Mass #
Atomic #
Electric charge
Na
# of atoms

Also referred to as a “salt”

Formation involves a transfer of electrons

Usually made up of a metal and a non-metal

Are good conductors when they can be
melted or dissolved

Typically have extremely high melting points
e- jumps from Na to Cl
Electron acceptor (Cl)
meets
electron donor (Na)
Ions attract
to form a
neutral pair


Smallest building blocks are ions, NOT
MOLECULES
Large numbers of ions can attract to form
clusters and eventually crystals
Ion pair
Ion cluster
Crystal
lattice

Cations – positively charged ions
◦ Na+

Al3+
Anions – negatively charged ions
◦ Cl-

Ca2+
O2-
Polyatomic ions – ions made up of more than
one type of atom
◦ NO3-
SO4-2
PO4-3

The number of e- gained, lost or shared ub
compound formations
◦
◦
◦
◦
Alkali metals +1
Alkaline earth metals +2
Oxygen group -2
Halogens -1

K+ and N3◦ K3N

Ca2+ and N3◦ Ca3N2

Ba2+ and NO3◦ Ba(NO3)2

Criss-cross rule

Binary – made of 2 ions

Write cation first
Change anion ending to –ide

Na+ and Cl-

◦ Sodium chloride

H+ and F-
◦ Hydrogen fluoride

CaBr2
◦ Calcium bromide

Name the cation
Polyatomic ion name is unchanged

NaNO3

◦ Sodium nitrate

Zinc carbonate
◦ ZnCO3




Also called covalent compounds
A molecule is a neutral group of atoms that
are held together by covalent bonds
The valence e- are shared by the atoms
Covalent bonding usually occurs between 2
non-metals
◦ H2O, CO2, O2, NO

Use prefixes
1
2
3
4
5
6
7
8
9
10
monoditritetrapentahexaheptaoctanonadeca-

P4O10
Tetraphosphorous decoxide

N2O3
Dinitrogen trioxide

As2O5
Diarsenic pentoxide

OF2
Oxygen difluoride





7 diatomic molecules
No noble gases
Halogens and N, O, H
They are all gases (not
noble gases) except for
Br and I
“Honcl brif”
 H2
 O2
 N2
 Cl2
 Br2
 I2
 F2

H2SO4
Sulfuric Acid

HF
Hydrofluoric Acid

H3PO4
Phosphoric Acid

H2SO3
Sulfurous Acid

H2CO3
Carbonic Acid

HNO3
Nitric Acid
CaBr2

Calcium bromide

Chromium (III) acetate

Barium sulfate
BaSO4

Copper (I) sulfide
Cu2S

Sulfur hexafluoride
Cr(C2H3O2)3
SF6

Cr2(C2O4)3
Chromium (III) oxalate

Hg(CN)2
Mercury (II) cyanide

Cu(ClO4)2
Copper (II) perchlorate

ZnC4H4O6
Zinc tartrate


The mass of a compound
In order to calculate molar mass (also called
molecular weight) you add up the masses of
each element in the compound
◦ Be aware of subscript numbers that designate the
amount of atoms per element

You get the masses from the periodic table

**be careful when rounding the mass

NaCl
◦ Na = 23 g/mol
◦ Cl = 35.5 g/mol

H2O
58.5 g/mol
18 g/mol
◦ H = 1 g/mol (but there are 2) = 2 g/mol
◦ O = 16 g/mol

HNO3

Ba(NO3)2
63 g/mol
◦ H = 1 g/mol
◦ N = 14 g/mol
◦ O = 16 g/mol (but there are 3) = 48 g/mol
261.3 g/mol
◦ Ba = 137.3 g/mol
◦ N = 14 g/mol (but there are 2) = 28 g/mol
◦ O = 16 g/mol (but there are 6) = 96 g/mol

All metal atoms in a metallic solid contribute
their valence e- to form a “sea” of e◦ These e- move easily and freely because they are
not tied to a specific atom
 Delocalized electrons
◦ Metallic cation is formed
All empty space
is evenly
distributed v.e-


The attraction of a metallic cation for
delocalized electrons
This accounts for a lot of theproperties of
metals
◦
◦
◦
◦
Range of melting points
Malleability
Ductile
Durable
 Hard to remove metallic cation because of the strong
e- attraction
◦ Mobile e-
 Explains why they are good conductors

Find the difference in electronegativities of
the two elements
Non-polar
0.5 Polar
Pure
Covalent
-share e- evenly
-2 non metals
and/or metalloids
1.7
Polar
Covalent
-Share e- but
not evenly
-One element
holds e- more
Ionic
-Metal and
non-metal

Count total valence electrons available

Place electrons around atoms

Ensure each atom has an octet (8)
◦ Or a pair for H (2)

Draw the Lewis Structure for the molecule

Count the total number of . . .
◦ Bonded regions around the central atom
 DOUBLE and TRIPLE bonds count as ONE REGION
◦ Unshared e- pair
 Count as ONE REGION
Molecular
Structure
CH4
NH3
H2O
Lewis Dot
structure
H
H-C-H
H
H-N-H
H
H-O-H
electron pairs around central atom
total
shared
unshared
4
4
0 “tetrahedral”
4
3
1 “trigonal
pyramidal”
4
2
2 “bent”
Total no. of No. of No. of
Molecular
Molecule electron shared unshared
shape
pairs
pairs pairs

A molecule is polar if
◦ There is a polar bond
◦ It is ASSYMETRICAL (not symmetric)
(+)
O
(+)
H
(-)
H
(+)
Polar
(+)
H
H
C
H
(+)
Non-Polar
(+)
H

Symmetric (non-polar)
◦ Linear
◦ Tetrahedral
◦ Trigonal planar
 If all elements around the center atom are the same

Asymmetric (polar)
◦ Bent
◦ Trigonal pyramidal

Van der Waals forces (London Dispersion forces)
◦ Weak forces between non-polar molecules
◦ These forces determine volatility
 Doesn’t take much nrg to break apart (liquid gas)
 Most likely to be a gas
 Like playing red rover and only holding pinkies together

Dipole-Dipole
◦ Attraction between polar molecules
 Most likely to be a liquid
 Play red rover and hold hands

Hydrogen Bonding (H-Bonds)
◦ Between hydrogen (H) and a highly electronegative
element
 F, O, N
◦ Extreme case of dipole-dipole
◦ Strongest of the intermolecular forces
 Play red rover and link elbows
 Needs A LOT of nrg to break bonds

Carbon has a mass of 12 g

Oxygen has a mass of 16 g

H2O molecules has a mass of 18 g

How do these #’s relate to the atom or
compound?
◦ Atomic mass


Amedeo Avogadro (1776-1856)
1 mole = 6.0221415 x 1023
◦
◦
◦
◦
◦
◦
Particles
Molecules
Atoms
Ions
Formula units
Etc, etc

Determine the mass percentage of each
element in the compound.
mass _ of _ element
100
mass _ of _ compound

Gives the lowest whole # ratio of elements in
a compound.

The empirical formula for C6H12O6 is

The empirical formula for C2H6 is

* most basic ratio of elements in the
compound
CH2O
CH3