Covalent Bonding
Covalent bonding deals with molecular compounds.
Molecule – a neutral group of atoms that are held together by covalent bonds.
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The valence electrons are shared by the atoms.
Covalent bonding usually occurs between 2 non-metals.
o Ex: H2O
CO2
O2
NO
Naming Molecular Compounds
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1
2
3
4
5
6
7
8
9
10
In naming molecular compounds we use prefixes in which we didn’t use for ionic
compounds
Prefix
MonoDiTriTetraPentaHexaHeptaOctaNonaDeca-
When used in naming, the end of the
prefix is dropped.
Ex: P4O10 = Phosphorous Decoxide
Not phosphorous decaoxide
Examples:
N2O3  dinitrogen trioxide
As2O5  diarsenic pentoxide
OF2  oxygen difluoride
** if there is only 1 of the atom that comes first in the formula we DO NOT use mono, we simply name
the element
Diatomic Molecules
H2
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O2
N2
Cl2
Br2
I2
F2
There are 7 diatomic molecules, they in clued the halogens, minus At, as well as N, O, H
No noble gases are diatomic because they are content and stable as is.
The 7 diatomic molecules are all gases that are not noble as well as Br and I
“hunckle brif”
Empirical Formula
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A formula that represents the lowest integral ratio of atoms of the elements in a compound
o C2H4
o C3H6
o C4H8
 Each of these have the same empirical formula: CH2
Think of empirical formula as the lowest common denominator of the chemical formula
o Ex: C2H6O2  CH3O
o NaC4H6O2  NaC4H6O2  this formula is already in its lowest form.
Naming Acids
Anion
--ide
--ite
--ate
Acid name
Hydro—ic
--ous
--ic
Example
HCl = hydrochloric acid
HNO2 = nitrous acid
HNO3 = nitric acid
***for –ite think tonsalitous…. Tonsal “ite” “ous”
*** for –ate think I “ate” something “ic”ky
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Only binary acids use “hydro”
o Binary = 2 components
Calculating Molar Mass (molecular weight, molecular mass, mass of compound)
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To calculate the molar mass of a compound, we must first find the mass of each element
involved in the formula
o Ex: H2O
 H = 1.00
 O = 16.00
Then we must account for how many of each element we have.
o 2 H  (1.00) x 2 = 2.00
o 1 O  (16.00) x 1 = 16.00
Then we add up the totals to get the Molecular Mass of the compound.
o 2.00 + 16.00 = 18.00
 H2O (water) has a molecular mass of 18.00
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Law of Conservation of Mass – in ordinary chemical reactions, matter can be changed in many
ways, but it can NOT be created or destroyed.
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Law of Definite Proportions – specific substances always contain elements in the same ratio by
mass.
o
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there are very definite proportions of each element in a particular compound
Law of Multiple Proportions – the ratio of masses of one element that combine with a constant
mass of another element can be exposed in ratios of small whole numbers
o
Ex: CO vs CO2  same elements but different proportions
o
H2O vs H2O2  the ratio of O to H is half in H2O of what it is in H2O2
Example:
In compounds of elements A and B, the ratio of masses of A & B are small whole
numbers.
Water 16g oxygen
H2O
2g hydrogen
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Hydrogen peroxide
H2O2
32g oxygen
2g hydrogen
8g oxygen
1g hydrogen
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16g oxygen
1g hydrogen
**Hydrogen peroxide MUST contain twice as much oxygen as water.