1. TOPIC 4 CHEMICAL BONDING AND STRUCTURE
4.3
COVALENT STRUCTURES

2. ESSENTIAL IDEA
Lewis (electron dot) structures show the electron domains in the valence shell and are used to predict molecular shape.
NATURE OF SCIENCE (1.10)
Scientists use models as representatives of the real world – the development of the model of molecular shapes (VSEPR) to explain observable properties.

3. THEORY OF KNOWLEDGE
Does the need for resonance structures decrease the value or validity of Lewis (electron dot) theory? What criteria do we use in assessing the validity of a scientific theory?

4. UNDERSTANDING/KEY IDEA 4.3.A
Lewis (electron dot) structures show all the valence electrons in a covalently bonded species.

5. UNDERSTANDING/KEY IDEA 4.3.B
The “octet rule” refers to the tendency of atoms to gain a valence shell with a total of 8 electrons.

6. APPLICATION/SKILLS
Be able to deduce the Lewis (electron dot) structures of molecules and ions showing all valence electrons for up to four electron pairs on each atom.

7. Lewis Dot Structures
Lewis Dot structures are used to represent the valence electrons of atoms in covalent molecules
Dots are used to represent only the valence electrons.
Dots are written between symbols to represent bonding electrons 7

8. Writing Dot Structures
Writing Dot structures is a process:
Determine the number of valence electrons each atom contributes to the structure
The number of valence electrons can usually be determined by the column in which the atom resides in the periodic table 8

9. Writing Dot Structures
Example SO32-
1 S = 6 e
= 6x3 = 18 e
(2-) charge = 2 e
Total = 26 e
Add up the total number of valence electrons
Adjust for charge if it is a poly atomic ion
Add electrons for negative charges
Reduce electrons for positive charges 9

10. Electron Dot Structures
Make the atom that is fewest in number the central atom.
Distribute the electrons so that all atoms have 8 electrons.
Use double or triple pairs if you are short of electrons
If you have extra electrons put them on the central atom 10

11. Electron Dot Structures
Example 2: SO3
1 S = 6 e
3 O = 6x3 = 18 e
no charge = 0 e
Total = 24 e
Note: a double bond is necessary to give all atoms 8 electrons 11

12. Lewis Dot Structure for SO3
The diagram below shows the dot structure for sulfur trioxide. The bonding electrons are in shown in red and lone pairs are shown in blue. 12

13. Electron Dot Structures NH4+
Example 3: NH4+
1 N = 5 e-
4 H = 4x1 = 4 e-
(+) charge = -1 e-
Total = 8 e-
Note: Hydrogen atoms only need 2 e- rather than 8 e- 13

14. Example: Carbon Dioxide
1. Central atom =
2. Valence electrons =
3. Form bonds.
C e- O e- x 2 O’s = 12 e- Total: 16 valence electrons
This leaves 12 electrons (6 pairs).
4. Place lone pairs on outer atoms.
Check to see that all atoms have 8 electrons around it
except for H, which can have 2.

15. Carbon Dioxide, CO2
C 4 e How many are in the drawing? O 6 e- X 2 O’s = 12 e- Total: 16 valence electrons
There are too many electrons in our drawing. We need to take away 4 electrons and then form DOUBLE BONDS between C and O.
Instead of sharing only 1 pair, a double bond shares 2 pairs.
Take away the lone pairs from the carbon. So one pair is taken away from each oxygen atom and replaced with another bond.

16. LEWIS STRUCTURE RULES
Add up the total number of valence electrons in the molecule.
Draw the skeletal structure.
Use a line between each element to symbolize an electron pair.
Distribute the remaining electrons around the elements in pairs to form octets. (Hydrogen can only ever have 2 electrons.)
If you do not have enough to form octets, make double or triple bonds.
Ions must have square brackets around them with the charge notated in the top right hand corner.
To be a correct Lewis structure, ALL electrons must be shown.

17. UNDERSTANDING/KEY IDEA 4.3.C
Some atoms, like Be and B, might form stable compounds with incomplete octets of electrons.

18. Violations of the Octet Rule
Violations of the octet rule usually occur with B and elements of higher periods. Some common examples include: Be, B, P, S, and Xe.
SF4
Be: 4
B: 6
P: 8 OR 10
S: 8, 10, OR 12
Xe: 8, 10, OR 12
BF3

19. OCTET RULE EXCEPTIONS Hydrogen will never have more than 2 electrons.
Be and B have less than 8 electrons.
Some elements like S and P can have expanded octets which hold more than 8 electrons.
Coordinate covalent bonds are formed when both electrons originate from the same atom.
An arrow is used to denote the direction in a coordinate covalent bond showing the atom from which both electrons originated.
This is common in double and triple bonds.

20. GUIDANCE
Coordinate covalent bonds should be covered.

21. Coordinate Covalent Bonds
Coordinate covalent bonds occur when one atom donates both of the electrons that are shared between two atoms
Coordinate covalent
bonds are also called
Dative Bonds

22. UNDERSTANDING/KEY IDEA 4.3.D
Resonance structures occur when there is more than one possible position for a double bond in a molecule.

23. RESONANCE
Resonance is a concept used to describe the structures when there are multiple ways to depict the same molecule.
If you can put a double bond in more than one position, you will be expected to draw the resonance structures.
The electrons are actually delocalized in the areas of the double bonds and are spread out equally among all bonding positions.
Bond strength and length are in between that of single and double bonds.

24. Double arrows are placed between all resonance structures.
Resonance structures allow us to depict all the possible positions of the double bonds.
The true structure, however, is an intermediate form known as a resonance hybrid.
Double arrows are placed between all resonance structures.
Ref: myweb.astate.edu

25. APPLICATION/SKILLS
Be able to deduce resonance structures. Examples include but are not limited to C6H6, CO32- and O3.

26. BENZENE

27. CARBONATE

28. OZONE

29. UNDERSTANDING/KEY IDEA 4.3.E
Shapes of species are determined by the repulsion of electron pairs according to the VSEPR theory.

30. APPLICATION/SKILLS
Be able to use the VSEPR theory to predict the electron domain geometry and the molecular geometry for species with two, three and four electron domains.

31. VSEPR THEORY Valence Shell Electron Pair Repulsion theory.
Video
VSEPR THEORY
Valence Shell Electron Pair Repulsion theory.
States that in a small molecule, the pairs of valence electrons are arranged as far apart from each other as possible.
So far we have dealt with structural formulas which only show the types of atoms, bonds and lone pairs of electrons. They do not show the shape of the molecule.

32. 12 BASIC SHAPES
LINEAR – two atoms bonded to the central atom, no lone pairs of electrons on the central atom.
BENT- two atoms bonded to the central atom with one or two lone pairs of electrons on the central atom.

33. TRIGONAL PLANAR(flat triangle) – three atoms bonded to the central atom, no lone pairs of electrons on the central atom.
TRIGONAL PYRAMIDAL – three atoms bonded to the central atom, one lone pair of electrons on the central atom.

34. TETRAHEDRAL – four atoms bonded to the central atom, no lone pairs of electrons on the central atom.
TRIGONAL BIPYRAMIDAL – five atoms bonded to the central atom (octet rule exception)

35. SEE SAW (also called unsymmetrical tetrahedron) – derivative of the trigonal bipyramidal with one lone pair of electrons.
Note that on the three trigonal bipyramidal shape derivatives, the shapes come from pulling off atoms from the flat triangle, not the top and bottom.

36. T-SHAPED – also a derivative of the trigonal bipyramidal shape with 2 lone pairs of electrons on the central atom.
LINEAR – derivative of the trigonal bipyramidal shape with three lone pairs of electrons on the central atom.
This linear is different from the
other linear as it has 3 pairs of
lone electrons on the central atom.

37. OCTAHEDRAL – six atoms bonded to the central atom (another octet rule exception).
SQUARE PYRAMIDAL – derivative of the octahedral with one lone pair of electrons on the top or bottom.
SQUARE PLANAR – derivative of the octahedral shape with two lone pairs of electrons – one on top and one on bottom.

38. APPLICATION/SKILLS
Be able to predict molecular polarity from bond polarity and molecular geometry.

39. MOLECULAR POLARITY The polarity of a molecule depends upon:
The polar bonds it contains.
The shape of the molecule.
If the bonds are equally polar and arranged symmetrically, then they cancel each other out and are non-polar.
If the molecule contains bonds of different polarities or the bonds are not arranged symmetrically, then the molecule will be polar.
You can usually tell by the shape and lone pairs of electrons if the molecule is polar or not.

40. APPLICATION/SKILLS
Be able to predict the bond angle from molecular geometry and presence of non-bonding pairs of electrons.

41. ELECTRON DOMAINS
Double and triple bonded electron pairs are orientated together and behave as a single unit known as an electron domain.
Lone pairs also count as electron domains so the 12 VSEPR shapes are narrowed down to 5 basic shapes.
For the electron domain shapes, you will need to know bond angles and hybridization.

42. LINEAR A linear molecule has two electron domains.
The angle is 180 degrees and it has “sp” hybridization.
Non polar

43. TRIGONAL PLANAR A trigonal planar molecule has 3 electron domains.
It has angles of 120 degrees and “sp2” hybridization. Non polar
The bent molecule can also have 3 electron domains. Polar

44. TETRAHEDRAL A tetrahedral molecule has four electron domains.
It has angles of degrees and “sp3” hybridization. Non polar
Trigonal pyramidal and “bent” with 2 lone pairs also have four electron domains. Polar

45. TRIGONAL BIPYRAMIDAL
A trigonal bipyramidal molecule has 5 electron domains.
It has angles of 90 and 120 degrees and “dsp3” hybridization.
Non polar (see saw and t-shaped are polar)

46. OCTAHEDRAL An octahedral molecule has 6 electron domains.
It has angles of 90 degrees and “d2sp3” hybridization. Non polar (square pyramidal – polar)

47. GUIDANCE
Allotropes of carbon (diamond, graphite, graphene and C60 buckminsterfullerene) should be covered.

48. ALLOTROPES OF CARBON
Some covalent structures are crystalline in nature like ionic lattices; however, they are linked together with covalent bonds.
The crystal is a single molecule with a regular repeating pattern of covalent bonds.
It is referred to as a giant molecular structure.
Allotropes are different forms of an element in the same physical state.
Different bonding within the structures gives rise to different physical properties.
Carbon has four allotropes.

49. GRAPHITE
Graphite
Each C atom is sp2 hybridized covalently bonded to 3 others forming hexagons in parallel layers with bond angles of 120 degrees.
The layers are held together by van der Waal’s forces so they can slide over each other.
Density is 2.26 g/cm3
Contains one non-bonded, delocalized electron per atom so graphite conducts electricity due to the movement of these electrons.
Not a good heat conductor
Very high melting point, most stable allotrope
Non lustrous – grey solid
Used as a lubricant and in pencils

50. DIAMOND
Diamond
Each C atom is sp3 hybridized covalently bonded to 4 others tetrahedrally in a regular repeating pattern with bond angles of degrees.
It is the hardest known natural substance.
Density is 3.51 g/cm3
All electrons are bonded so it does not conduct electricity.
Does conduct heat better than metals.
Very high melting point, brittle
Lustrous crystal
Used in jewelry and tools

51. FULLERENE, C60
Fullerene
Each C atom is sp2 hybridized covalently bonded in a sphere of 60 carbon atoms consisting of 12 pentagons and 20 hexagons.
The structure is a closed spherical cage in which each carbon atom is bonded to 3 others.
Density is g/cm3
It easily accepts electrons to form negative ions so it is a semiconductor at normal temp and pressure due to some electron mobility.
Very low heat conductivity
Low melting point
Yellow crystalline solid – soluble in benzene
Related forms are used to make nanotubes
for the electronics industry,
catalysts and lubricants.

52. GRAPHENE
Graphene
Each C atom is sp2 hybridized covalently bonded to three other carbons forming hexagons with bond angles of 120.
The structure is a two dimensional single layer described as a honeycomb or chicken wire structure.
Density is 1.5 g/cm3
Contains one non-bonded, delocalized electron per atom so graphene conducts electricity due to the movement of these electrons.
Excellent heat conductor – better than diamonds
Very high melting point, thinnest material to exist, stronger than steel
Almost completely transparent
Used in touch screens, high performance electronics, etc

53. UNDERSTANDING/KEY IDEA 4.3.F
Carbon and silicon form giant covalent/network covalent structures.

54. APPLICATION/SKILLS
Be able to explain the properties of giant covalent compounds in terms of their structures.

55. Silicon, Si Silicon is a group 4 element with 4 valence electrons.
In its elemental state, each silicon atom is covalently bonded to four others in a tetrahedral arrangement.
This results in a giant lattice structure like a diamond.

56. Silicon dioxide SiO2
SiO2 commonly known as silica or quartz, forms a giant covalent structure.
It is also a tetrahedral structure, but the bonds are between Si and O.
Each Si atom is covalently bonded to 4 oxygen atoms and each O to two Si atoms.
The formula SiO2 is the ratio of atoms within the giant molecule.
The structure is strong, insoluble in water, does not conduct electricity or heat and has a high melting point – all properties of glass and sand which are different forms of silica.

57. Citations
International Baccalaureate Organization. Chemistry Guide, First assessment Updated Brown, Catrin, and Mike Ford. Higher Level Chemistry. 2nd ed. N.p.: Pearson Baccalaureate, Print. Most of the information found in this power point comes directly from this textbook. The power point has been made to directly complement the Higher Level Chemistry textbook by Catrin and Brown and is used for direct instructional purposes only.