Ionic & Covalent
Compounds
Bonding
Chemical bonds are forces that cause a
group of atoms to behave as a unit.
Bonds result from the tendency of a system
to seek its lowest possible energy.
Bond breaking always requires energy, and
bond formation always releases energy.
Types of Bonds
The type of bonding depends upon the
nature of the atoms that are combined.
A metal and a non-metal will form ionic bonds
when electrons are transferred from the metal to
the non-metal. The resulting attraction between
oppositely charged ions forms a stable crystal.
Types of Bonds
When metals bond with each other, the
valence electrons are shared by the atoms in the
entire crystal. The electrons are no longer
associated with a specific nucleus, and are free to
move throughout the sample.
Lewis Structures
Lewis Structures, also known as Lewis dot
diagrams, show how the valence electrons are
arranged among the atoms in the molecule.
For ionic compounds, it shows the end
result when the metal loses its electrons to the
non-metal.
Lewis Structures – Ionic Compounds
A dot (. ) is used to represent each valence
electron. Consider sodium chloride. The dot
diagram for the sodium atom is:
: :
Na .
The dot diagram for the chlorine atom is:
: Cl .
Lewis Structures – Ionic Compounds
: :
Note that electrons (dots) representing
valence electrons are written singly or in pairs on
either side of the atomic symbol, and above and
below the symbol.
Na . . Cl :
It doesn’t matter which side has the unpaired
electron.
Lewis Structures – Ionic Compounds
Sodium donates its single valence electron to
the chorine atom, and chlorine accepts the extra
electron.
: :
.
.
Na
Cl :
: :
After transfer, both ions have a noble gas
electron configuration.
Na+ : Cl:
Lattice Energy
When metals react with non-metals, a large
amount of energy is usually released. These
reactions are usually exothermic.
The electron transfer from the metal to the
non-metal usually requires energy, and is
endothermic.
The large release of energy comes from
forming an ionic crystal.
Lattice Energy
The crystal will have oppositely charged ions
in contact with each other and like charged ions
separated from each other. The alternating
charged ions form a crystal lattice. As the lattice
forms, large amounts of energy are released.
Lattice Energy
Lattice Energy
Lattice energy is defined as the energy
required to convert a mole of an ionic solid into
its constituent gaseous ions.
The greater the lattice energy, the stronger
the bonding, and the more stable the
compound.
Lattice Energy
Based on Coulomb’s Law, the lattice energy
increases with smaller ionic size.
Lattice Energy
Based on Coulomb’s Law, the lattice energy
increases significantly with greater ionic charge.
E = constant (q1)(q2)/r
Melting Points
As ionic crystals melt, the lattice is disrupted,
and the ions become free to move, and can
conduct electricity.
Ionic Bonding
The strength of
the attraction
between ions
increases significantly
with increasing ionic
charge, and results in
high melting points.
Naming Inorganic Compounds
1. Binary Compounds
Binary compounds contain only two elements.
The elements are either a metal with a nonmetal (ionic bonding), or two non-metals
(covalent bonding).
Naming Binary Compounds
a) Metal + Non-metal:
When metals react with non-metals, the metal
loses electrons and the non-metal gains
electrons. The resulting attraction between
oppositely charged ions creates ionic bonds.
Common Ionic Charges
The charges of ions of elements in groups
1A-7A (the main groups) are usually predictable.
Group 1A metals form +1 ions, group 2A
metals form +2 ions, etc.
The non-metals of group 5A have a -3
charge, those of group 6A have a -2 charge, and
the halogens form ions with a -1 charge.
Typical Ionic Charges
Naming Binary Compounds
For example, NaCl is called
sodium chloride
Where “chlor” is the root for the element
chlorine.
Naming Binary Compounds
Three common transition metals also have
only one ionic charge, and are also named the
same way.
They are: zinc ion (always +2), silver ion
(+1) and cadmium ion (+2)
ZnS is zinc sulfide, as “sulf” is the root for
sulfur.
Writing Formulas of Binary
Compounds
Compounds have no net charges, so the
formulas of ionic compounds must contain
equal numbers of positive and negative charges.
Magnesium bromide, made from magnesium
ion (Mg2+) and bromide ion (Br1-) has the
formula
MgBr2
Binary Compounds with Variable
Charge Metals
Most transition metals and the metals on the
lower right side of the periodic table can have
several ionic charges.
The properties of the ion vary greatly with
charge, so the charge must be specified in
naming the ion or its compounds.
Typical Ionic Charges
Binary Compounds with Variable
Charge Metals
Binary Compounds with Variable
Charge Metals
If an ion has variable charges, you must
specify the charge in naming the metal.
If an ion has only one charge, it is incorrect
to specify its charge.
Naming Fe2O3

Fe2O3 is an iron oxide, but we must specify the
charge of the iron ion.
We know each oxide has a -2 charge, so
three oxide ions have a total charge of -6.
The two iron ions therefore have a charge of
+6, with each iron having a charge of +3.
The name of the compound is iron(III) oxide.
Covalent Bonding
Covalent compounds
exist as discrete
molecules, whereas
ionic compounds
consist of an
aggregate of cations
and anions.
Covalent Bonds
When two (or more) non-metals form
bonds, electrons are shared. The result is a
covalent bond.
Covalent bonds form because the attraction
of electrons for the nuclei in the atoms is greater
than the electron-electron repulsion or the
nucleus-nucleus repulsion.
Types of Bonds
There is usually
an optimum bond
length or internuclear
distance where
attractions between
electrons and the
nuclei are optimized
and repulsions are
minimized.
Covalent Bonding
Bond Energy
Bond formation releases energy, and bond
breaking requires energy.
Molecules
Molecules are neutral combinations of two
or more atoms that are bonded together and act
as a unit.
Molecules may be elements (H2, O2, O3, F2),
or compounds containing atoms of different
elements bonded together. Molecules typically
contain elements that are non-metals.
Molecules
Scientists studying the nature of matter
focused much of their research in 1800s on the
composition of compounds.
Since molecules are much too small to
observe, they typically observed the reactions of
larger amounts of matter and used mass
measurements to develop their theories of
matter.
The Law of Definite Proportion

Joseph Proust (1754-1826) determined the
chemical composition of many compounds. He
found that a given compound always contains
the exact same proportion of elements by mass.
This is known as the law of definite proportion.
For example, all samples of water contain 88.8%
oxygen by mass, and 11.2% hydrogen by mass.
The Law of Multiple Proportions


This chemical law applies when two (or more)
elements can combine to form different
compounds.
Common examples are carbon monoxide and
carbon dioxide, or water and hydrogen peroxide.
John Dalton (1766-1844) conducted
experiments on these types of compounds, and
determined that there is a simple relationship
between the masses of one element relative to
the others.
The Law of Multiple Proportions

When two elements form a series of
compounds, the ratios of the masses of one
element that combine with a fixed mass of the
other element are always in a ratio of small
whole numbers.
The meaning of this law is difficult to
understand unless it is illustrated using a specific
series of compounds.
The Law of Multiple Proportions

Consider the compounds of water and hydrogen
peroxide. At this point in history, chemists
knew the compounds were different, and that
they both contain (or can be broken down into)
the elements hydrogen and oxygen. They did
not yet know the formulas for either compound,
nor was the concept of atoms fully developed.
The Law of Multiple Proportions

Analysis of 100 grams of the compounds produced the
following data:
Compound
Mass of
Mass of
oxygen/100g hydrogen/100g
of compound of compound
Grams of
oxygen/gram
of hydrogen
water
88.8 grams O
11.2 grams H
7.93 gO/gH
Hydrogen
peroxide
94.06 grams O
5.94 grams H
15.8 gO/gH
The Law of Multiple Proportions

The Law of Multiple Proportions is illustrated when the
numbers in the last column are compared.
Compound
Grams of
oxygen/gram
of hydrogen
water
7.93 gO/gH
Hydrogen
peroxide
15.8 gO/gH
15.8/7.93 = 2/1
The small whole
number ratio suggests
that there is twice as
much oxygen in
hydrogen peroxide as
there is in water.
The Law of Multiple Proportions

The key feature is that small whole numbers are
generated. The results support the hypothesis that
molecules consist of various combinations of atoms,
and that atoms are the smallest unit of matter. The
ratio doesn’t produce fractions, since there is no such
thing as a fraction of an atom.

For the example cited, we would propose that hydrogen
peroxide contains twice as many oxygen
atoms/hydrogen atoms than does water. We cannot,
however, determine the actual formula of either
compound.
Empirical Formulas
The studies of water and hydrogen peroxide
lead to empirical formulas. These are based on
experiment, and represent to simplest way of
expressing the ratio of atoms in a compound.
The early scientists analyzed new chemical
compounds to determine their composition and
chemical formulas. Modern analytical
laboratories still provide this service.
Molecular Formulas
Molecular formulas show the exact number
of each type of atom in a molecule. For
example, hydrogen peroxide has a molecular
formula of H2O2. Its empirical formula shows
that there is one hydrogen for every oxygen, so
it is OH or HO. Neither of these formulas
provides the structure or arrangement of the
bonds in the molecule.
Structural Formulas
Structural formulas provide the arrangement
of atoms in the molecule. The structural
formula for hydrogen peroxide is:
H-O-O-H
This formula shows the arrangement of the
atoms, but doesn’t show bond angles or the
shape of the molecule.
Naming Covalent Binary
Compounds
When two non-metals form a compound,
they share electrons, rather than transfer them.
The resulting bond is called a covalent bond.
The naming of these compounds is fairly
simple. The first element is named first, and
the second element is named as the root + ide.
Prefixes are used to indicate the number of
each atom present.
Naming Covalent Binary
Compounds
These prefixes are
used only for
compounds containing
two non-metals.
The prefix mono is
never used for the first
element in the
compound.
Naming Covalent Binary
Compounds

The prefix mono is never used for the first
element. CO2 is carbon dioxide.

If the prefix ends in an a or o, and the element
that follows begins with a vowel, the last letter
of the prefix is usually dropped. N2O5 is called
dintrogen pentoxide (and not pentaoxide).
Naming Covalent Binary
Compounds
Note that these prefixes are only used for
binary covalent compounds. It is incorrect to
use them for compounds containing a metal and
a non-metal.
Naming Covalent Binary
Compounds
There are some compounds of metalloids or
metals in very high oxidation states that are
sometimes named using this system.
Naming Covalent Binary
Compounds
For example, compounds such as SnCl4 or
PbCl4 are covalent in nature, and not ionic
solids. They may be called tin(IV) chloride or
tin etrachloride [or lead(IV) chloride or lead
tetrachloride.
Binary Compounds with
Hydrogen
With metals, hydrogen can form ionic
compounds in which the hydrogen has a -1 ionic
charge. These compounds are named like any
binary ionic compound.
NaH is sodium hydride
CaH2 is calcium hydride
Binary Compounds with
Hydrogen
With non-metals, the bonding is covalent.
Hydrogen never forms a positive ion in nature.
Many of the compounds containing
hydrogen have common names that do not
follow the usual nomenclature rules.
Binary Compounds with
Hydrogen
Examples include:
water
H2O
ammonia
NH3
phosphine
PH3
hydrogen sulfide H2S
Note that the order in which the
elements are written is also irregular.
Binary Compounds with
Hydrogen- Acids
Hydrogen also forms binary compounds that
act as acids in water. These compounds dissociate
in water to donate a proton to water.
HCl(g) + H2O(l)  H3O+(aq) + Cl–(aq)
hydrogen chloride
hydronium
Naming Binary Acids
The naming of the pure compound and its
aqueous acid solution differ.
HCl is a gas called hydrogen chloride.
HCl(aq) is an acid called hydrochoric acid.
Naming Binary Acids

Name the following acids:
H2S(aq) , HBr(aq)
Organic Nomenclature
Compounds containing carbon, hydrogen
and sometimes oxygen, nitrogen, sulfur and the
halogens, have a separate system of
nomenclature.
Unusual Ions
Mercury forms two ions, mercury(I) and
mercury(II). The mercury(I) ion is polyatomic,
and exists as two mercury(I) ions bonded
together. Its formula is Hg22+.
Oxygen in compounds usually exists as the
oxide ion, O2-. Oxygen also exists as the
peroxide ion, O22-, with each oxygen having a -1
charge.
Naming Polyatomic Ions
There are many ions, such as sulfate or
nitrate, that contain more than one element.
Many of these ions contain oxygen and a
non-metal.
These ions can be found in a group of acids
called the oxy acids (such as sulfuric acid, nitric
acid, etc.).
Polyatomic Ions
The bonding within these polyatomic ions
(such as nitrate, sulfate and phosphate) is
covalent. The ionic charge results from the loss
of one or more H ions to water, resulting in a
negative charge on the anion formed.
In water, the covalently bonded hydrogen is
donated to water, forming hydronium ions and
the corresponding anion.
Naming the Oxy Acids
The easiest way to learn the names of the
ions is to memorize a short list of oxy acid
names and their formulas.
The names of the ions are derived from the
names of the acids.
Keep in mind that the acids must be aqueous
solutions.
Common Oxy Acids
Acid
HNO3
Name
H2SO4
Sulfuric acid
HClO3
Chloric acid (or iodic or bromic acid)
H3PO4
Phosphoric acid
H2CO3
Carbonic acid
Nitric acid
Naming Complex Ions
Once the list of acids is learned, the names
of other acids and ions can be derived.
Removal of the hydrogens in the acid as H+
ions results in ions that end in ate.
HNO3 minus one H+ ion gives NO31-, the
nitrate ion.
The oxy acids that end in ic, produce ions
that end ate.
Naming Complex Ions
Sulfuric acid is H2SO4. Removing two H+
ions produces SO42-, the sulfate ion.
Keep in mind that the formula of the ions
must include the charge.
If only one of the H+ ions is removed from
sulfuric acid, HSO41- is produced. This is called
the hydrogen sulfate ion, also commonly known
as the bisulfate ion.
Naming Complex Ions
Carbonic acid, H2CO3, produces two ions:
HCO31-, the hydrogen carbonate or
bicarbonate ion
and
CO32-, the carbonate ion
Naming Complex Ions
Some of the oxy acids previously listed also
exist with one more oxygen in the formula.
HClO3, HBrO3 and HIO3 , in aqueous
solution are chloric, bromic and iodic acid
respectively.
Adding an oxygen to the formulas provides
the formulas for the per root ic acid.
HClO4 is perchloric acid. The ion, ClO41- is the
perchlorate ion.
Naming Complex Ions
Several of the oxy acids listed previously can
have one less oxygen atom in the formula.
These acids have names that end in ous, and ions
that end in ite.
HNO3 is nitric acid. HNO2(aq) is nitrous
acid. The ion NO21- is the nitrite ion.
Naming Complex Ions
Sulfuric acid, phosphoric acid, chloric,
bromic and iodic acids all can have one less
oxygen atom. The acids are sulfurous acid,
phosphorous acid, chlorous acid, bromous acid
and iodous acid.
The ions are called sulfite, phosphite,
chlorite, bromite and iodite ion.
Naming Complex Ions
The halogen oxy acids HClO3, HBrO3, and
HIO3 also exist with two less oxygen atoms in
the formula. The name of the resulting acid has
the name
hypo root ous acid.
HClO(aq) is hypochlorous acid, and ClO1- is
the hypochlorite ion.
Naming Complex Ions

If you memorize the list of acids ending in ic,
you can derive the names and formulas for many
other acids and ions.
Acid
HNO3
H2SO4
HClO3
H3PO4
H2CO3
Name
Nitric acid
Sulfuric acid
Chloric acid (or iodic or bromic acid)
Phosphoric acid
Carbonic acid
Naming Complex Ions



In naming the ions from the acids on the list,
remember that ic  ate.
If there is one additional oxygen atom, the acid
has the name per root ic, and the ion has the
name per root ate.
If there is one less oxygen atom, the acid has a
name ending in ous. The ions will have names
ending in ite. (ous ite)
Naming Complex Ions

If an acid has two less oxygen atoms than the
“ic” list, its name has the form hypo root ous.
The ion will have the name hypo root ite.
Other Common Formulas
CH3COOH
CH3COO1NH3
NH4+
OH1H3O+
MnO41CrO42Cr2O72-
Acetic acid
Acetate ion
Ammonia
Ammonium ion
Hydroxide ion
Hydronium ion
Permanganate ion
Chromate ion
Dichromate ion
Percent Composition
A chemical formula can be used to calculate
the percent composition of a compound.
Likewise, the percent composition can be used
to determine the empirical formula of a
compound. This is extremely useful when trying
to determine the formula of a new, or unknown
compound.
Chemical Composition
Usually, the compound is combusted in the
presence of oxygen. Any carbon in the compound
is collected as carbon dioxide (CO2), and any
hydrogen is collected as water (H2O).
Chemical Composition
Similar techniques exist to analyze for other
elements.
The formula obtained for the compound is
the simplest whole number ratio of the elements
in the compound, or the empirical formula. It may
differ from the actual formula. For example,
hydrogen peroxide is H2O2, but chemical
analysis will provide an empirical formula of
HO.
Percent Composition
To calculate the % composition of a known
compound, you determine the total mass of the
molecule, and the mass due to each of the
elements in the compound.
% by mass of element A =
total mass of A (100%)
molecular mass
Percent Composition
To calculate the % composition of a known
compound, you determine the total mass of the
molecule, and the mass due to each of the
elements in the compound.
% by mass of element A =
total mass of A (100%)
molecular mass
% Composition Problem

Problem: Calculate the percent composition of
ammonia.
Determining Formulas
It is generally more useful to obtain percent
composition data (usually from a laboratory),
and determine the empirical formula of a
compound. This will be the simplest whole
number ratio of the elements, and provides no
information about the structure of the
compound.
Determining Empirical Formulas
If given % composition:
1. Assume a quantity of 100 grams of the
compound.
2. Determine the number of moles of each
element in the compound by dividing the
grams of each element by the appropriate
atomic mass.
3. To simplify the formula into small whole
numbers, divide the moles of each element
by the smallest number of moles.
Determining Empirical Formulas
4. If necessary, multiply each number of moles
by a factor that produces whole number
subscripts.
5. If you know the approximate molar mass of
the compound, determine the molecular
formula.
% Composition Problem

Caffeine contains 49.48% carbon, 5.15%
hydrogen, 28.87% nitrogen, and 16.49% oxygen.
The compound has a molar mass of 194.2.
Determine the empirical and molecular formula
of caffeine.