1. Chapters 6 and 16 Covalent Bonding
Objectives – Section 16.1
Use electron dot structures to show the formation of single, double and triple covalent bonds
Describe and give examples of coordinate covalent bonding, resonance structures, and exceptions to the octet rule.

2. CA Standards
Students know atoms combine to form molecules by sharing electrons to form covalent or metallic bonds or by exchanging electrons to form ionic bonds.
Students know chemical bonds between atoms in molecules such as H2, CH4, NH3, H2CCH2, N2, Cl2, and many large biological molecules are covalent.
Students know how to draw Lewis dot structures.

3. Single Covalent Bonds
Let’s look at Hydrogen as the simplest model of a covalent bond
H· ·H H:H
Hydrogen Hydrogen Hydrogen
atom atom molecule
H : H
Each Hydrogen has one electron and they share them to form a single covalent bond.
The single covalent bond can be represented by the pair of electrons or as a dash as shown below
H:H or H-H
Each dash represents a pair of shared electrons.

4. Conventions for naming
The chemical formulas of ionic compounds describe formula units (Example: NaCl is a formula unit)
The chemical formulas of covalent compounds describe molecules. (Example H2O is a molecule)
Ionic compounds do not have molecular formulas because they are not composed of molecules.
What does that mean?
Example: Ionic copper(II) oxide is composed of equal numbers of Cu2+ and O2- ions in a crystal lattice.
The formula unit shows the lowest whole-number ratio of Cu2+ to O2-, which is 1:1, CuO.
In contrast, individual H2 atoms do exist, and their subscripts show actual number of atoms, not a lowest-number ratio.

5. Ionic versus Molecular Compounds

6. Covalent Molecules
Combinations of atoms of the nonmetallic elements in groups 4A, 5A, 6A and 7A of the periodic table are likely to form covalent bonds.
Chemist Gilbert Lewis summarized this tendency in his formulation of the octet rule for covalent bonding:
Sharing of electrons occurs if the atoms involved acquire the electron configuration of noble gases.
The configurations often contain an octet (eight) valence electrons. [H2 is of course an exception to this rule.]

7. Covalent Bonding – Diatomic Gas Fluorine

8. In a water molecule, two hydrogen atoms form one single covalent bond each with one oxygen atom.
Note how the O atom ends up with eight electrons around it.
Covalent molecules will form if each atom will end up with 8 electrons around it (except H).
Each dot is one electron.
Each line is two electrons.

9. In an ammonia molecule, NH3, three H atoms form single covalent bonds with one N atom.
Note how the N has eight electrons around it.
When you consider the N and the three H’s, you can see the 2s and 2p orbitals are now full with eight electrons.

10. A methane molecule has four carbon-hydrogen bonds
A methane molecule has four carbon-hydrogen bonds. In each bond, C and H share the 1s e- from the hydrogen and a 2s or 2p e- from the carbon.
Normally C would start with s22s22p2 configuration, but by promoting one 2s e- to 2p, resulting in 1s22s12p3 it can create a stable octet with the four H atoms.

11. Carbon bonding – 4 covalent bonds
This is what Carbon “normally” looks like.
This is what it looks like when it “promotes” an e- to 2p so it can bond with four other atoms.
CH4 is much more stable than
CH2 so having four covalent bonds is better for Carbon.

12. Covalent Bonding
Different bonding models for methane, CH4. Models are NOT reality. Each has its own strengths and limitations.

13. Review: The Octet Rule and Covalent Compounds
Covalent compounds tend to form so that each atom, by sharing electrons, has an octet of electrons in its highest occupied energy level.
Covalent compounds involve atoms of nonmetals only.
The term “molecule” is used exclusively for covalent bonding

14. A single bond The Diatomic Fluorine Molecule2p
Each has seven valence electrons F 1s 2s 2p F F

15. Some double bonds: Carbon Dioxide
A carbon dioxide molecule has two C=O bonds
Note how C and the O’s each have 8 electrons now

16. An exception to the rule: The Diatomic Oxygen Molecule
Oxygen has six valence electrons.
You would think O2 would form a double bond by the looks of it, but experiments show its nonstandard and has two unpaired e- O 1s 2s 2p O 1s 2s 2p O O

17. A triple bond: The Diatomic Nitrogen Molecule1s 2s 2p
Each has five valence electrons N 1s 2s 2p N N

18. I Bring Clay For Our New House

19. Lewis Dot Structures
Lewis structures show how valence electrons are arranged among atoms in a molecule.
Lewis structures reflect the central idea that stability of a compound relates to noble gas electron configuration (atoms will react if they can arrange themselves to have 8 electrons around them).
Shared electron pairs are covalent bonds and can be represented by two dots (:) or by a single line ( - )

20. The HONC Rule Hydrogen (and Halogens) form one covalent bond
Oxygen (and sulfur) form two covalent bonds
One double bond, or two single bonds
Nitrogen (and phosphorus) form three covalent bonds
One triple bond, or three single bonds, or one double bond and a single bond
Carbon (and silicon) form four covalent bonds.
Two double bonds, or four single bonds, or a triple and a single, or a double and two singles

21. Completing a Lewis Structure: CH3Cl
Draw carbon as the central atom (it wants the most bonds, 4)
Add up available valence electrons:
C = 4, H = (3)(1), Cl = 7 Total = 14
Join peripheral atoms
to the central atom with electron pairs. H .. .. .. .. .. H C Cl .. ..
Complete octets on
atoms other than hydrogen with the remaining electrons H
Check: Final structure should have 14 e-

22. Coordinate Covalent Bonds
A covalent bond in which one atom contributes both bonding electrons is called a coordinate covalent bond.
This is signified by showing coordinate covalent bonds as arrows that point from the atom donating the pair of electrons to the atom receiving the bond.
Many polyatomic cations and anions contain both covalent and coordinate bonds. NH4+ is an example.

23. Carbon Monoxide – coordinate covalent bonding
In a coordinate covalent compound, one atom contributes both electrons of a bonding pair. In carbon monoxide, which atom contributes two electrons in one of the carbon-oxygen bonds?
:C O:
Triple bond – one of them is a coordinate bond

25. exceptions to octet rule

26. How to represent ions (covalent bonds within the ion)

27. Electron dot structure of the sulfite ion (SO32-) each O has 6 each S has 6 plus 2 extra = 26 e- Sulfur has to create one coordinate covalent bond to make this work.

29. Bond Dissociation Energies
Large amounts of heat are given off when hydrogen atoms combine to make H2, which implies that the product is more stable than the reactants.
If you try to break H2 apart, it will require a large amount of energy to do it.
The same thing is true for Carbon. A typical C-C single covalent bond has a bond dissociation energy of 347 kJ.
Since Carbon forms such strong C-C bonds, that explains why it’s compounds are so stable.
See table of bond dissociation energies on the next page. Note which one is weakest.

30. Bond Length and Bond Energy What trend do you notice?

31. Resonance
Occurs when more than one valid Lewis structure can be written for a particular molecule, such as ozone, below.
These are resonance structures.
The actual structure is an average or
a blend of the resonance structures.

32. Resonance in Benzene, C6H6
Each of these junctions represents where a Carbon is

33. Exceptions to the Octet Rule
NO2 (also known as smog in LA) has one unpaired electron, so it is an exception to the octet rule.

34. Oxygen: an exception to the octet rule
O O
The measured distance between oxygen atoms indicates that O2 does have some double bond character. This suggests that oxygen is a hybrid of the two structures shown on this page. O O

35. Other exceptions to the octet rule F – B – F F
BF3 is deficient by 2 e-.
Sometimes Phosphorus or Sulfur expand the octet to include 10 or 12 electrons.
Examples are PCl5 and SF6
P S

36. Diamagnetic vs. Paramagnetic Molecules
When e- spins are paired one up/one down as shown in these box diagrams, their moving electric charges create a magnetic fields that oppose each other, cancelling out the magnetic effects.
Normal molecules with all e- in pairs are weakly repelled by a magnetic field. These are called diamagnetic substances.
But when there is a lone unpaired e-, then the substance shows a strong attraction to an external magnetic field. These substances are paramagnetic. (But they are not permanent magnets like we see with Fe).

37. Section 16.3: Polar bonds and molecules
Covalent bonds involve sharing electrons between two atoms.
Sometimes the sharing is equal and the electron resides halfway in between the atoms, as in a diatomic gas like N2,Cl2, etc. This is called a nonpolar covalent bond.
This sharing isn’t always equal, because one atom may pull harder than the other atom, and then the electron will not be in the middle.
If the bonding electron is shared unequally, this is called a polar covalent bond (or just a polar bond).

38. Polar Bonds
The greater the electronegativity value, the greater the ability of an atom to attract electrons to itself. A high electronegativity atom is not “stealing” electrons as in the ionic case, but it is moving them in its direction.
Consider HCl. Hydrogen has an electronegativity of 2.1 and Chlorine has an electronegativity of These values are quite different, so the covalent bond in HCl is polar. The shared electron is pulled in the direction of Cl, because it is more electronegative. This can be represented as follows:
d+ d-
H – Cl H - Cl

39. Bond Polarity Electronegativity Differences and Bond Types
Electronegativity Difference Range
Most probable type of bond
Example
0.0 – 0.4
Nonpolar covalent
H – H (0.0)
0.4 – 1.0
Moderately polar covalent
d+ d-
H – Cl (0.9)
1.0 – 2.0
Very polar covalent
H – F (1.9)
≥ 2.0
Ionic
Na+ Cl- (2.1)

40. Sample Problem 16-4
Find which type of bond (nonpolar covalent, moderately polar covalent, very polar covalent or ionic) will form between each of the following?
N (3.0) and H (2.1) Δ = 0.9, moderately polar covalent
B) F (4.0) and F (4.0) Δ = 0.0, nonpolar covalent
C) Ca (1.0) and O (3.5) Δ = 2.5, ionic
D) Al (1.5) and Cl (3.0) Δ = 1.5, very polar covalent

42. Polar Molecules
The presence of a polar bond in a molecule makes the entire molecule polar. That means one end of the molecule is slightly negative and the other end slightly positive.
A molecule that has two poles is called a dipolar molecule or a dipole. If this kind of molecule is placed in an electric field, they orient themselves with respect to the positive and negative plates creating the field.

43. Polar Molecules
Some molecules are polar, but their polarities line up in such a way that they cancel.
Carbon dioxide is one such example
O = C = O

45. Intermolecular Attractions- van der Waals forces
The weakest intermolecular attraction is the van der Waals forces. These consist of two possible types, London dispersion forces and dipole interactions.
London dispersion forces, (weakest of all intermolecular interactions) are caused by the motion of electrons. The strength of dispersion forces increases as the number of electrons increases.
For halogens, which have more e- in their outer shell, the major attraction between them is dispersion forces.
These forces are weaker for F and Cl (gases at STP). They are stronger for Bromine, a liquid at STP, and even stronger for Iodine, a solid at STP.

46. Dipole interaction forces
The second type of van der Waals force is the dipole interaction, when polar molecules are attracted to one another.
The positive region of one molecule is attracted to the negative region of another.
HCl molecules

47. Hydrogen Bonds
A hydrogen bond is an attractive force where a hydrogen which is covalently bonded to a very electronegative atom (meaning the H has a slight + charge on it) is also weakly bonded to an unshared electron pair of another atom (pair has – charge to it).
This happens because when H bonds to O, F or N, the very polar bond leaves the H very electron deficient, with essentially an exposed nucleus with no electrons. The H nucleus is then attracted to a negatively charged unshared electron pair on another atom.
The resulting hydrogen bond is only about 5% of the strength of a regular covalent bond, but it is still the strongest of the intermolecular forces.
This is what causes water to be a liquid at room temperature

48. Hydrogen bonds in water
Hydrogen has valence e- that are not shielded from the nucleus by another layer of electrons.
Water has this type of interaction because the hydrogens have a slightly + charge and the oxygen has a slightly – charge.
This relatively strong interaction is called a hydrogen bond.

51. Hydrogen Bonding videos

52. Intermolecular Forces Summary
Weakest – London dispersion forces
Middle – dipole-dipole interactions
Strongest – hydrogen bonds
But all three are still much weaker than a covalent bond (max 5% of the strength of an average covalent bond)

53. Intermolecular Attractions and Molecular Properties
The physical properties of a compound depend on the type of bonding it has – ionic or covalent.
Here are some comparisons of physical properties
Characteristics of Ionic and Covalent Compounds
Characteristic
Ionic Compound
Covalent Compound
Representative unit
Formula unit
Molecule
Bond formation
Transfer of one or more electrons between atoms
Sharing of electron pairs between atoms
Type of elements
Metallic & nonmetallic
Nonmetallic
Physical state
Solid
Solid, liquid and gas
Melting point
High (usually > 300 C)
Low (usually < 300 C)
Solubility in water
Usually high
High to low
Electrical conductivity of aqueous solution
Good conductor
Poor conductor or nonconducting

54. Section 6.5 – Molecular Compounds
Binary molecular compounds are composed of two nonmetallic atoms.
Because atoms can combine in different ratios (for example CO and CO2 ) we use prefixes to help distinguish between compounds.
CO is carbon monoxide
CO2 is carbon dioxide
CCl4 is carbon tetrachloride
Note the –ide ending (similar to how an anion works, but these aren’t ionic compounds)
Prefix
Number
mono- 1
di- 2
tri- 3
tetra- 4
penta- 5
hexa- 6
hepta- 7
octa- 8
nona- 9
deca- 10

55. Section 6.5 – Molecular Compound Naming
To convert a name to a formula, write the correct symbols for the two elements, then add appropriate subscripts.
If there is just one of the first atom, you don’t need to write the mono-, it is assumed.
But for the second atom, if there is one, use mono-
Ex: tetraiodine nonoxide is I4O9
sulfur trioxide SO3
phosphorus pentafluoride PF5
Ex: N2O dinitrogen monoxide
PCl phosphorus trichloride
SF sulfur hexafluoride
H2O dihydrogen monoxide

56. Molecular Bonding - Acids
Here is a list of some of the most common acids which have covalent bonds and their names, which don’t always follow the standard naming convention.
HCl Hydrochloric acid
H2SO Sulfuric acid
HNO3 Nitric acid
CH3COOH Acetic acid (also written HC2H3O2)
H3PO4 Phosphoric acid
H2CO Carbonic acid
These are the most common ones and good to memorize