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Chapter 4
Structure and
Properties of Ionic
and Covalent
Compounds
Denniston
Topping
Caret
4th Edition
4.1 Chemical Bonding
• Chemical Bond - the force of attraction
between any two atoms in a compound.
• Interactions involving valence electrons are
responsible for the chemical bond.
• Lewis symbol (Lewis structure) - a way to
represent atoms (and their bonds) using the
element symbol and valence electrons as
dots.
Lewis Symbols
Principal Types of Chemical Bonds:
Ionic and Covalent
• Ionic bond - a transfer of one or more
electrons from one atom to another.
• forms attractions due to the opposite
charges of the atoms.
• Covalent bond - attractive force due to
the sharing of electrons between atoms.
IONIC BONDING
Let’s examine the formation of NaCl
Na + Cl  NaCl
Sodium has a
low ionization energy
it readily loses this
electron .
When Sodium loses
the electron, it gains
the Ne configuration.
Na  Na+ + e-
Chlorine has a high
electron affinity.
When chlorine gains
an electron, it gains
the Ar configuration.

..
..


: Cl  e   : Cl :
..
 .. 
Essential Features of Ionic Bonding
• Atoms with low I.E. and low E.A. tend to form
positive ions.
• Atoms with high I.E. and high E.A. tend to form
negative ions.
• Ion formation takes place by electron transfer.
• The ions are held together by the electrostatic
force of the opposite charges.
• Reactions between metals, and metals and
nonmetals (representative) tend to be ionic.
COVALENT BONDING
Let’s look at the formation of H2
H + H  H2
• Each hydrogen has one electron in it’s
valance shell.
• If it were an ionic bond it would look like
this:
H   H   H  H :


• However, both hydrogen atoms have the same
tendency to gain or lose electrons.
• Both gain and loss will not occur.
• Instead, each atom gets a noble gas figuration
by sharing electrons.
H  H  H :H
Each Hydrogen
atom now has two
electrons around it
and has a He
configuration
The shared
electron
pair is a
Covalent Bond
Features of Covalent Bonds
• Covalent bonds tend to form between
atoms with similar tendency to gain or lose
electrons.
• The diatomic elements have totally
covalent bonds (totally equal sharing.)
H2, N2, O2, F2, Cl2, Br2, I2
Polar Covalent Bonding and
Electronegativity
1
The Polar Covalent Bond
• Polar covalent bonding - bonds made
up of unequally shared electron pairs.
• A truly covalent bond can only occur
when both atoms are identical.
• Electronegativity is used to determine if
a bond is polar and who gets the
electrons the most.
somewhat positively charged
somewhat negatively charged
..
..
H  F: H : F:


These two
electrons
are not shared
equally.
• The electrons spend
more time with fluorine.
• This sets up a polar
bond
• Electronegativity - a measure of the
ability of an atom to attract electrons in
a chemical bond.
electronegativity increases
electronegativity increases
• The greater the difference in electronegativity
between two atoms, the greater the polarity of a
bond.
• Which would be more polar, a H-F bond or a
H-Cl bond? The HF bond is more polar than
the HCl bond.
4.2 Naming Compounds and
Writing Formulas of Compounds
• Nomenclature - the assignment of a correct
and unambiguous name to each and every
chemical compound.
• We will learn two systems
– one for naming ionic compounds and
– one for naming covalent compounds.
Outline of Nomenclature Topics
I.
Ionic Compounds
A. Writing Formulas of Ionic Compounds from the
Identities of the Component Ions.
B. Writing Names of Ionic Compounds from the
Formula of the Compound
C. Writing Formulas of Ionic Compounds from the
Name of the Compound
IV. Covalent Compounds
A. Naming Covalent Compounds
B. Writing Formulas of Covalent Compounds
I. Ionic Compounds
• Metals and nonmetals usually react to form
ionic compounds.
• The metals are the cations and the nonmetals
are the anions.
• The cations and anions arrange themselves in
a regular three-dimensional repeating array
called a crystal lattice.
• Formula - the smallest whole number ratio of
ions in the crystal.
• Here is an
example: this is
the crystal lattice
of sodium
chloride (table
salt. Do not
write or draw
this.)
A. Writing Formulas of Ionic Compounds
• Determine the charge of the ions (usually can be
obtained from the group number.)
• Cations and anions must combine to give a formula
with a net charge of zero, it must have the same
number of positive charges as negative charges.
B. Writing Names of Ionic Compounds
from the Formula
Stock System:
1. Name cation followed by the name of anion.
2. Give anion the suffix -ide.
•
Examples:
•
NaCl is sodium chloride.
•
AlBr3 is aluminum bromide.
3. If the cation of an element has several ions of different
charges (as with Transition metals) use a Roman
numeral after the metal name.
4. Roman numeral gives the charge of the metal.
Examples:
•
FeCl3 is iron(III) chloride
•
FeCl2 is iron(II) chloride
•
CuO is copper(II) oxide
Common Nomenclature System
•
Use -ic to indicate the higher of the charges.
•
Use -ous to indicate the lower of the charges.
•
Examples:
•
•
FeCl2 is ferrous chloride, FeCl3 is ferric chloride
•
Cu2O is cuprous oxide, CuO is cupric oxide
Note that the common system requires you to know the
common charges and use the Latin names of the metals.
• Monatomic ions - ions consisting of a single atom.
• Examples
•
K+
potassium ion
•
Ba2+
barium ion
• Polyatomic ions - ions composed of 2 or more atoms
bonded together. Within the ion, the atoms are bonded
using covalent bonds. The ions will be bonded to other
ions with ionic bonds.
• Examples:
•
NH4+ ammonium ion
•
SO42- sulfate ion
Table 4.3 Common Polyatomic Cations and Anions
Ion
NH4+
NO3SO42OHCNPO43CO32HCO3CH3COO- (or C2H3O2-)
Name
ammonium
nitrate
sulfate
hydroxide
cyanide
phosphate
carbonate
bicarbonate
acetate
Tables 4.2 and 4.3 in your textbook gives a more complete list of
monatomic and polyatomic ions.
C. Writing Formulas of Ionic Compounds
from the Name
• Determine the charge on the ions.
• Write the formula so the compounds are
neutral.
• Example:
Barium chloride:
Barium is +2, Chloride is -1
Formula is BaCl2
II. Covalent Compounds
• Covalent compounds are usually formed
from nonmetals.
• Molecular Compounds - compounds
characterized by covalent bonding.
• not a part of a massive three dimensional
crystal structure.
A. Naming Covalent Compounds
1. The names of the elements are written in the order in
which they appear in the formula.
2. A prefix indicates the number of each kind of atom
•
mono-
1
hexa-
6
•
di-
2
hepta-
7
•
tri-
3
octa-
8
•
tetra-
4
nona-
9
•
penta-
5
deca-
10
3. If only one atom of a particular kind is present in the
molecule, the prefix mono- is usually omitted from the
first element.
Example: CO is carbon monoxide
4. The stem of the name of the last element is used
with the suffix –ide
5. The final vowel in a prefix is often dropped
before a vowel in the stem name.
B. Writing Formulas of Covalent Compounds
• Use the prefixes in the names to determine the
subscripts for the elements.
• Example:
• diphosphorus pentoxide
• P2 O 5
• Some common names are used.
4.3 Properties of Ionic and
Covalent Compounds
• Physical State
– Ionic compounds are solids at room temperature
– Covalent compounds are solids, liquids and gases
• Melting and Boiling Points
– melting point - the temperature at which a solid is
converted to a liquid
– boiling point - the temperature at which a liquid is
converted to a gas
4
• Melting and Boiling Points
– Ionic compounds have higher melting points and boiling
points than covalent compounds due to the large amount of
energy required to break the attractions between ions.
• Structure of Compounds in the Solid State
– Ionic compounds are crystalline
– Covalent compounds are crystalline or amorphous - have
no regular structure.
• Solutions of Ionic and Covalent Compounds
– Ionic compounds often dissolve in water, when they do they
dissociate (form positive and negative ions in solution) and
are Electrolytes -ions present in solution allow the solution
to conduct electricity; Covalent compounds usually do not
dissociate and do not conduct electricity – nonelectrolytes.
4.4 Drawing Lewis Structures on
Molecules and Polyatomic Ions
Lewis Structure Guidelines
1. Use chemical symbols for the various elements to
write the skeletal structure of the compound.
– the least electronegative atom will be placed in the central
position
2. Determine the total number of valence electrons
associated with each atom in the compound.
– for polyatomic cations, subtract one electron for every
positive charge;
– for polyatomic anions, add one electron for every negative
charge.
3. Connect the central atom to each of the surrounding
atoms using electron pairs. Then give each atom an
octet.
– Remember, hydrogen needs only two electrons
4. Count the number of electrons you have and compare
to the number you used.
• If they are the same, you are finished.
• If you used more electrons than you have add a bond for
every two too many you used. Then give every atom an
octet.
• If you used less electrons than you have….(see later
when discuss exceptions to the octet rule)
5. Check that all atoms have the octet rule satisfied and
that the total number of valance electrons are used.
Lewis Structure, Stability, Multiple Bonds,
and Bond Energies
• Single bond - one pair of electrons are shared between two
atoms
• Double bond - two pairs of electrons are shared between two
atoms
• Triple bond - three pairs of electrons are shared between two
atoms
• Bond energy - the amount of energy required to break a bond
holding two atoms together.
triple bond > double bond > single bond
• Bond length - the distance separating the nuclei of two adjacent
atoms.
single bond > double bond > triple bond
H : H or H - H
.. ..
O :: O or O  O
 
.. ..
N  N or N  N
Resonance - two or more Lewis structures that
contribute to the real structure.
Lewis Structures and Resonance
• Write the Lewis structure at CO32- on your paper.
• If you look at the people around you they probably
put the double bond in different places.
• Who is right? You all are.
..
:O:
:
:O:
::
..
:O:
:
..
..
..
..
..
: O : C :: O  : O : C : O :  : O :: C : O :




• Experimental evidence shows all bonds are the same
length. Shorter than single bond but longer than a
double bond.
• None of the three exist but the actual structure is an
average or hybrid of these three Lewis structures.
Lewis Structures and Exceptions to the Octet Rule
1. Incomplete Octet - less then eight electrons
around an atom other than H.
•
Let’s look at BF3
2. Odd Electron - if there is an odd number of
valence electrons it isn’t possible to give every
atom eight electrons.
•
Let’s look at NO
3. Expanded Octet - elements in 3rd period and
below may have 10 and 12 electrons around it.
Expanded octet is the most common exception.
•
Write the Lewis structure of SF6
Molecular Geometry: VSEPR Theory
• VSEPR theory - valance shell electron pair
repulsion theory. This is used to predict the
shape of the molecules.
• Electrons around the central atom (both
bonding and nonbonding pairs) arrange
themselves so they can be as far away from
each other as possible.
• We will do several examples in order to see
how the geometry is determined.
• These are examples, do not write
them or draw them.
• Let’s do BeH2 on the overhead.
• Let’s do BF3 on the overhead.
• Let’s do CH4 on the overhead
• Let’s do NH3 on the overhead
Trigonal Pyramidal
• Let’s do H2O on the overhead
Bent
The basic procedure to follow to determine the
shape:
1. Write the Lewis structure.
2. Count the number of bonded atoms and nonbonded electrons around the central atom.
• 2 legs - linear
• 3 legs - trigonal planer
• 4 legs - tetrahedron
3. Look at the atoms and name the shape.
• These would include: linear, trigonal planer, bent,
trigonal pyramid, tetrahedon.
Lewis Structures and Polarity
• Polar molecules are molecules that are polar or
behave as a dipole (two “poles”). One end is
positively charged the other is negatively
charged. Typically the result of polar bonds or
geometry.
• Nonpolar molecules do not have dipoles or
react in an electric field.
4.5 Properties Based on
Molecular Geometry
8
• Intramolecular forces - attractive forces within
molecules. (Chemical bonds)
• Intermolecular forces - attractive forces between
molecules.
• We will look at how intermolecular forces affect:
– I. Solubility.
– II. Boiling points and melting points.
• Solubility - the maximum amount of solute that dissolves
in a given amount of solvent at a specific temperature.
9
“Like dissolves like”
– polar molecules are most soluble in polar solvents
– nonpolar molecules are most soluble in nonpolar
solvents.
• Example: ammonia (NH3) in water.
• Example: water is polar and oil is nonpolar.
II. Boiling Points and Melting Points
• The greater the intermolecular force the
higher the melting point and boiling
point.
• Two factors to consider:
• Larger molecules have higher m.p. and b.p.
than smaller molecules.
• Polar molecules have higher m.p. and b.p.
than nonpolar molecules (of similar
molecular mass.)