TYPES OF COMPOUNDS
Chemical Family Resemblances
Binary salts
Binary salts are made of a metal and a
nonmetal – only two different elements.
Examples: NaCl, MnO2
 Binary salts are named with the name of the
metal first, then the name of the nonmetal
with the “-ide” ending.
Example: K2O
potassium oxide

FORMULAS
 The
formula unit is the simplest ratio of
ions in the salt.
Ga2O3
2:3 ratio of gallium atoms to oxygen atoms
2 gallium atoms and 3 oxygen atoms make
one formula unit
formulas
 Electrons
and charge are conserved in a
formula unit.
– 2 gallium atoms have a total of 6 valence
electrons and no charge
– 3 oxygen atoms have a total of 18 valence
electrons and no charge
– so gallium oxide (Ga2O3) has 18+6=24
valence electrons and no charge
conservation
 Conservation
of electrons and charge in
gallium sulfide (Ga2S3)
conservation
oxidation numbers
 Oxidation
number of an ion is equal to
the charge on an ion after it gains or
loses electrons.
 All atoms gain or lose electrons to try to
attain a noble gas configuration (8
valence electrons)
 Noble gases have no oxidation numbers
oxidation numbers
 Metals
– lose all valence electrons,
positive (+) oxidation numbers
 Metals lose electrons so as to expose full
valence shell in next lower level
– Alkali metals and hydrogen are +1
– Alkaline earths are +2
– Aluminum and friends are +3
oxidation states
– Tin and lead are +2 or +4
– Transition metals vary
 Nonmetals
– gain electrons, negative (-)
oxidation numbers
 Enough electrons are gained to complete
the valence shell
– Oxygen is always –2, and sulfur is –2 unless
with oxygen
ternary salts
– Halogens are –1 unless with oxygen
– Nitrogen and phosphorus are –3 unless with
oxygen or halogens
 Ternary
salts are composed of more than
two elements
 Ternary salts contain polyatomic ions
– Polyatomic ions contain more than one atom
example: CO3-2 carbonate
polyatomic ions
 Polyatomic
anions have a (-) charge, and
polyatomic cations a (+) charge
 Polyatomic ions act as a unit – the
subscripts of the formula may not be
changed
 Names and formulas
– Most names end in “-ate” or “-ite”, which
means the ion contains oxygen
naming polyatomic ions
– Examples: sulfate (SO4-2), sulfite (SO3-2)
– The ending and prefix (if present) indicate
the relative number of oxygen atoms in the
formula.
perchlorate
chlorate
chlorite
hypochlorite
ClO4–
ClO3–
ClO2–
ClO–
polyatomic cations
 The
“-ium” ending means a positive ion
(hydronium, H3O+, and ammonium,
NH4+)
 Multiple ions are indicated by
parentheses and a subscript
– Example: magnesium hydroxide is
Mg(OH)2
– Ammonium sulfide: (NH4)2S
formulas with
polyatomic ions
 Formulas
are made the same way as the
binary salts, with the criss-cross method
+
Na 2
CO3
-2
+2 (OH-)
CaCa
2
Naming ternary salts
 Ternary
salts are named with the metal
name first, then the name of the
polyatomic ion
K3PO4
potassium phosphate
Transition metal salts
 Many
transition and “other” metals have
more than one oxidation number
 These numbers are found on some
periodic tables
 Metals to know: Fe (+2, +3), Cu (+1, +2),
Ag (+1), Zn (+2), Sn (+2, +4), Pb (+2, +4),
Bi (+3, +5)
transition metal salts
 Oxidation
number of transition metal is
indicated by a Roman numeral in
parentheses
 FeCl3 is iron (III) chloride
 Name these: CrO chromium (II) oxide
Cr2O3
chromium (III) oxide
CrO3
chromium (VI) oxide
transition metal salts
 The
Roman numeral is not needed if
there is only one oxidation state for the
metal (i.e. Zn, Ag, Sc)
 The Roman numeral is also used for
“other” metal salts like tin (II) fluoride
(SnF2, formerly used in toothpaste)
 Transition metal salts are often brightly
colored
hydrates
 Hydrates
are salts that have water
incorporated into the crystal structure
 The water is usually associated with the
cation
 The number of water molecules in the
crystal are specified in the formula
MgCl2. 6H2O
hydrates
 The
dot means they are not chemically
bonded
 Names of hydrates – “hydrate” plus a
prefix is added to the salt name
MgCl2. 6H2O
is magnesium chloride hexahydrate
 Prefix indicates the number of water
molecules
hydrate prefixes
mono = 1
tri = 3
penta = 5
hepta = 7
nona = 9
di = 2
tetra = 4
hexa = 6
octa = 8
deca = 10
Formation of hydrates
 Hydrates
can be formed when certain
salts are crystallized from water.
 Example – CuSO4. 5H2O {copper (II)
sulfate pentahydrate}
 Hygroscopic compounds become
hydrates by taking water from the air.
Formation of Hydrates
 Example
– sodium carbonate becomes
sodium carbonate decahydrate
(Na2CO3. 10H2O)
 Deliquescent compounds take enough
water from the air to form concentrated
solutions – examples: calcium chloride
(CaCl2), sodium hydroxide (NaOH)
More about polyatomic ions
 Bonding
– Polyatomic ions form ionic bonds with
metals
– They are held together with covalent bonds
 Formal
oxidation states can be assigned
to each atom in the ion using the
oxidation state rules
Oxidation state rules
 Oxygen
is always -2
 Hydrogen is always +1
 Sulfur is -2 unless with oxygen
 Nitrogen and phosphorus are -3 unless
with oxygen
 Halogens are -1 unless with oxygen
 When with oxygen oxidation states of
other atoms vary
Oxidation states of atoms
in polyatomic ions
 The
sum of all the oxidation states must
add to the charge of the ion
 Carbonate CO3-2
– oxygen – always (–2) charge
– total negative charge = (-6)
 overall
(+4)
charge is (-2), so carbon must be
Oxidation states of atoms
in polyatomic ions
 Try
these:
 Arsenate AsO4-3
O is -2, As is +5
 Cyanate NCO-1
O is -2, N (not next to O) is -3, so C must
be +4
Dot structures of
polyatomic ions
 All
valence electrons must be counted,
with extras added for a negative charge.
Carbonate CO3-2
Dot structures of
polyatomic ions
 Make
single bonds between all atoms
Dot structures of
polyatomic ions
 Pair
all electrons and make double
bonds where necessary to fulfill the octet
rule
Dot structures of
polyatomic ions
 Add
brackets and the charge
Dot structures of
polyatomic ions
 For
positive ions, leave out valence
electrons for positive charge
 Example: Ammonium (NH4+)
has 5 + 4 -1 = 8 valence electrons
Molecular substances
 Made
of molecules, which are loosely
held together
 Tend to be liquids, gases or low melting
solids
 Liquids can be purified by distillation
 Solids can be purified by
recrystallization
 Most are insulators
Molecular substances
 Molecular
elements
 Most nonmetals are molecular
 Diatomic gases – H2, N2, O2, F2, Cl2, Br2,
I2 (BrINClHOF)
 Bromine also exists as a liquid, and
iodine exists as a solid
allotropes
 Many
elements exist in more than one
molecular form
 oxygen (O2) and ozone (O3)
 carbon:
– charcoal, soot (random arrangement)
– graphite (flat sheets)
– diamond (three dimensional crystal lattice)
allotropes
fullerenes (hollow balls)
 linear acetylenic carbon
 -(-CC-CC-CC-)x phosphorus (P4):

– black (three dimensional, semiconductor)
– red (concatenated P4 molecules, used in matches)
– white (individual P4 molecules, unstable in air)

Different allotropes have different properties
Formulas and names of
small molecules
 Many
have common names (i.e. water,
ammonia)
 Systematic names use prefixes for each
element – same set of prefixes as for
hydrates
 P2O5 – diphosphorus pentoxide
 N2O – dinitrogen monoxide
Formulas and names of
small molecules
 “mono”
is not used for the first element
in a compound
 CO2 – carbon dioxide
 CO – carbon monoxide
 SO3 – sulfur trioxide
 CCl4 – carbon tetrachloride
Organic compounds
 Covalent
carbon containing compounds
– usually also contain H; may also
contain O, N, S, halogens, P
 Many names are derived from alkane
names
 Alkanes are hydrocarbons (containing
only C and H) with all single bonds
alkanes
Alkanes are named for the number of carbons
in a chain
 CH4 – methane
C2H6 – ethane
 C3H8 – propane
C4H10 – butane
 C5H12 – pentane
C6H14 – hexane
 C7H16 – heptane
C8H18 – octane
 C9H20 – nonane
C10H22 – decane
 General formula for an alkane: CnH2n+2

alkanes
 Carbon
always makes 4 bonds in organic
compounds, and hydrogen makes only 1
bond
 Oxygen makes 2 bonds (two lone pairs),
and nitrogen makes three bonds (one
lone pair)