1. Chapter 9
Covalent Bonding

2. 9.1 The Covalent Bond
Covalent Bond – A chemical bond that results from the sharing of valence electrons.
Between nonmetallic elements of similar electronegativity.
A group of atoms united by a covalent bond is called a molecule.
A substance made up of molecules is called a molecular substance.
Molecules may have as few as two atoms. (CO), (O2)

3. Covalent Bonds

4. Polar Covalent Bonds: Unevenly matched, but willing to share.

5. To describe the composition of a molecular compound we use a molecular formula.
The molecular formula tells you how many atoms are in a single molecule of the compound.
The molecular formula for oxygen is O2, which tells you that this molecule contains 2 atoms of oxygen.

6. The molecular formula for sucrose (table sugar) is C12H22O11
The molecular formula for sucrose (table sugar) is C12H22O11. This indicates that one sucrose molecule contains 12 carbon atoms, 22 hydrogen atoms and 11 oxygen atoms.
The empirical formula shows the ratio of atoms in the molecule. You get it by dividing all subscripts by the smallest subscript and rounding. C6H12O6 becomes CH2O.

7. Molecules are often shown using a structural formula.
A structural formula specifies which atoms are bonded to each other in a molecule.
One type is the Lewis structures.
The Lewis structure uses dots as before and has a circle drawn around the dots that represent valence electrons that are joined in a covalent bond.

8. The principle that describes covalent bonding is the Octet rule.
Atoms within a molecule will share electrons so as to acquire a full set (8) of valence electrons.
If each atom has 7 valence electrons, than they will share one so that they both will have 8. The sharing of that electron is a Single covalent bond.

9. Single Covalent bonding
Fluorine has seven valence electrons
A second atom also has seven
By sharing electrons
Both end with full orbitals F F

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

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

12. If more than two atoms join in a covalent bond the atoms with the fewest valence electrons will pair up with the atom that has the most up to a total of 8.
If a pair of electrons is not shared they are called an unshared pair.
If two atoms share two pairs of electrons it is called a double covalent bond.
If two atoms share three pairs of electrons it is called a triple covalent bond.

13. The Lewis structures that you have seen so far use a pair of dots to represent covalent bonds.
Another kind of Lewis structure uses dashes to represent covalent bonds.
A single dash represents a single covalent bond.
A double dash represents a double covalent bond
A triple dash represents a triple covalent bond.

14. Covalent compounds (molecules) tend to have low melting and boiling points
When an electron pair is shared by the direct overlap of bonding orbitals, a sigma bond results. Single bonds are sigma bonds.
The overlap of parallel orbitals forms a pi bond. Multiple bonds involve both sigma and pi bonds.

15. Bond length depends on the sizes of the bonded atoms and the number of electron pairs they share.
Bond dissociation energy is the energy needed to break a covalent bond.
Double bonds have larger bond dissociation energies than single. Triple even larger
Bond length and bond dissociation energy are directly related.

16. 9.2 Naming Molecules
Covalent compounds (Molecules) are named by adding prefixes to the element names.
Covalent’ (in this context) means both elements are nonmetals.
The names of molecular compounds will include prefixes that indicate the number of atoms in the molecule. (CO2 would be called Carbon Dioxide, CCl4 is called Carbon Tetrachloride)
The following prefixes are used:
mono-1 di-2
tri-3 tetra-4
penta-5 hexa-6
hepta-7 octa-8
nona-9 deca-10
Note: When a prefix ending in ‘o’ or ‘a’ is added to ‘oxide’, the final vowel in the prefix is dropped.

17. Remember to use the “ide” suffix. (Cl is chloride not chlorine)
Exceptions to the rule:
The prefix “mono” is usually not written with the first word of the compounds name.
CO2 is not called monocarbon dioxide, just carbon dioxide.
Some compounds have common names. (H2O is water or dihydrogen monoxide)

18. Naming Binary Covalent Compounds: Examples1
mono 2 di 3
tri 4
tetra 5
penta 6
hexa 7
heptaa 8
octa 9
nona 10
deca
* Second element in ‘ide’ from
* Drop –a & -o before ‘oxide’
N2S4
dinitrogen tetrasulfide
NI3
nitrogen triiodide
XeF6
xenon hexafluoride
CCl4
carbon tetrachloride
P2O5
diphosphorus pentoxide
SO3
sulfur trioxide

19. Naming Acids
An acid is a molecular substance that dissolves in water to produce hydrogen ions (H+).
Acids behave like an ionic compound because when they dissolve in water, they create cations and anions.
The cation is always H+ and the anion depends on the particular acid.
Example: When the acid HCL dissolves in water it forms H+ cations and CL- anions.
HNO3 acid will form H+ and NO3- ions.

20. Naming Acids
The name of an acid usually results from the name of the anion(-).
For acids, where the anion ends with the suffix “ide”(Chloride) the name of the acid begins with the prefix hydro.
The acids name also includes the root name of the anion(-) and the word acid.
In addition you need to change the suffix of the anion from “ide” to “ic”.
By these rules HCL is named hydrochloric acid.
Other acids are named without the prefix “hydro”. The acid HNO3- derives its name from NO3- which is nitrate. HNO3 is named Nitric Acid.
The following are names and formulas of common acids:

21. Naming Acids Anion Corresponding Acid
F-, Fluoride HF, hydrofluoric Acid
CL-, Chloride HCL, hydrochloric acid
Br-, Bromine HBr, hydrobromic acid
I-, iodide HI, hyroiodic acid
S2-, sulfide H2S, hydrosulfuric acid
NO3-, nitrate HNO3, nitric acid
CO3-2, carbonate H2CO3, carbonic acid
SO4-2, sulfate H2SO4, sulfuric acid
PO4-3, phoshate H3PO4, Phosphoric acid
C2H3O2- , acetate HC2H3O2, acetic acid

22. 9.3 Molecular Structures
Structural formulas (Lewis Structures) use letter symbols and bonds to show the position of atoms.
Drawing Lewis Structures:

23. H O Water Each hydrogen has 1 valence electron and wants 1 more
The oxygen has 6 valence electrons
and wants 2 more
They share to make each other “happy” H O

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

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

26. 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.

27. C O Carbon dioxide CO2 - Carbon is central atom ( I have to tell you)
Carbon has 4 valence electrons
Wants 4 more
Oxygen has 6 valence electrons
Wants 2 more C O

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

29. Carbon dioxide
Attaching the second oxygen leaves both oxygen 1 short and the carbon 2 short 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. Carbon dioxide
The only solution is to share more O C O

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

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

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

36. 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

37. 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

38. 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

39. 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

40. How to draw them To figure out if you need multiple bonds
Add up all the valence electrons.
Count up the total number of electrons to make all atoms happy.
Subtract.
Divide by 2
Tells you how many bonds - draw them.
Fill in the rest of the valence electrons to fill atoms up.

41. N H Examples NH3 N - has 5 valence electrons wants 8
H - has 1 valence electrons wants 2
NH3 has 5+3(1) = 8
NH3 wants 8+3(2) = 14
(14-8)/2= 3 bonds
4 atoms with 3 bonds N H

42. H H N H Examples Draw in the bonds All 8 electrons are accounted for
Everything is full H H N H

43. Examples HCN C is central atom N - has 5 valence electrons wants 8
C - has 4 valence electrons wants 8
H - has 1 valence electrons wants 2
HCN has = 10
HCN wants = 18
(18-10)/2= 4 bonds
3 atoms with 4 bonds -will require multiple bonds - not to H

44. H C N HCN Put in single bonds Need 2 more bonds
Must go between C and N H C N

45. H C N HCN Put in single bonds Need 2 more bonds
Must go between C and N
Uses 8 electrons - 2 more to add H C N

46. H C N HCN Put in single bonds Need 2 more bonds
Must go between C and N
Uses 8 electrons - 2 more to add
Must go on N to fill octet H C N

47. Another way of indicating bonds
Often use a line to indicate a bond
Called a structural formula
Each line is 2 valence electrons H O H H O H =

48. H C N H C O H Structural Examples
C has 8 electrons because each line is 2 electrons
Ditto for N
Ditto for C here
Ditto for O
H C N H
C O H

49. Coordinate Covalent Bond
A Coordinate Covalent Bond is when one atom donates both electrons in a covalent bond.
Carbon monoxide CO O C

50. Coordinate Covalent Bond
When one atom donates both electrons in a covalent bond.
Carbon monoxide CO C O

51. Coordinate Covalent Bond
When one atom donates both electrons in a covalent bond.
Carbon monoxide CO C O O C

52. Resonance
Resonance is a condition when more than one dot diagram with the same connections is possible.
Resonance structures differ only in the position of the electron pairs, never the atom positions.

53. 9.4 Molecular Shape (VSEPR)
Valence Shell Electron Pair Repulsion.
Predicts three dimensional geometry of molecules.
The 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.
And the shape of molecules

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

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 basic 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. C H O H C O H 3 atoms no lone pair
The farthest you can the electron pair apart is 120º.
Shape is flat and called trigonal planar.
Will require 1 double bond C H O
120º H C O H

60. 2 atoms no lone pair
With three atoms the farthest they can get apart is 180º.
Shape called linear.
Will require 2 double bonds or one triple bond C O
180º

61. Molecular Orbitals
The overlap of atomic orbitals from separate atoms makes molecular orbitals
Each molecular orbital has room for two electrons
Two types of MO
Sigma ( σ ) between atoms
Pi ( π ) above and below atoms

62. Sigma bonding orbitals
From s orbitals on separate atoms +
+ + + + +
Sigma bonding molecular orbital
s orbital
s orbital

63. Sigma bonding orbitals
From p orbitals on separate atoms  
p orbital
p orbital    
Sigma bonding molecular orbital

64. Pi bonding molecular orbital
Pi bonding orbitals
P orbitals on separate atoms        
Pi bonding molecular orbital

65. Sigma and pi bonds All single bonds are sigma bonds
A double bond is one sigma and one pi bond
A triple bond is one sigma and two pi bonds.

66. Hybridization
The mixing of several atomic orbitals to form the same number of hybrid orbitals.
When an atom approaches another atom to form a bond, the orbitals of its electrons may be changed.
The changing of the orbitals forms what is called hybrid orbitals. This is a cross between all of the orbitals involved (s,p,d,f).

67. Hybridization
When an “s” and “p” orbital come together they form a linear shaped molecule.
When an “s” and “p2” orbital come together they form a trigonal planer molecule.
When an “s” and “p3” orbital come together they form tetrahedral, pyramidal or bent molecules.
Atoms with higher orbital numbers form longer bonds. Multiple bonds are shorter than single bonds because bonds with more electrons attract the nuclei of the bonding atoms more strongly.

68. 9.5 Electronegativity and Bond Type
Difference
Bond Type
<.4
Non-Polar Covalent
0.5 – 1.9
Polar Covalent
>2.0
Ionic

69. 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.

70. 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 moderately polar
Covalent polar
>2.0 Ionic

71. 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

72. Polar Molecules
Polar molecules have a partially positive end and a partially 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.

73. Polar Molecules
Symmetrical shapes are those without lone pair on central atom
Tetrahedral
Trigonal planar
Linear
Will be nonpolar if all the atoms are the same
Shapes with lone pair on central atom are not symmetrical
Can be polar even with the same atom

74. Intermolecular Forces
They are what make solid and liquid molecular compounds possible.
The weakest are called van der Waal’s forces - there are two kinds
Dispersion forces
Dipole Interactions

75. Dispersion Force F2 is a gas Br2 is a liquid I2 is a solid
Depends only on the number of electrons in the molecule
Bigger molecules more electrons
More electrons stronger forces
F2 is a gas
Br2 is a liquid
I2 is a solid

76. Dipole interactions
Occur when polar molecules are attracted to each other.
Slightly stronger than dispersion forces.
Opposites attract but not completely hooked like in ionic solids.

77. Dipole interactions H F d+ d- H F d+ d-
Occur when polar molecules are attracted to each other.
Slightly stronger than dispersion forces.
Opposites attract but not completely hooked like in ionic solids.
H F
d+ d-
H F
d+ d-

78. Properties of Molecular Compounds
Made of nonmetals
Poor or nonconducting as solid, liquid or aqueous solution
Low melting point
Two kinds of crystals
Molecular solids held together by IMF
Network solids- held together by bonds
One big molecule (diamond, graphite)