1. Covalent Bonding

2. Remember… Atoms want to be stable so they will obey the octet rule.
The octet rule states that atoms gain, lose, and share valence electrons to achieve an outer octet of 8 valence electrons.
When forming ionic compounds, metals transfer valence electrons to nonmetals until all atoms obey the octet rule.

3. Atoms in covalent compounds share valence electrons until all atoms obey the octet rule. Atoms involved in covalent bonding are nonmetals and some metalloids.
Def: Covalent bond forms when atoms share valence electrons.
Def: A molecule forms when 2 or more atoms are covalently bonded.

4. Lewis Symbols

5. Bonding
Electrons that are paired will not share to form bonds. Single, unpaired electrons will share to form covalent bonds.
What kinds of bonds can these atoms make?

7. Bonding So, if carbon can make 4 bonds, let’s combine it with F.
Since carbon can make two sets of two bonds, lets combine it with oxygen.

8. Let’s look at some of the molecules you made.
These are called Lewis dot structures or dot structures.

9. Definitions
Def. Single covalent bond forms when 2 atoms share 2 electrons.
Def. Double covalent bond forms when 2 atoms share 4 electrons.
Def. Triple covalent bond forms when 2 atoms share 6 electrons.

10. Some atoms bond with themselves to form diatomic molecules.
Formula
Name H2
Hydrogen
Cl2
Chlorine N2
Nitrogen
Br2
Bromine O2
Oxygen I2
Iodine F2
Fluorine

11. Lewis Structures for Diatomic Molecules

12. Lengths & Strengths of Covalent Bonds
Single bond C C
Double bond C C
Triple bond C C

13. Bonds and Energy When bonds form or break, energy is involved.
When bonds form, the molecule that forms is more stable than the atoms it is made of. That means that, when the bond forms, energy is released. When energy is released, it is called an exothermic reaction

14. Bonds and Energy
When bonds break, energy is required to break the bond. That means that energy must be absorbed for the bond to break. This is an endothermic reaction when energy is absorbed.

15. Naming Binary Covalent/Molecular Compounds.
Covalent compounds are also called molecular compounds. They are named using prefixes to indicate how many of each element is included.

16. Prefixes
*Only mono- is NOT used with the name of the first element in the compound.
Number of Atoms
Prefix 1
mono- 5
penta- 8
octa- 2
di- 6
hexa- 9
nona- 3
tri- 7
hepta- 10
deca- 4
tetra-

17. Practice Naming Covalent Compounds
SO3 P4O10
B2F6 N2O5

18. Formula Writing of Covalent Compounds
Carbon tetrachloride
Diphosphorus trioxide
Dinitrogen monoxide
Chlorine trifluoride

19. Acids
Acids are compounds that, when dissolved in water, produce hydrogen ion (H+). There are two common types of acids – binary and oxyacids/oxoacids.

20. Binary Acids
Binary acids contain hydrogen and 1 type of nonmetal and follow this naming rule:
Hydro _______ ic acid
Formula
Name HF
Hydrofluoric acid
HCl
Hydrochloric acid
HBr
Hydrobromic acid HI
Hydroiodic acid
H2S
Hydrosulfuric acid

21. Oxyacids
Oxyacids contain hydrogen and a polyatomic acid The naming rule is: if the polyatomic ion ends in –ate, the acid name ends in –ic. If the polyatomic ion ends in –ite, the acid name ends in –ous. Add one hydrogen to the front of the acid for each NEGATIVE charge.

22. Polyatomic ion formula
Oxyacids
Polyatomic ion formula
Polyatomic ion name
Oxyacid Name
Oxyacid formula
NO3-
Nitrate
Nitric acid
HNO3
SO42-
H2CO3
BrO2-
Acetate
Phosphorous acid

23. Ways to Represent Covalent Compounds

24. Definitions
Def: Bonding pairs of electrons are electrons shared between atoms.
Def: Lone pairs of electrons are electrons located on an atoms and not shared with another atom.
Def: Resonance structures are different dot structures needed to correctly represent a molecule.
Def: Formal charge is the charge assigned to an atom in a molecule based on arrangement of valence electrons.

25. Creating Dot Structures
The Steps.
Count the total number of valence electrons available for bonding.
2. Arrange the atoms with one central atom. This will be the one with the lowest electronegativity.
3. Place a pair of electrons (bond) between each pair of atoms.
4. Use remaining electrons to complete octets on atoms.
5. Any electrons left are lone pairs. Pace them where needed to complete octets.