Ions and Ionic Compounds
 Remember
an ion is an atom that has lost
or gained electrons
• Cations – positive – lost electrons
• Anions – negative – gained electrons
 Ionic
Compounds form when 2 or more
atoms are joined by the loss and gain of
electrons
• ALWAYS cation + anion
 Cation
is always first and anion is always
second.
 Remember…
• for the representative elements you can use the
location on the periodic table to determine how
many electrons will be lost or gained.
 Ex. Hydrogen is +1
 Ex. Fluorine is -1
 For
the transition elements you need
more information.
 Transition
Elements show their charge by
the roman numeral that follows their
name
• Ex. Copper (I) is Cu+1
 Transition
elements always lose electrons
so their charge is ALWAYS positive.
 The
name of cations (positive ions) is the
SAME as the name of the element.
 Ex. K+1 is potassium
• Transition metals do need the charge as a Roman
numeral following the name.
 So, if the element is in the middle – it needs a roman
numeral!
 Ex. Cu+1 is copper (I)
 The
name of anions (negative ions) is
different. Anions end in –ide.
• Ex. F-1 is fluoride (fluorine)
 Copper(II)
 Potassium
 Barium
+ Sulfur =
+ Nitrogen =
+ Sulfur =
 Lithium
+ Oxygen =
 Calcium
+ Nitrogen =
 Copper(II)
+ Iodine =
 Potassium
+ Sulfate =
 Magnesium
+ Hydroxide =
 Ammonium
+ Sulfite =
 Calcium
+ Phosphate =
 Aluminum
+ Nitrate =
 Potassium
+ Chromate =
 Rubidium
+ Perchlorate =
 Potassium
+ Permanganate =
 Lead(IV)
+ Hydroxide =
 Compound
must be neutral
• Charge must equal ZERO.
 The
Sum of the cation and anion charges
must be zero.
• Ex. (+1) + (-1) = 0
• Ex. (+2) + (-2) = 0
 We needed one cation and one anion to make the sum
ZERO.
 If
the charges do not add up to be zero,
then you must add additional cations or
anions so that the sum does equal zero.
• Ex. (+1) + (-2) ≠ 0
 so… you must add another (+1) ion.
• (+1) + (+1) + (-2) = 0
 We needed TWO (+1) cations and ONE (-2) anion to
make the sum ZERO.
 The
number of cations and anions you need
to make the sum ZERO is the ratio of cations
to anions in an ionic compound.
 This
ratio is called the FORMULA UNIT for
the ionic compound.
 The
ratio is represented in a formula as
subscripts for the cation and anion.
 Since
the charges add up to equal zero, NO
CHARGE should be written in the formula.
 Copper(II)
+ Sulfur =
• Cu+2 + S-2 = CuS
 Potassium
+ Nitrogen =
• K+1 + N-3 = K3N
 Barium
+ Sulfur =
• Ba+2 + S-2 = BaS
 Lithium
+ Oxygen =
 Calcium
+ Nitrogen =
 Copper(II)
+ Iodine =
 Potassium
+ Sulfate =
 Magnesium
+ Hydroxide =
 Ammonium
+ Sulfite =
 Calcium
+ Phosphate =
 Aluminum
+ Nitrate =
 Potassium
+ Chromate =
 Rubidium
+ Perchlorate =
 Potassium
+ Permanganate =
 Lead(IV)
+ Hydroxide =
 Name
cation first
• Remember Transition Metals should have a
roman numeral representing the charge in the
name.
 Name
the anion second.
• Elements that are anions will always end –ide.
• Do not change the ending of a polyatomic ion.
 Ex. Fluorine  Fluoride (ending change)
 Ex. Sulfate  Sulfate (no change)
 Cation
+ Anion
 Ions are joined by the transfer of
electrons
• Creates Electrostatic forces (attraction of
opposite charges) that hold the ions together
 Ionic
Compounds are composed of a
continuous arrangement of oppositely
charged ions.
• NOT a single separate unit
 Bonds
between atoms will be ionic when
there is a LARGE difference in
electronegativity between the atoms.
• > 1.7 ∆EN
 Metal
+ Nonmetal
• Opposite sides of the periodic table – large
differences in electronegativity
 Metal
+ Nonmetal (usually)
 Crystalline SOLIDS at room temp. (most)
• Crystal – repeating geometric pattern
 Brittle
 HIGH
melting points
 HIGH boiling points
• Some are so high it takes extreme conditions to
get them to change to gas
 Conduct
electrical currents when melted
or dissolved in water
Molecules and Covalent Bond
 Compounds
that form when atoms
SHARE electrons
• Forms a covalent bond
 NEVER
contain ions
• NEVER have charges
 Contain
only nonmetals
 Nonmetal + Nonmetal
 Can be solids, liquids, or
temperature
gases at room
• State is determined by the bond strength
(compounds with stronger bonds tend to be solids)
 LOW melting points
 LOW boiling points
 POOR conductors of
electricity under any
conditions
 Can be Polar or Nonpolar
• Depends on how the electrons are shared
• Polar compounds are better conductors
 Since
molecules do not have charges, their
names and formulas are determined
differently.
 Two elements may form covalent bonds in
more than one way creating DIFFERENT
chemical compounds
• Ex. H2O and H2O2
 Other examples:
• CH4 and C2H6
• C6H12O6 and C12H22O11
 Whenever
two elements form more than
one compound, the different masses of
one element that combine with the same
mass of the other element are in the ratio
of small whole numbers.
• Ex. H2O and H2O2
 contain the same mass of Hydrogen
 The mass of Oxygen is in a ratio of 1:2
 Since
two elements can combine in more
than one way, we must use information
besides their identities to determine
their formulas and names.
• Scientists perform quantitative analysis to
determine the mass ratios of elements to
determine their formulas
• We can use electron dot diagrams (Lewis
structures) to help us determine possible
arrangements
 Names
• Elements are named in the order they appear in the
formula.
• The final element’s ending is changed to –ide.
• Prefixes are added to each element to identify the
number of each atom in the formula
 EXCEPTION: Mono- is not used for the first element. If
there is only 1 of the first element, no prefix is added.
Prefixes
1. Mono2. Di3. Tri4. Tetra5. Penta-
6. Hexa7. Hepta8. Octa9. Nona10.Deca-
 The
prefixes in the name can be used to
write the formula for a given compound.
• Symbols for each element are written in the
order they appear.
• Subscripts are added based on the prefix of
each element.
 CO2
• Carbon dioxide
 CCl4
• Carbon tetrachloride
 N2O5
• Dinitrogen pentaoxide
 Carbon
Monoxide
• CO
 Sulfur
dioxide
• SO2
 Diphosphorus
• P2S3
trisulfide
 Acids
are molecular compounds that
behave more like ionic compounds
 ALWAYS contain Hydrogen
• Hydrogen is always the first element in the formula
 Can
contain polyatomic ions
 Since
they form compounds like ionics but
consist of only nonmetals, there are special
rules for naming and writing formulas for
acids.
 Formulas
for acids are determined the
same as ionic compounds.
• Sum of the charges must equal ZERO
• Hydrogen is always +1
• Element or polyatomic ion it combines with must
be NEGATIVE
 Example: Hydrogen
• H+1 + Cl-1 = HCl
 Example: Hydrogen
• H+1 + SO4-2 = H2SO4
+ Chlorine
+ Sulfate
 Two
different ways to name acids
• Hydrogen + Element
 Ex. HCl
• Hydrogen + Polyatomic Ion
 H2SO4
 Name begins with Hydro• Ex. Hydro Root of the element
• Ex. Hydrochlor Ends with –ic acid
• Ex. Hydrochloric acid
 Practice
• HF
• H2S

Name begins with polyatomic ion Root
• Sulf-

Name ends with –ic acid or –ous acid
• -ic acid is used when the polyatomic ion name ends with
–ide or –ate
 Sulfuric acid (H2SO4)
• -ous acid is used when the polyatomic ion name ends
with –ite
 Sulfurous acid (H2SO3)

Practice
•
•
•
•
HClO3
HCN
HClO2
H3PO3
Metallic Bonds
 Metal
+ Metal
 Consists of positive metal ions with a “sea
of electrons”
• Electrons are free floating – are not attached to
any one atom or ion
 HIGH
melting points
 HIGH boiling points
 HIGH conductivity
 Malleable, Ductile, High luster (shine)
• All properties are a result of the “free” electrons.
1.
2.
3.
4.
5.
6.
7.
8.
9.
10.
11.
12.
13.
14.
sodium hydroxide
sodium bromide
barium hydroxide
calcium oxide
lithium sulfide
carbon monoxide
sulfur dioxide
iron (II) sulfate
silver (I) chloride
copper (II) hydroxide
ammonium sulfide
nickel (II) fluoride
mercury (I) sulfate
iron (III) oxide
15.
16.
17.
18.
19.
20.
21.
22.
23.
24.
25.
magnesium phosphate
nickel (II) carbonate
diphosphorous
pentoxide
aluminum phosphate
nitrogen dioxide
phosphorus trichloride
dinitrogen pentoxide
germanium
tetrachloride
scandium bromide
bromine monoiodide
antimony pentasulfide