1. Covalent Bonding
Unit 6

2. Review
Octet Rule: The maximum number of electrons in the outer energy level is 8, atoms will form compounds to reach eight electrons in their outer energy level.

3. Covalent Bond
Two or more elements combine by sharing electrons, generally occurs with elements close together on the periodic table.
The positive nucleus of both atoms has almost equal attraction to the electrons being shared
The shared electrons spend most of their time between the 2 atoms
The force of attraction between both nuclei and the shared electrons hold the atoms together in a covalent bond
Can occur between atoms of the same element. (diatomic molecules)

4. Diatomic Molecules
When found alone in nature, these elements exist only as 2 atoms covalently bonded.
Memorize the 7 diatomic molecules:
Hydrogen-H2
Nitrogen-N2
Oxygen-O2
Fluorine-F2
Chlorine-Cl2
Bromine-Br2
Iodine-I2

5. Molecule vs. Compound
Molecule: formed when 2 or more atoms bond covalently.
Ex: H2O and CO2
Compound: formed when 2 or more atoms bond chemically. It can be ionic or covalent.
Ex: NaCl, MgO, H2O

6. Lewis Structures
A model that uses electron dot structures to show how electrons are arranged in molecules.
Pairs of dots or lines represent bonding pairs
Ex: H:H or H-H
Elements that share a pair of electrons form a single bond
Elements that share more than one pair of electrons form multiple bonds

7. Multiple Bonds
Double Bond: When two pairs of electrons (4 electrons) are shared between 2 atoms.
Ex: 2 oxygen atoms
Triple Bond: when 3 pairs of electrons (6 electrons) are shared between 2 atoms
Ex: 2 nitrogen atoms
Use 1 line to represent single bond
Use 2 lines to represent double bond
Use 3 lines to represent triple bond

8. Rules for Drawing Lewis Structures
Sum the valence electrons for all atoms.
Write the symbols for the atoms to show which atoms are attached to which, and connect them with a single bond.
Complete the octets around all the atoms bonded to the central atom.
Place any leftover electrons on the central atom.
If there are not enough electrons left to give the central atom an octet, try multiple bonds.

9. Naming Molecules Non-organic or Inorganic
uses prefixes that stand for the number of atoms present.
Prefixes: 1-mono, 2-di, 3-tri, 4-tetra, 5-penta, 6-hexa, 7-hepta, 8-octa, 9-nona, 10-deca
Rules:
First element written normal, use prefix if the number of atoms is greater than 1.
Second element, always add prefix indicating the number of atoms present and ending is –ide.
Ex: SO2= Sulfur Dioxide
N2O3=Dinitrogen trioxide

10. Organic Chemistry and Organic Molecules
Hydrocarbons: A compound composed of hydrogen and carbon.
Names consist of a prefix that indicates the number of carbon atoms followed by and ending that indicates the number of bonds between carbons.
Organic Prefixes:
1 carbon atom = meth 6 carbon atoms = hex
2 carbon atoms = eth 7 carbon atoms = hept
3 carbon atoms = prop 8 carbon atoms = oct
4 carbon atoms = but 9 carbon atoms = non
5 carbon atoms = pent 10 carbon atoms = dec

11. Alkanes Carbon to carbon bonds are single. End in –ane
Ex: C-C
End in –ane
General formula is CnH2n+2 where n = number of carbons.

12. Alkene One carbon to carbon bond is a double bond End in –ene
Ex: C=C
End in –ene
General formula is CnH2n

13. Alkyne One carbon to carbon bond is a triple bond End in –yne
Ex: C=C
End in –yne
General formula is CnH2n-2

14. Practice Problems
Propane
Heptene
Octyne

15. Characteristics of Ionic vs. Covalent Bonds

16. Hydrocarbons 2 main groups:
Aromatic: composed of a stable ring of carbon atoms with special characteristics.
Aliphatic: No ring of carbons. Includes alkanes, alkenes, and alkynes.

17. Aliphatic Characteristics
As alkane molecules grow larger and heavier, their boiling point and freezing point increase.
Small, light molecules exist as a gas at room temp.
As molecules get heavier it changes to liquid and gradually becomes a solid.

18. Saturated vs. Unsaturated
Saturated hydrocarbons: the compound only contains single bonds.
Unsaturated hydrocarbons: the compound has at least one double or triple bond.

19. Electronegativity
The tendency for an atom to attract electrons to itself when bonded.
Covalent Bonds: Have electronegativities less than 1.7
Ionic Bonds: Have electronegativities greater than 1.7

21. Molecular Polarity Not all covalent bonds are the same.
The pair of bonding electrons being shared are pulled between the nuclei of the atoms sharing them.
Look at electronegativity differences to determine type of covalent bond.
Types:
Nonpolar covalent
Polar covalent

22. Polar vs. Nonpolar Electronegativity Difference Type of Bond 0-0.4
Moderately polar
Very polar
1.7+
Ionic

23. Nonpolar When atoms are chemically similar like in diatomic molecules.
Bonding electrons are shared equally
Ex: H2 or O2
Bonds are symmetrically arranged (everything cancels out)
Ex: CO2 or CH4

24. Polar When atoms are chemically different.
Bonding electrons are shared unequally.
Dipole molecule: One of the molecules is positive while the other end is negative.
Ex: H2O

25. Isomers
Compounds with the same formula but different molecular structure or atomic arrangement.
Ex: C4H10 or Butane
Normal butane Isobutane
n-butane i-butane