Types of Compounds
 Metal Binary Compounds – Metal and Non-Metal,
forms an Ionic Bond.
 Non-Metal Binary Compounds – two Non-Metals,
forms a Covalent Bond.
 Ternary Compounds – Contain Polyatomic ions. The
formula will have three or more elements in it.
Metal Binary Compounds
 Name the first element- a metal.
 Replace the ending on second element (non-metal)
with an “-ide” ending.
Examples:
NaCl
Sodium + Chlorine
= Sodium Chloride
MgS
Magnesium + Sulfur = Magnesium Sulfide
Naming Compounds with a Transition metal
 When some atoms can have more than one possible
charge, you name the charge on the atom.
 Copper is +1 or +2
Iron is +2 or +3
Cu +1 is Copper I
Fe +2 is Iron II
Cu +2 is Copper II
Fe +3 is Iron III
CuCl = Copper (I) Chloride
CuCl2 = Copper (II) Chloride
FeCl2 = Iron (II) Chloride
FeCl3 = Iron (III) Chloride
Non-Metal Binary Compounds
 Name the first element
 Replace the ending on the second element with “-ide”
 Use prefixes for the number of atoms in the formula.
CO2 is Carbon + Oxygen = Monocarbon Dioxide
N2O is Nitrogen + Oxygen = Dinitrogen Monoxide
H2O is Hydrogen + Oxygen = Dihydrogen Monoxide
Prefixes
1 atom = Mono2 atoms = Di3 atoms = Tri4 atoms = Tetra5 atoms = Pent-
6 atoms = Hex7 atoms = Hept8 atoms = Oct9 atoms = Non10 atoms = Deca-
Ternary Compounds
Compounds with Polyatomic Ions
 Name the first part of the compound- an element or
polyatomic ion.
 Name the second part of the compound- an element or
polyatomic ion. Examples:
MgSO4
Magnesium Sulfate
NH4OH
Ammonium Hydroxide
K3PO4
Potassium Phosphate
C.7.B write the chemical formulas of common
polyatomic ions, ionic compounds containing main
group or transition metals, covalent compounds, acids
Writing Formulas: Ionic Compounds
 Write chemical symbol for each part of the compound.
 Write the charge for the element.
Do the charges add together and equal zero?
 Yes, Stop this is the formula. The number of electrons
given away is the same as what is being taken by the
second atom.
 No, Cross the absolute value of the charge to the
opposite element as a subscript. Multiply the new
subscript by the charge and see if the new values will
add together and equal zero. If yes, Stop you have the
formula
 Potassium Bromide
K +1
Br -1
(+1) + (-1) = 0
 Magnesium Chloride
Mg +2
Cl -1
(+2)
+ (-1) = +1
Mg 1
Cl 2
1 (+2) + 2 (-1) = 0
Formula
Yes
KBr
No
Yes
MgCl2
Transition Elements
 Same rules as normal ionic compounds. The charge for
the transition metal will come from the name of the
compound.
 Iron III Chloride
 Fe +3 Cl -1
(+3) + (-1) = +2
No
 Fe1
Cl 3
1 (+3) + 3 (-1) = 0
Yes FeCl3
Polyatomic Ions
 The rules for polyatomic ions will be the same as ionic
compounds. Place the polyatomic ion in parenthesis.
 Keep the parenthesis at the end of the process if you
have a number greater than one outside of the
parenthesis. If you did not cross a number or if you
only crossed a one do not keep the parenthesis.
 Magnesium Sulfate
 Mg +2
(SO4) -2 Yes
MgSO4
 Iron III Phosphate
 Fe +3
(PO4) -3 Yes
FePO4
Sodium Hydroxide
Na +1
(OH) -1
Yes
NaOH
Do not keep the parenthesis because there is no number crossed.
Calcium Hydroxide
Ca +2
(OH) -1
Ca 1
(OH)2
1 (+2) + 2 (-1) = 0
Yes
Ca(OH)2
Keep the parenthesis because there is a number greater than one
outside the parenthesis
C.7.C construct electron dot formulas to
illustrate ionic and covalent bonds
 There are three main types of Chemical bonding.
Ionic, Covalent, and Metallic.
 Ionic Bonding occurs when there is a transfer of
electrons.
 Covalent Bonding occurs when atoms share
electrons.
 Metallic Bonding consist of the attraction of free
floating valance electrons for positively charged metal
ions.
Electronegativities are used to determine what type of
bond is formed when atoms come together in a
chemical reaction. To find the type of bond find the
difference in the electro negativities.
 If the difference is greater than 1.67 an ionic bond is
formed.
 If the difference is less than 1.67 a covalent bond is
formed.
 All atoms want to obtain eight electrons in the valence
energy level. To do so they will give, take, or share
electrons.
 NaCl Sodium Chloride

Sodium: (1.01)

Na: 1s22s22p63s1



Chlorine: (2.83)
Cl: 1s22s22p63s23p5
Sodium transfers the 3s1 to Chlorine to complete
the 3p energy level.
The electronegativity difference is 1.72
An ionic bond is formed.
Rules for Ionic Bonds
 The element with the fewest atoms goes in the center.
 The other atoms go around the central atom.
 Show the transfer of the electrons with a positive for
the atom that lost the electrons and a negative for the
atoms that gain the electrons.
 AsI3
Arsenic Triiodide
Arsenic (2.20)
Iodine (2.21)
As: 1s22s22p63s23p64s23d104p3
I: 1s22s22p63s23p64s23d104p65s24d105p5
The electronegativity difference is .01
A covalent bond is formed. The atoms share the
electrons.
Rules for showing Covalent Bonds
 The element with the fewest atoms goes in the center.
 The other elements go around the central atom.
 A bonding pair can only form where there is an
unpaired electron.
 Shared pairs or bonding pairs are shown with a dash.
One dash equals two electrons.
C.7.E predict molecular structure for molecules with linear,
trigonal planar, or tetrahedral electron pair geometries
using Valence Shell Electron Pair Repulsion (VSEPR) theory
Molecular Geometry
 The shape that a covalently bonded substance will take
is referred to as its Molecular Geometry.
 The shape is determined by the central atom, and the
number of shared and unshared electron pairs around
the atom.
 Electron pairs around the central atom will spread out
as far as possible to minimize the repulsive forces.
 This gives bond angles depending on the shape.
Linear molecule
Total number Number of
of electron
shared pairs
pairs.
2
2
Number of
unshared
pairs
0
Shape
Linear
Bond Angle
180 0
Trigonal planar molecule
Total number Number of
of electron
shared pairs
pairs.
3
3
Number of
unshared
pairs
0
Shape
Trigonal
Planar
Bond Angle
120 0
Tetrahedral molecule
Total number Number of
of electron
shared pairs
pairs.
4
4
Number of
unshared
pairs
0
Shape
Tetrahedral
Bond Angle
109.5 0
Trigonal Pyramidal
Total number Number of
of electron
shared pairs
pairs.
4
3
Number of
unshared
pairs
1
Shape
Trigonal
Pyramidal
Bond Angle
107.3 0
Bent molecule
Total number Number of
of electron
shared pairs
pairs.
4
2
Number of
unshared
pairs
2
Shape
Bent
Bond Angle
104.5 0
Linear
Tetrahedral
Trigonal Planar
Trigonal Pyramidal
Bent
C.7.A name, acids using International Union of
Pure and Applied Chemistry (IUPAC)
nomenclature rules
Naming Acids without Oxygen
 Acids without Oxygen are named with the prefix
“Hydro” and end in “ic”
 Examples:
 HCl is
Hydrochloric Acid
 HF
is Hydrofluoric Acid
 HBr is
Hydrobromic Acid
Naming Acids with Oxygen
 Some acids with oxygen have several forms and use
suffixes with “-ic” and “-ous” endings.
 The “-ic” or regular ending for an acid comes from the
polyatomic ion with the “-ate” ending. This gives the
regular count for the oxygen for this type of acid.
 Example: H2SO4
 SO4 is Sulfate so this acid is called Sulfuric Acid
 Once you know the “-ic” ending, count the number of
oxygens in the other forms to find the name for the
acid. (REMEMBER: The regular “-ic” form comes
from the polyatomic ion that ends with “-ate”)
 Two less oxygen Hypo ________ “-ous” Acid
 One less oxygen
________ “-ous” Acid
 Regular “ic” form
________ “-ic” Acid
 One more oxygen Per ________ “-ic” Acid
 The other names for the acids will come from the
count based from the “regular acid name”
 H2SO4 “-ate” ending so it is Sulfuric Acid
 H2SO3 “-ite” ending so it is Sulfurous Acid
 H2SO2 two less oxygen will have a prefix and “ous”ending. Hyposulfurous Acid.
 H2SO5 one more oxygen will have a prefix “Per” and
the regular “-ic” ending. Persulfuric Acid