1. COVALENT BONDS

2. Covalent clip

3. I.The Covalent Bond A. Why do atoms form bonds?
1. To get 8 valence shell electrons
(noble gas configuration)
2. more stable configuration -
less potential energy
B. Definition of Covalent Bond
1. bond resulting from sharing electrons
2. share electrons by overlapping of
orbitals

4. I.The Covalent Bond A. Why do atoms form bonds?
1. To get 8 valence shell electrons
(noble gas configuration)
2. more stable configuration -
less potential energy
B. Definition of Covalent Bond
1. bond resulting from sharing electrons
2. share electrons by overlapping of
orbitals

5. II.Heisenberg Uncertainty Principle
It is impossible to know exactly the position and velocity of a particle at the same time.
The better we know one, the less we know the other.
The act of measuring changes the properties.
Look at a fan

6. More obvious with the very small
To measure where a electron is, we use light.
But the light moves the electron
And hitting the electron changes the frequency of the light.
Watch the balloon

7. After Before Photon changes wavelength Photon
Electron Changes velocity
Moving Electron

8. Covalent Bonds – Sharing of Electrons

9. O The Lewis dot structure for Oxygen
Oxygen is in group VIA so it has 6 valence electrons

10. Cl The Lewis dot structure for Chlorine
chlorine is in group VIIA so it has 7 valence electrons

11. Ca The Lewis dot structure for calcium
calcium is in group IIA so it has 2 valence electrons

12. Making calcium chloride+ Ca Cl
Ca( Cl )2

13. Lewis dot structure of a compoundH
NH3 H N H

14. H O Water - H2O Each hydrogen has 1 valence electron
Each hydrogen wants 1 more
The oxygen has 6 valence electrons
The oxygen wants 2 more
They share to make each other happy H O

15. H O Water Put the pieces together The first hydrogen is happy
The oxygen still wants one more H O

16. H O H Water The second hydrogen attaches
Every atom has full energy levels H O H

17. O2
Oxygen is also one of the diatomic molecules

18. O
How will two oxygen atoms bond?

19. O
Each atom has two unpaired electrons

20. O

21. O

22. O

23. O

24. O

25. O

26. O Oxygen atoms are highly electronegative.
So both atoms want to gain two electrons.

27. O Oxygen atoms are highly electronegative.
So both atoms want to gain two electrons.

28. O

29. O O

30. O O

31. O O

32. Both electron pairs are shared.

33. O O
6 valence electrons
plus 2 shared electrons
= full octet

34. O O
6 valence electrons
plus 2 shared electrons
= full octet

35. O O
two bonding pairs,
making a double bond

36. O O =
For convenience, the double bond
can be shown as two dashes.

37. This is the oxygen molecule,=
this is so cool!!
This is the oxygen molecule, O2

38. This is the oxygen molecule,=
this is so cool!!
This is the oxygen molecule, O2

39. Types of Covalent Bonding
Nonpolar
Polar

40. Non polar and Polar Bonds
When the atoms in a bond are the same, the electrons are shared equally.
This is a nonpolar covalent bond.
When two different atoms are connected, the electrons may not be shared equally.
This is a polar covalent bond.
How do we measure how strong the atoms pull on electrons?

41. Electronegativity
A measure of how strongly the atoms attract electrons in a bond.
The bigger the electronegativity difference the more polar the bond.
Covalent nonpolar
Covalent polar
1.7> Ionic

42. Covalent:
Non-polar
0>0.3
Polar:
0.3 < 1.7
Ionic
> 1.7

44. Electronegativity
A. Nonpolar Covalent Bonds
1. equal sharing of electrons
2. difference in electronegativity is
( )
3. examples ( H O N Cl2)

45. B. Polar Covalent Bonds
1. unequal sharing of electrons
a. electrons spend more time by
one of the atoms
b. atoms take on a charge
(+dipole or - dipole)
c. difference in electronegativity is
( )
2. examples (HBr H2S CBr4 )

46. C. Ionic Nature of Covalent Bonds
1. polar covalent bonds have an
ionic nature
a. when ionic nature is greater
than 50% the bond is
considered ionic
1) electronegativity difference
is greater than 1.7

47. 3. diatomic molecules (contain two atoms)
a. periodic table (1,1-7--->group 17)
elements that form diatomic molecules
H N O F Cl Br2
b. compounds can also be diatomic
HCl CO NO HF HI HBr

48. C. Examples and Diagrams
H F HF
Make a drawing D. What is a molecule?
1. two or more atoms bonded covalently
2. examples- Make a drawing
(usually two or more nonmetals)
H2O NH3 CH4 N2O5 HCl C6H12O6

49. Isomers- Build the molecule
On the right using springs for double bonds.
Figure 3.32 – Page 312

50. Use molecular model kit to build 1. PH3 2. H2S 3. HCl 4. CCl4
Lewis structure
Use molecular model kit to build
1. PH H2S
3. HCl CCl4
5. SiH CH4
7. ClF PCl5
9. SO N2O5

51. Lewis structure
1. PH3
2. H2S
3. HCl
4. CCl4
5. SiH4

52. ClF3
PCl5
SO2
N2O5