,c
Start per 3,5,6,

44. 6.2 Polyatomic Ions Learning Target
How do we draw a Lewis structure of a polyatomic ion?

45. 6.2 Lewis Structures of Polyatomic ions
see Practice Hint on p. 203
Example:
For positive charged ions subtract (add/subtract) electrons from the count
For negative charged ions add (add/subtract) electrons from the count
Ammonium ion
Sulfate ion
Practice:
p. 203 #1,2; p. 207 #7d

46. 6.2 Lewis Structure of Polyatomic Ions
Practice Draw the Lewis structures the all the important polyatomic ions we highlighted on our ions sheet. (see Practice Hint on p. 203)
Ammonium
Sulfate
Hydroxide
Nitrate
Nitrite
Cyanide
Peroxide
Sulfite
Carbonate
Bicarbonate
Phosphate

47. 6.2 Think about it……
Different non-metal elements have a tendency to form specific numbers of bonds to become isoelectronic with a noble gas. Take a look at all the Lewis structure examples we have done so far.
How many covalent bonds do the following elements tend to form?
Hydrogen?
Carbon?
Nitrogen?
Oxygen, Sulfur?
The halogens?
Start per 1,2

48. 6.3 Molecules # of Valence Electrons # of bonds Hydrogen 1 Carbon 4
Nitrogen 5 3
Oxygen, Sulfur 6 2
Halogens 7

49. 6.2 Diatomic molecules Only nobles gases are monatomic
H2 = Hydrogen
N2 = Nitrogen
F2 = Fluorine
O2= Oxygen
I2 = Iodine
Cl2 = Chlorine
Br2 = Bromine
Memory aid:
Have No Fear Of Ice Cold Beer
See Ions Sheet

50. 6.3 Resonance Structures
Draw the Lewis structure of ozone, O3. Notes: Anyoneoftwoormorepossibleconfigurationsofthesamecompoundthathaveidenticalgeometrybutdifferentarrangementsofelectrons. Any one of two or more possible configurations of the same compound that have identical geometry but different arrangements of electrons. Read Resonance Structures p. 206

51. 6.2 Drawing and Naming Molecules
Objectives
Name a molecule when given its formula
E.g. CO2  carbon dioxide
Write the formula of a molecule when given its name
E.g. carbon dioxide  CO2

52. 6.2 Naming rules Binary compounds p. 206
Less electronegative element is first
Name of second element ends with ‘ide’
Name of the second element is preceded by a prefix (see p. 207 Table 5)
Name of first element is preceded by a prefix only if there is more than one atom
Practice
p. 207 #9a-d, 10 a-d

53. 6.2 Names of Compounds Practice Name the following compounds: SiF3
PtF6
P2O3
V4O10
P. 217 # 25 a-e

54. 6.3 Learning Target
How can we use Lewis structures to predict the shape of molecules?
The aspirin molecule 

55. 6.2 Lewis Structures
A Lewis Structure tells us how electrons are shared in a molecule.
A Lewis Structure does NOT tell us about the shape of a molecule.

56. 6.3 Molecular Shapes Notes:
Valence Shell Electron Pair Repulsion Theory
Atheorythatpredictsmolecularshapesbasedontheideathatpairsofelectronssurroundinganatomrepeleachother
P. 209
‘Electron domain’

57. 6.3 Molecular Shapes Linear Bent Trigonal planar trigonal pyramidal
tetrahedral

58. 6.3 Molecular Shape Learning Target
Why is the geometry of the water molecule bent, not straight?
Repulsion between ‘electron domains’

59. 6.3 Molecular Shapes Linear geometry AB2 e.g. CO2 Carbon dioxide
Bond angle = 180o
Balloon demo

60. 6.3 Molecular Shapes Trigonal Planar geometry AB3 e.g.BCl3
Bond Angle = 120o

61. 6.3 Molecular Shapes tetrahedral geometry AB4
e.g. Methane (CH4), organic molecules
Bond Angle = 109.5o

62. 6.3 Molecular Shapes trigonal pyramid geometry
AB3 (3 bonding pairs, one lone pair)
Example: Ammonia, NH3
Bond Angle = 107o

63. 6.3 Molecular Shapes Bent geometry AB2 (2 bonding pairs, 2 lone pairs)
e.g. water, H2O
Bond Angle = 104.5o

64. 6.3 Molecular Shapes Remember…. According to VSEPR theory,
1. What is the name of the shape of the carbon dioxide (CO2) molecule?
2. What is the name of the shape of the methane (CH4) molecule?
PhET sim: polarity