1. Ch. 8 Covalent Bonding
Chemistry

2. Background Information
Covalent Bonding
Homework: p. 229 #14-16, 18-21, p. 247 #43-44, 46-47, 49-50, 64
DUE FRIDAY!!!!!!!

3. Review Remember an ionic bond?
Holding together of cations and anions by electrostatic forces, or the “giving and taking” of electrons
So if an ionic bond is the giving and taking of electrons, what would a covalent bond be?
Molecule held together by sharing electrons
Another way to achieve stability and have both electrons achieve a full octet

4. Molecules Molecule – group of atoms (neutral) joined by covalent bonds
So what would a diatomic molecule be?
A molecules consisting of two of the same atoms
form a “seven” on the periodic table and hydrogen

5. Molecules Diatomic molecules occur naturally together
you will NOT find just one of the lone atoms by itself
Ex: Cl2, NOT Cl (if we are talking about it occurring naturally)

6. Molecular Formula What is a chemical formula?
Shows kinds and numbers of atoms in a compound
What would a molecular formula be then?
Shows how many atoms of each element are in a molecule

7. Molecular Formula
It’s basically a chemical formula, for covalent (molecular) compounds
One difference: MOLECULAR FORMULAS ARE NOT REDUCED TO THE LOWEST WHOLE NUMBER RATIOS!!!
Ex: NH3 (ammonia) or C2H6 (ethane)
What is one more difference?
Not ions, sharing of bonds (covalent)

8. Objective 1
Demonstrate covalent bonding using electron dot structures and orbital diagrams
Homework: p. 229 #14-16, 18-21, p. 247 #43-44, 46-47, 49-50, 64

9. Rules
1. Decide which element goes in the “middle” (usually least electronegative)
What are the trends again?
Increase from bottom left to upper right
2. Draw valence electrons around each atom
Count how many valence electrons each element has (remember to add/subtract for ions!)
3. Draw covalent bonds between unpaired electrons.
4. Remember: C, N, O, S can double bond
C, N, and O NEVER disobey the octet rule
5. Disobey the octet rule only if you have to!
Place extra electrons on the central atom – it will be the one that disobeys the octet rule

10. Electron Dot Diagrams Remember these?
What would the electron dot diagram for C be?
For Cl?
Why were electron dot diagrams useful for ionic compounds?
Show which electrons are accepted/ donated and shows bonding

11. Electron Dot Diagrams
So why do you think electron dot diagrams are useful for covalent (molecular) compounds?
To show which electrons are shared!
Let’s try Hydrogen!
Electron dot diagram:
Molecular compound:
This shared pair of
electrons is called a
pair or shared electrons

12. Electron Dot Diagrams Electron dot diagram of Fluorine:
Molecular Compound:
These valence electrons not shared (and not in bonding are referred to as…
Unshared electrons (or pair)

13. Let’s try another
How would we do water?

14. Let’s try another
How many shared pairs: 2
Unshared:

15. Let’s try another
How would we do methane?

16. Let’s try another
Shared pairs: 4
Unshared pairs:
None!

17. What about one not bonding with H?
Let’s try BF3
Shared pairs: 3
Non shared
On B – none
On F - 9

18. Here’s a tough one Let’s try making O2
Remember: all electrons MUST be paired
What did you notice?
A double bond needed to be formed
Will a double bond be more or less stable than a single bond?

19. Here’s a tough one
Double bond – a bond that shares two pairs of electrons

20. Here’s another
Let’s try N2
A triple bond

21. Here’s one more:
CO2
Atoms don’t have to be the same type of element to participate in double or triple bond
Hydrogen will never participate. Why?
Only 1 electron!

22. Another: Let’s try CO Both electrons were shared from the O
Coordinate covalent bond - When one atoms contributes both bonding electrons

23. With Ions
Let’s try NH4+
Remember to take away one electron for the +1 charge!

24. With Ions
Let’s try OH-1
Remember to add one electrons for the -1 charge!

25. Another… PCl5 Forced to disobey the octet!
Octet rule cannot be satisfied in molecules who have an odd number of valence electrons

26. One More BF3 (again) What do you notice?
It cannot obey the octet rule because the “middle” atom does not have enough electrons!

27. Ozone O3 Where do you put the double bond?
It doesn’t matter!
Resonance structure – occurs when it is possible to draw 2 or more valid electron dot structures
We are moving ELECTRONS, NOT ATOMS!

28. Another example of resonance

29. Objective 2 Use the VSEPR Theory to predict the shapes of molecules

30. VSEPR Theory
Valence Shell Electron Pair Repulsion Theory – what does it sound like it is?
Repulsion between electron pairs cause molecular shapes to adjust so valence electron pairs are as far apart as possible
Determines the 3D shape of a molecule
Think of a magnet – if you put two negative sides together, what happens?

31. VSEPR Theory Does water look like this? Or like this?
How do you know? Why do you think it looks like this?
Places atoms and valence electrons off the central O as far away as possible

32. VSEPR Theory
Both bonding electrons and non-bonding electrons (off the central atom) are important for VSEPR
Possible molecular shapes:

33. VSEPR Theory How can I tell what molecular shape an atom is?
Geometry not considering lone pairs
Geometry considering lone pairs
How many single bonds off the central atom?

34. VSEPR Theory How can I tell what molecular shape an atom is?
Dashed line means these electrons are away from you
Wedged line means these electrons are towards you

35. VSEPR Theory How can I tell what molecular shape an atom is?
Bond Angle

36. VSEPR Theory Let’s try methane (CH4) Central atom? Valence Electrons?
Draw Bonds?
Lone pairs off central atom?
none
Molecular Geometry?
tetrahedral
Bond angle?
109.5

37. VSEPR Theory So let’s try ammonia(NH3) Central atom?
Valence Electrons? 5
Draw Bonds?
Lone pairs off central atom?
One pair
Molecular Geometry?
Trigonal pyramidal
Bond angle?
107

38. VSEPR Theory
Remember those double and triple bonds? We can draw those in VSEPR theory too
Only the first bond is counted for bonding purposes (i.e.: a double and triple bond is counted only once)
The first bond/single bond – sigma bond (s)
Second/third bond – pi bond (p)

39. VSEPR Theory Remember those double and triple bonds? Let’s try CO2
Central atom? C
Valence Electrons?
C – 4
Draw Bonds?
Lone pairs off central atom?
none
Molecular Geometry?
linear
Bond angle?
180
sigma pi

40. Hybridization
Hydridization – mix of several atomic orbitals (s and p subshells) to form hybrid oribals
Describes the type of bonds formed, which the VSEPR theory CANNOT do!!

41. 8.3
Hybridization
In methane, each of the four sp3 hybrid orbitals of carbon overlaps with a 1s orbital of hydrogen.

42. Hybrid Orbitals 8.3 Hybridization Involving Double Bonds
In an ethene molecule, two sp2 hybrid orbitals from each carbon overlap with a 1s orbital of hydrogen to form a sigma bond. The other sp2 orbitals overlap to form a carbon–carbon sigma bond. The p atomic orbitals overlap to form a pi bond. Inferring What region of space does the pi bond occupy relative to the carbon atoms?

43. Hybrid Orbitals 8.3 Hybridization Involving Triple Bonds
In an ethyne molecule, one sp hybrid orbital from each carbon overlaps with a 1s orbital of hydrogen to form a sigma bond. The other sp hybrid orbital of each carbon overlaps to form a carbon–carbon sigma bond. The two p atomic orbitals from each carbon also overlap. Interpreting Diagrams How many pi bonds are formed in an ethyne molecule?

44. Hybrid Orbitals sp3 – four things “attached” to central atom
sp2 - three things attached, double bonds
sp - two things attached, triple bonds
Let’s try some….
sp3
sp sp

45. Hybrid Orbitals
A few more…
sp2 sp

46. Objective 3 Identify compounds as non-polar, polar, or covalent
p. 244 #36-37 p. 247 #57-60, 70

47. Review - electronegativity
Remember electronegativity?
Increases across, decreases down
Electronegativity is important in differentiating between covalent and ionic bonds
Covalent bonds – sharing of electrons
Ionic bonds – give/take of electrons
Let’s consider NaCl
Which atom is more electronegative?
Cl (3.0) vs. Na (0.9)
By a lot or a little?
A lot
Do you think this would be covalent or ionic? Why?
Ionic – big difference between electronegativities

48. Bond Polarity
Nonpolar covalent bonds – bonding electrons are shared equally
Polar covalent bonds – covalent bond between atoms, electrons shared unequally
Ionic bonds – complete giving/taking of electrons
Will the more or less electronegative atom attract the electrons better?
more electronegative atoms – slight/partial negative charge
less electronegative atom – slight/partial positive charge

49. Bond Polarity
Nonpolar covalent bonds – bonding electrons are shared equally
Polar covalent bonds – covalent bond between atoms, electrons shared unequally
Ionic bonds – complete giving/taking of electrons

50. Bond Polarity
Nonpolar covalent bonds – bonding electrons are shared equally (less than 0.4)
Polar covalent bonds – covalent bond between atoms, electrons shared unequally ( )
Ionic bonds – complete giving/taking of electrons (greater than or equal to 2.1)

51. Bond Polarity
Special “arrows” are used to show the most electronegative side of the bond OR
“delta” signs denotes the partial positive or negative charge

52. Bond Polarity Let’s try one: CO2 EN of C: EN of O: Draw structure: 2.5
3.5
Draw structure:

53. Molecular Polarity
If a polar bond is an unequal sharing of electrons, what would a polar molecule be?
One end of the molecule is more negative, one end is positive
Nonpolar molecule?
The molecule has no polarity between bonds
OR molecule has polarity between bonds, but the molecule is symmetrical!

54. Molecular Polarity Let’s look at CO2 EN of C: EN of O:
2.5
EN of O:
3.5
Does it have polar bonds?
yes
Is it symmetrical?
Polar or nonpolar molecule?
Non polar

55. Molecular Polarity Let’s look at N2 EN of N: Does it have polar bonds?
3.0
Does it have polar bonds? no
Polar or nonpolar molecule?
Non polar

56. Molecular Polarity Let’s look at H2O EN of H: EN of O:
2.1
EN of O:
3.5
Does it have polar bonds?
yes
Is it symmetrical ?
No, look at the Lewis structure otherwise you will think WRONG
Polar or nonpolar molecule?
polar

57. Molecular Polarity (click me) Polar bonds? Yes Symmetrical?
Yes – polar molecule
No – non polar molecule No
Non polar
(click me)

58. Molecular Polarity
Its’s hard to tell if a molecule is symmetrical or not.
Here’s some hints
Tetrahedral, linear, trigonal planer, trigonal bipyramidal, and square planer will always be nonpolar if the elements off the central atom are the same
Do dipoles cancel out?

59. Molecular Polarity Let’s look at BF3 EN of F: EN of B:
4.0
EN of B:
2.0
Does it have polar bonds?
yes
Is it symmetrical ?
Polar or nonpolar molecule?
Non-polar

60. Objective 4
Given the characteristics of ionic and covalent compounds, complete Table 8.4

61. Ionic vs. Covalent

62. Ionic vs. Covalent
Why are ionic compounds good conductors and covalent poor?
Hint: electricity flows through ions
Ionic compounds have ions, covalent do not
Why are ionic compounds usually more soluble than covalent compounds in water?
Hint: think of electronegativities in water
Water is very polar, ionic compounds have ions, ions dissociate in water easily