1. Chemical Bonds
Chemistry
Chapter 6

2. Bonding Electrons
The electrons responsible for the chemical properties of atoms are those in the outer energy level.
Valence electrons - The s and p electrons in the outer energy level.

3. Keeping Track of Electrons
Atoms in the same column
Have the same outer electron configuration.
Have the same valence electrons.
Easily found by looking up the group number on the periodic table.
Main group number indicates valence electrons.

4. X Electron Dot diagrams A way of keeping track of valence electrons.
How to write them
Write the symbol.
Put one dot for each valence electron
Pair-up s, then place p’s (follow Hund’s rule) X

5. Electron Dot diagram for Nitrogen
Nitrogen has 5 valence electrons.
First we write the symbol.
Then add 1 electron at a time to each side. N
Until they are forced to pair up (Hund’s rule).

6. Stable Electron Configurations
All atoms react to achieve noble gas configuration.
Noble gases have 2 s and 6 p electrons.
8 valence electrons .
Also called the octet rule. Ar

7. Bonding Ionic – giving or taking of electrons
Covalent – sharing of electrons
Metallic – delocalizing electrons

8. Electron Dot – Ionic BondingNa Cl

9. Na+
Cl-

10. Covalent Bonding + + How does Hydrogen bond with itself?
The nuclei repel + +

11. But they are attracted to electrons
They share the electrons + +

12. Electron Dot Bonding – Covalent Bonding
Fluorine has seven valence electrons F

13. Fluorine has seven valence electrons
A second atom also has seven F F

14. F F Fluorine has seven valence electrons A second atom also has seven
By sharing electrons F F

15. F F Fluorine has seven valence electrons A second atom also has seven
By sharing electrons
Both end with full orbitals F F

16. F F Fluorine has seven valence electrons A second atom also has seven
By sharing electrons
Both end with full orbitals F F
8 Valence electrons

17. F F Fluorine has seven valence electrons A second atom also has seven
By sharing electrons
Both end with full orbitals F F
8 Valence electrons

18. Covalent Single Bond A sharing of two valence electrons.
Only nonmetals and Hydrogen.
Sigma Bonds

19. H O Water 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

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

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

22. Covalent Multiple Bonds
Sometimes atoms share more than one pair of valence electrons.
A double bond is when atoms share two pair (4) of electrons.
A triple bond is when atoms share three pair (6) of electrons.
Pi bonds – any bond past the original single bond (sigma)

23. Carbon dioxide
CO2 - Carbon is central atom ( least electronegative)
Carbon has 4 valence electrons
Wants 4 more
Oxygen has 6 valence electrons
Wants 2 more C O

24. Carbon dioxide
Attaching 1 oxygen leaves the oxygen 1 short and the carbon 3 short C O

25. Carbon dioxide
Attaching the second oxygen leaves both oxygen 1 short and the carbon 2 short O C O

26. Carbon dioxide
The only solution is to share more O C O

27. Carbon dioxide
The only solution is to share more O C O

28. Carbon dioxide
The only solution is to share more O C O

29. Carbon dioxide
The only solution is to share more O C O

30. Carbon dioxide
The only solution is to share more O C O

31. Carbon dioxide
The only solution is to share more O C O

32. O C O Carbon dioxide The only solution is to share more
Requires two double bonds
Each atom gets to count all the atoms in the bond O C O

33. O C O Carbon dioxide The only solution is to share more
Requires two double bonds
Each atom gets to count all the atoms in the bond
8 valence electrons O C O

34. O C O Carbon dioxide The only solution is to share more
Requires two double bonds
Each atom gets to count all the atoms in the bond
8 valence electrons O C O

35. O C O Carbon dioxide The only solution is to share more
Requires two double bonds
Each atom gets to count all the atoms in the bond
8 valence electrons O C O

36. Electron Dot Diagrams Calcium and bromine Potassium and iodine
Nitrogen and nitrogen
Boron and fluorine
Carbon and chlorine
Hydrogen and hydrogen
Lithium and fluorine
Chlorine and chlorine

37. Ionic Bonding
Ionic bond – bond created by the giving or taking of electrons.
Cations (metals) and anions (non-metals) are held together by opposite charges.
Ionic compounds are called salts.
Simplest ratio is called the formula unit.
Difference in electronegativities > 1.67 create an ionic bond.

38. Properties of Ionic Compounds
Crystalline structure.
A regular repeating arrangement of ions in the solid.
Ions are strongly bonded.
Structure is rigid.
High melting points- because of strong forces between ions.

39. Do they Conduct? Conducting electricity is allowing charges to move.
In a solid, the ions are locked in place.
Ionic solids are insulators.
When melted, the ions can move around.
Melted ionic compounds conduct.
First get them to 800ºC.
Dissolved in water they conduct.

40. Crystalline structure

41. Ionic solids are brittle+ -

42. Ionic solids are brittle
Strong Repulsion breaks crystal apart. + - + - + - + -

43. Covalent bonds Nonmetals hold onto their valence electrons.
They can’t give away electrons to bond.
Still want noble gas configuration.
Get it by sharing valence electrons with each other.
By sharing both atoms get to count the electrons toward noble gas configuration.

44. Polar Bonds
Nonpolar covalent bond - when the atoms in a bond are the same, the electrons are shared equally.
Polar covalent bond - when two different atoms are connected, the atoms may not be shared equally.
How do we measure how strong the atoms pull on electrons?

45. How to show a bond is polar
Isn’t a whole charge just a partial charge
d+ means a partially positive
d- means a partially negative
The Cl pulls harder on the electrons
The electrons spend more time near the Cl d+ d- H Cl

46. Electronegativity
A measure of how strongly the atoms attract electrons in a bond.
The bigger the electronegativity difference the more polar the bond.
Nonpolar Covalent
Polar Covalent
>1.67 Ionic

47. Electronegativity Bonding
Calcium and bromine
Potassium and iodine
Carbon and oxygen
Nitrogen and oxygen
Iron and oxygen
Hydrogen and hydrogen
Lithium and fluorine
Chlorine and chlorine

48. Metallic Bonds How atoms are held together in the solid.
Metals hold onto there valence electrons very weakly.
Think of them as positive ions floating in a sea of electrons.

49. Sea of Electrons + Electrons are free to move through the solid.
Metals conduct electricity. +

50. Metals are Malleable Hammered into shape (bend).
Ductile - drawn into wires.

51. Malleable +

52. Malleable
Electrons allow atoms to slide by. + + + + + + + + + + + +

53. VSEPR Valence Shell Electron Pair Repulsion.
Predicts three dimensional geometry of molecules.
Name tells you the theory.
Valence shell - outside electrons.
Electron Pair repulsion - electron pairs try to get as far away as possible.
Can determine the angles of bonds.

54. VSEPR
Based on the number of pairs of valence electrons both bonded and unbonded.
Unbonded pair are called lone pair.
CH4 - draw the structural formula
Has 4 + 4(1) = 8
wants 8 + 4(2) = 16
(16-8)/2 = 4 bonds

55. H H C H H VSEPR Single bonds fill all atoms.
There are 4 pairs of electrons pushing away.
The furthest they can get away is 109.5º. H H C H H

56. H C H H H 4 atoms bonded Basic shape is tetrahedral.
A pyramid with a triangular base.
Same shape for everything with 4 pairs. H
109.5º C H H H

57. N H N H H H H H 3 bonded - 1 lone pair
Still basic tetrahedral but you can’t see the electron pair.
Shape is called trigonal pyramidal. N H N H H H
<109.5º H H

58. O H O H H H 2 bonded - 2 lone pair
Still basic tetrahedral but you can’t see the 2 lone pair.
Shape is called bent. O H O H
<109.5º H H

59. 3 atoms no lone pair
The farthest you can the electron pair apart is 120º H C O H

60. H H C C O H O H 3 atoms no lone pair
The farthest you can the electron pair apart is 120º.
Shape is flat and called trigonal planar. H
120º H C C O H O H

61. 2 atoms no lone pair
With three atoms the farthest they can get apart is 180º.
Shape called linear.
180º O C O

62. Combines bonding with geometry
Hybrid Orbitals
Combines bonding with geometry

63. Hybridization
The mixing of several atomic orbitals to form the same number of hybrid orbitals.
All the hybrid orbitals that form are the same.
sp3 -1 s and 3 p orbitals mix to form 4 sp3 orbitals - tetrahedral.
sp2 -1 s and 2 p orbitals mix to form 3 sp2 orbitals leaving 1 p orbital – trigonal planar.
sp -1 s and 1 p orbitals mix to form 2 sp orbitals leaving 2 p orbitals - linear.

65. sp3 geometry This leads to tetrahedral shape.
Every molecule with a total of 4 atoms and lone pair is sp3 hybridized.
Gives us trigonal pyramidal and bent shapes also.
109.5º

66. How we get to hybridization
We know the geometry from experiment.
We know the orbitals of the atom
hybridizing atomic orbitals can explain the geometry.
So if the geometry requires a tetrahedral shape, it is sp3 hybridized.
This includes bent and trigonal pyramidal molecules because one of the sp3 lobes holds the lone pair.

67. sp2 hybridization C2H4 double bond acts as one pair trigonal planar
Have to end up with three blended orbitals
use one s and two p orbitals to make sp2 orbitals.
leaves one p orbital perpendicular

69. H H C C H H

70. What about two sp hybridization when two things come off
one s and one p hybridize
linear

71. CO2 C O

72. N2

73. Polar Molecules
Molecules with ends

74. Polar Molecules Molecules with a positive and a negative end
Requires two things to be true
The molecule must contain polar bonds
This can be determined from differences in electronegativity.
Symmetry can not cancel out the effects of the polar bonds.
Must determine geometry first.

75. Is it polar? HF
H2O
NH3
CCl4
CO2