Review
0 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.
Covalent Bond
0 Two or more elements combine by sharing electrons,
generally occurs with elements close together on the
periodic table.
0 The positive nucleus of both atoms has almost equal
attraction to the electrons being shared
0 The shared electrons spend most of their time between
the 2 atoms
0 The force of attraction between both nuclei and the
shared electrons hold the atoms together in a covalent
bond
0 Can occur between atoms of the same element. (diatomic
molecules)
Diatomic Molecules
0 When found alone in nature, these elements exist only
as 2 atoms covalently bonded.
0 Memorize the 7 diatomic molecules:
0
0
0
0
0
0
0
Hydrogen-H2
Nitrogen-N2
Oxygen-O2
Fluorine-F2
Chlorine-Cl2
Bromine-Br2
Iodine-I2
Molecule vs. Compound
0 Molecule: formed when 2 or more atoms bond
covalently.
0 Ex: H2O and CO2
0 Compound: formed when 2 or more atoms bond
chemically. It can be ionic or covalent.
0 Ex: NaCl, MgO, H2O
Lewis Structures
0 A model that uses electron dot structures to show
how electrons are arranged in molecules.
0 Pairs of dots or lines represent bonding pairs
0 Ex: H:H or H-H
0 Elements that share a pair of electrons form a single
bond
0 Elements that share more than one pair of electrons
form multiple bonds
Multiple Bonds
0 Double Bond: When two pairs of electrons (4
electrons) are shared between 2 atoms.
0 Ex: 2 oxygen atoms
0 Triple Bond: when 3 pairs of electrons (6 electrons)
are shared between 2 atoms
0 Ex: 2 nitrogen atoms
0 Use 1 line to represent single bond
0 Use 2 lines to represent double bond
0 Use 3 lines to represent triple bond
Rules for Drawing Lewis
Structures
1. Sum the valence electrons for all atoms.
2. Write the symbols for the atoms to show which
atoms are attached to which, and connect them with
a single bond.
3. Complete the octets around all the atoms bonded to
the central atom.
4. Place any leftover electrons on the central atom.
5. If there are not enough electrons left to give the
central atom an octet, try multiple bonds.
Naming Molecules
Non-organic or Inorganic
0 uses prefixes that stand for the number of atoms present.
0 Prefixes: 1-mono, 2-di, 3-tri, 4-tetra, 5-penta, 6-hexa, 7-
hepta, 8-octa, 9-nona, 10-deca
0 Rules:
1. First element written normal, use prefix if the number of
atoms is greater than 1.
2. Second element, always add prefix indicating the number
of atoms present and ending is –ide.
Ex: SO2= Sulfur Dioxide
N2O3=Dinitrogen trioxide
Organic Chemistry and
Organic Molecules
0 Hydrocarbons: A compound composed of hydrogen and
carbon.
0 Names consist of a prefix that indicates the number of
carbon atoms followed by and ending that indicates the
number of bonds between carbons.
0 Organic Prefixes:
1 carbon atom = meth
2 carbon atoms = eth
3 carbon atoms = prop
4 carbon atoms = but
5 carbon atoms = pent
6 carbon atoms = hex
7 carbon atoms = hept
8 carbon atoms = oct
9 carbon atoms = non
10 carbon atoms = dec
Alkanes
0 Carbon to carbon bonds are single.
0 Ex: C-C
0 End in –ane
0 General formula is CnH2n+2 where n = number of
carbons.
Alkene
0 One carbon to carbon bond is a double bond
0 Ex: C=C
0 End in –ene
0 General formula is CnH2n
Alkyne
0 One carbon to carbon bond is a triple bond
0 Ex: C=C
0 End in –yne
0 General formula is CnH2n-2
Practice Problems
0 Propane
0 Heptene
0 Octyne
Characteristics of Ionic vs.
Covalent Bonds
Characteristic
Type of Particle
Bond Formed
Types of elements
Physical State
Melting Point
Solubility in water
Conductivity
Ionic
Ion
transfer emetal/nonmetal
solid
high
high
good
Covalent
Molecule
share enonmetal/nonmetal
solid,liq,gas
low
lower
poor to none
Hydrocarbons
0 2 main groups:
0 Aromatic: composed of a stable ring of carbon atoms
with special characteristics.
0 Aliphatic: No ring of carbons. Includes alkanes, alkenes,
and alkynes.
Aliphatic Characteristics
0 As alkane molecules grow larger and heavier, their
boiling point and freezing point increase.
0 Small, light molecules exist as a gas at room temp.
0 As molecules get heavier it changes to liquid and
gradually becomes a solid.
Saturated vs. Unsaturated
0 Saturated hydrocarbons: the compound only contains
single bonds.
0 Unsaturated hydrocarbons: the compound has at least
one double or triple bond.
Electronegativity
0 The tendency for an atom to attract electrons to itself
when bonded.
0 Covalent Bonds: Have electronegativities less than 1.7
0 Ionic Bonds: Have electronegativities greater than 1.7
Molecular Polarity
0 Not all covalent bonds are the same.
0 The pair of bonding electrons being shared are pulled
between the nuclei of the atoms sharing them.
0 Look at electronegativity differences to determine
type of covalent bond.
0 Types:
0 Nonpolar covalent
0 Polar covalent
Polar vs. Nonpolar
Electronegativity Difference
Type of Bond
0-0.4
Nonpolar
0.4-1.0
Moderately polar
1.0-1.7
Very polar
1.7+
Ionic
Nonpolar
0 When atoms are chemically similar like in diatomic
molecules.
0 Bonding electrons are shared equally
0 Ex: H2 or O2
0 Bonds are symmetrically arranged (everything
cancels out)
0 Ex: CO2 or CH4
Polar
0 When atoms are chemically different.
0 Bonding electrons are shared unequally.
0 Dipole molecule: One of the molecules is positive
while the other end is negative.
0 Ex: H2O
Isomers
0 Compounds with the same formula but different
molecular structure or atomic arrangement.
0 Ex: C4H10 or Butane
Normal butane
n-butane
Isobutane
i-butane