1. Unit 6: Covalent Bonding

2. Covalent Bonding
A chemical bond is the force that holds two atoms together and makes them function as a unit
Atoms form bonds to become most stable and obtain an octet
Covalent Bonding: Electrons are shared between two or more elements. Always between 2 non-metals
Metalloids can act as nonmetals
Never involves polyatomic ions
The bonding results from mutual attraction of the two nuclei for the shared electrons
Not all electrons are shared in a covalent bond. The unshared electrons are called lone pairs.
How is covalent bonding different than ionic bonding?

3. Properties of Covalent Molecules
Poor conductors
Low melting and boiling points than ionic
Because covalent bonds are usually weaker than ionic bonds
Soft
Most are liquid or gas phase
Typically more flammable than ionic compounds
Many are insoluble or partially soluble (don’t dissolve in water well)

4. Prefixes for Covalent Bonding
Number Indicated
Mono- 1
Di- 2
Tri- 3
Tetra- 4
Penta- 5
Hexa- 6
Hepta- 7
Octa- 8

5. Rules for Naming Covalent Molecules
The first element in the formula is named first, and the full element name is used.
The second element is named as though it were an anion (ending gets ide).
Prefixes are used to denote the numbers of atoms present. (AKA used to represent the subscripts)
The prefix mono- is never used for naming the first element. For example, CO is called carbon monoxide, NOT monocarbon monoxide.
Prefix(not mono)1st element space prefix2nd element- ide

6. Naming Covalent Compounds Examples
BF3
Rule 1: Name the first element, using the full element name: boron.
Rule 2: Name the second element as though it were an anion: Fluoride
Rule 3 and 4: Use prefixes to denote numbers of atoms. One boron atom: do not use mono- in first position. Three fluorine atoms: use prefix tri-.
The name is boron trifluoride.

7. Covalent Compounds Name the following covalent compounds: NO N2O5 CO2
SiO3
CCl4
IF5
PCl5
P4H6
Common Molecules to know:
NH3 is ammonia
CH4 is methane

8. Writing the Formulas for Covalent
Use the prefixes as subscripts (do not crisscross)
Do NOT simplify covalent molecules

9. Write the formulas for the following Covalent Compounds
Carbon Monoxide
Carbon tetrafluoride
Dinitrogen Trioxide
Selenium dioxide
Nitrogen Monoxide

10. Diatomic Molecules
Diatomic molecules: A molecule composed of two of the same atoms
You must memorize the following diatomic molecules
These are the only elements able to form diatomic molecules
These elements exist naturally as diatomic gases
Name
Formula
Hydrogen H2
Nitrogen N2
Oxygen O2
Fluorine F2
Chlorine
Cl2
Bromine
Br2
Iodine I2

11. Lewis Structures
A Lewis Structure is a representation of a molecule showing how valence electrons are arranged among the atoms in the molecule or ion.
In writing Lewis Structures, we ONLY include valance electrons
Electrons involved in bonding are called bonding pair. Electrons not involved in bonding are called lone pairs or unshared pairs
Keep in mind the octet rule when drawing Lewis structures
Exceptions to the octet rule:
Hydrogen and helium only need a duet (lithium and beryllium only need a duet as well, but they are present in ionic bonds)
Boron follows the octet, but most of the time it can be satisfied with only 6 valance electrons

12. Steps for Writing Lewis Structures
Calculate the sum of the valence electrons from all of the atoms. Do not worry about keeping track of which electrons come from which atoms. It is the total number of valence electrons that is important.
The least electronegative element is the central atom (fluorine is the most electronegative) (central atom can never be hydrogen or a halogen)
Use one pair of electrons to form a bond between each pair of bound atoms. For convenience, a line (instead of a pair of dots) is often used to indicate each pair of bonding electrons
Arrange the remaining electrons in pairs to satisfy the duet rule for hydrogen and the octet rule for all other elements. Always start with the central atom.
If you run out of electrons, take lone pairs from the central atom and turn them into double or triple bonds
If the molecule has a charge, put the structure in brackets and indicate the charge as a superscript

13. Lewis Structures Write the Lewis Structure for water:
Step 1: find the sum of the valence electrons for H2O:
= 8 valence electrons
Using a pair of electrons per bond, we draw in the two O-H bonds, using a line to indicate each pair of boding electrons:
H-O-H
We arrange the remaining electrons around the atom to achieve a noble gas electron configuration for each atoms. Remaining electrons = number of total valance electrons – number of electrons used in bonds (2 per bond)
Water has used 4 electrons for bonds so we have 4 remaining electrons (8-4=4)
How many lone pairs does water have?

14. Lewis Structure Draw the Lewis Structure for the following molecules:
CCl4
PH3

15. Lewis Structures with Multiple Bonds
Let’s do the Lewis Structure for Carbon Dioxide
Total valence electrons: = 16
Form Bonds: O-C-O
Remaining electrons: 16-4 = 12
Distribute remaining electrons:
Is this correct? (Did we use 16 electrons? Is every octet filled?)
No not all octet is filled
We need a double or triple bond
Correct Lewis Structures:
Have to draw in the incorrect lewis structure; have to draw in the correct lewis structures (2 double or triple bond)

16. Bonds
Single Bond- a covalent bond in which one pair of electrons is shared by two atoms
Double Bond- A covalent bond in which two pairs of electrons are shared by two atoms (stronger than a single bond)
Triple Bond- A covalent bond in which three pairs of electrons are shared by two atoms (stronger than single and double bonds)
No more than 3 pairs can be shared
Having multiple possible valid structures is referred to as having resonance

17. Examples
Draw Lewis structure for the following molecules; Do any of the following molecules have resonance? HF N2
NH4+
SO2
CF4 NO
CH2O
NO3-

18. Valence Shell Electron Pair Repulsion Theory
Planar triangular
Valence Shell Electron Pair Repulsion Theory
Tetrahedral
Trigonal pyramidal
Bent

19. VSEPR Theory
VSEPR (pronounced “vesper”) stands for Valence Shell Electron Pair Repulsion
VSEPR theory assumes that the molecular geometry or shape of a molecule is determined by the strong repulsion of electron pairs
Electrons around the central atom repel each other in order to have the maximum bond angle as possible
In other words in order to be as far apart as possible
Based on Electron Dot (Lewis structures)

20. Electron Domains
Electron domains are regions of high electron density (ex: bonds and lone pairs)
A bond is considered ONE electron domain regardless if the bond is single, double, or triple
A lone pair is considered an electron domain
Lone pairs repel electrons more than bonded pairs
In order to determine the VSEPR geometry or the shape of the molecule you must:
Draw the correct Lewis structure
Determine the number of electron domains around the central atom

21. VSEPR Overview You will need to know the following for each shape:
Shape name
What the shape looks like
Number of electron domains
Bond angle
You will need to know the following shapes:
Linear
Trigonal planar
Bent with 3 domains
Bent with 4 domains
Tetrahedral
Trigonal pyramidal

22. Linear Electron Domains = 2 Bonds on central atom =2
Unshared electron pairs on central atom = 0
Bond angles = 180 °
Examples: BeI2, CO2, HCN

23. Trigonal Planar Electron Domains = 3 Bonds on central atom =3
Unshared electron pairs on central atom = 0
Bond angles = 120 °
Examples: BF3, SO3 , NO3-1

24. Bent (or V shape) with 3 domains
Electron Domains = 3
Bonds on central atom = 2
Unshared electron pairs on central atom = 1
Bond angles = <120°
Lone pairs repel more than bonded pairs
Examples: SO2 , O3 , NO2-1

25. Tetrahedral Electron Domains = 4 Bonds on central atom = 4
Unshared electron pairs on central atom = 0
Bond angles = °
Examples: CH4 , SO42- , NH4+1

26. Trigonal Pyramidal Electron Domains = 4 Bonds on central atom = 3
Unshared electron pairs on central atom = 1
Bond angles = 107 °
Examples: NH3 , SO32-

27. Bent or V shaped with 4 domains
Electron Domains = 4
Bonds on central atom = 2
Unshared electron pairs on central atom = 2
Bond angles = °
Examples: H20, OF2

28. VSEPR Summary Shape Picture # of Electron Domains (#Bonds and #LP)
Bond Angle
Linear
Trigonal Planar
V shape with 3 domains
Tetrahedral
Trigonal Pyramidal
Bent with 4 domains

29. Bond Polarity
Electronegativity- The ability of an atom to attract an electron in a bond
Increases up and to the right
fluorine has the highest electronegativity (ignore noble gases)
Polar- having opposite ends
Nonpolar bonds- Share electrons equally
Polar bond- the electron is shared unequally
Creating a partially positive and a partially negative end
(lower case delta to indicate partial charge)
Why would the electron not be shared equally?
Different atoms have different electronegativity
It is a tug-of-war with the electron and the more electronegative element is winning

30. Bond Polarity
Polar bonds are indicated by a dipole moment (arrow with a tail pointing towards the more electronegative atom)
For example the bond between hydrogen and chlorine is polar; the arrow shown is the dipole moment
Lower case delta with a positive and negative indicates the partial charge
Polar Covalent bond: electronegativity difference ≥ 0.5 and < 1.7
Nonpolar Covalent bond: electronegativity difference < 0.5
Ionic bond: electronegativity difference ≥ 1.7
Write these as one single inequality:

31. Bond Polarity
Determine if the bond is polar, nonpolar, or ionic AND draw the dipole moment if there is one
Carbon and Oxygen
Hydrogen and Carbon
Hydrogen and Oxygen
Hydrogen and Fluorine
Oxygen and Fluorine
Electronegativity H
2.1 C
2.5 O
3.5 F
4.0

32. Molecular Polarity
Molecules are considered to be polar if they have an overall dipole (partially positive end and partially negative end)
Polar molecules must have one or more polar bonds
Polar molecules have a net dipole; meaning all of the dipole moments add together to one dipole
A molecule can have a polar bonds and be non- polar
This happens when the polar bonds cancel each other out by symmetry
Nonpolar molecules have no net dipole
Water is VERY POLAR

33. Polar vs. Non-Polar
No Net dipole moment
This is a nonpolar molecule

34. Molecular Polarity
Molecular polarity affects many properties of a compound
Solubility (ability to dissolve)
“Like dissolves like”: molecules of similar polarity and size will dissolve in each other
Polar dissolves polar
Nonpolar dissolves nonpolar
Polar will not dissolve nonpolar
Melting point
Boiling point
Surface tension