Chapter 4
Forces Between
Particles
3.1—NOBLE GAS CONFIGURATIONS
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Abbreviated electron configurations using the previous noble
gas to indicate all subshells are full to that point.
Valence Electrons
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Valence electrons—electrons in an atom that have the highest
n value.
The number of valence electrons is the same as the Roman
numeral group number for the main group or representative
elements.
Examples: Calcium, Ca, is in group II A. The number of
valence electrons is 2. Phosphorus, P, is in group V A. The
number of valence electrons is 5.
Lewis Dot Structures
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A representation of an atom or ion that shows the valence
electrons as dots arranged around the elemental symbol.
How to draw a Lewis dot structure:
• Write the elemental symbol
• Pretend the symbol has a box around it.
• Starting on one side of the box, place a dot on each side
correlating to the number of valence electrons.
X
Lewis Dot Structure Examples
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Determine the Lewis dot structures for the following elements.
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Potassium
Aluminum
Sulfur
Carbon
Chlorine
4.2—Ionic Bonding
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The octet rule—Main group atoms tend to gain, lose, or share
electrons to achieve an electron configuration identical to the
nearest noble gas.
Ion—formed when an atom gains or loses electron(s).
• An atom that loses electrons becomes a positively charged
ion called a cation.
• An atom that gains electrons becomes a negatively
charged ion called an anion.
Monatomic Ion Examples
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Write the symbol for the ions that would form from the
following atoms.
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Bromine
Oxygen
Magnesium
Nitrogen
Aluminum
General Rule for Ions of Main Group
Elements
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Main Group elements (IA-IIIA) will form ions having the same
positive charge as the group number.
Main Group elements (VA-VIIA) will form ions with a negative
charge equal to group number minus 8.
For example, strontium, Sr, belongs to group IIA and forms
Sr2+ ions and phosphorus, P, belongs to group VA and forms
P3- ions.
Ionic Bond Formation
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An ionic bond forms from the attraction of a cation and anion.
For main group elements an ionic bond occurs because of a
transfer of valence electron(s) from a metal atom to a
nonmetal atom. Both atoms are changed into ions with noble
gas configurations.
***Note: The ions are called isoelectronic with the noble
gases, which means they have the same electron
configuration as the noble gas.
4.3—Ionic Compounds
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Ionic compounds are compounds formed from ionic bonds.
A binary ionic compound is the result of a compound formed
from two monatomic ions. Note: there may be more than one
of each ion present.
Binary Ionic Compound Formulas
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The formula for the ionic compound shows the type and
number of each ion in the compound.
The cation is written first in the formula and the number of
each ion is a subscript to the right of the element symbol.
The number of each ion is determined by the fact that the
compound has a total charge of zero. In other words, the
same number of positives as negatives.
Binary Ionic Compound Formula Examples
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Determine the formula for the ionic compounds formed from the
following elements
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Sodium and fluorine
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Sodium and sulfur
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Aluminum and oxygen
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Magnesium and nitrogen
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Calcium and chlorine
4.4—Naming Binary Ionic Compounds
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Binary ionic compounds are named using the following
pattern:
name = metal name + stem of nonmetal name + -ide
The stem names and ionic symbols for some common
nonmetals are given in the following table:
Naming Binary Ionic Compound Examples
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Give the name for each of the following ionic compounds.
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K2O
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Mg3N2
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BeS
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AlBr3
Metals That Form More Than One Ion
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Some transition metals form more than one type of cation.
(e.g. Iron forms both Fe2+ and Fe3+ ions.)
Naming is the same as before except a Roman numeral, in
parentheses, follows the metal name to indicate the charge of
the metal ion.
For example, the compounds FeCl2 and FeCl3 are iron(II)
chloride and iron(III) chloride, respectively.
FeCl2
FeCl3
4.7—Polyatomic Ions
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Polyatomic ions—ions made from more than one type of
atom.
Ionic Compounds Containing Polyatomic
Ions
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When writing formulas for ionic compounds containing
polyatomic ions, the rules are essentially the same as those
for binary ionic compounds.
The formula of the cation is written first followed by the
formula of the anion and the number of each ion is a subscript
to the right of their symbols/formulas.
When more than one polyatomic ion is required in the formula,
parentheses are placed around the polyatomic ion before the
subscript is inserted.
Na3PO4
Mg3 PO4 2
NH4 3 PO4
Examples Using Polyatomic Ions
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Write the formula for the compounds formed from each of the
following.
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potassium ions and chlorate ions
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calcium ions and phosphate ions
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ammonium ions and sulfate ions
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aluminum ions and nitrate ions
Naming Using Polyatomic Ions
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The names of ionic compounds that contain a polyatomic ion
are obtained using the following pattern:
name = name of cation + name of anion
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Write the name for each of the following compounds.
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KClO3
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Ca3(PO4)2
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(NH4)2SO4
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Al(NO3)3
4.5—The Smallest Unit of Ionic Compounds
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An ionic compound’s crystal lattice is a rigid three-dimensional
arrangement of its ions.
Formulas for ionic compounds represent only the simplest
combining ratio of the ions in the compounds, not the precise
number of ions found in a crystal lattice.
Formula weight is the sum of the atomic weights of the atoms
shown in the formula of an ionic compound. This is similar to
molecular weight.
Example: sodium chloride, NaCl
FW = 22.99 amu + 35.45 amu = 58.44 amu
4.6—Covalent Bonding
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Covalent bond—a bond formed when atoms share valence
electrons. The shared electrons are counted in the octet of
each atom that shares them as illustrated below for fluorine,
F2.
The covalent bond may be represented by the shared pair of
electrons as dots or by a single line between the bonded
atoms.
Covalent bonds are generally formed from two nonmetals.
Covalent Bonds
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The sharing of electrons takes place when electron-containing
orbitals of atoms overlap.
An example of orbital overlap is shown in this example for the
formation of an H2 molecule:
Covalent Bonding Examples
Drawing Lewis Structures
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Step 1: Decide on the central atom—usually the one that
makes the most bonds (or the one written first in the
formula). C, N, P, and S are common central atoms.
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Group 7A atoms are usually terminal except when
bonded with O (in oxoacids) or other group 7A atoms.
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H is terminal because it only forms one bond.
Step 2: Determine the number of valence electrons for the
molecule or ion.
Step 3: Form single bonds (represented by lines) between
each pair of bonded atoms. Single bonds indicate one shared
pair of electrons.
Drawing Lewis Structures (Continued)
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Step 4: Use remaining electrons as lone pairs around each
terminal atom (except H) so that they are surrounded by 8
electrons. Place leftover electrons on the central atom.
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Central atoms in the 3rd Period and higher can have
more than 8 electrons (expanded octet).
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Boron is a central atom that can handle 6 electrons
instead of 8.
Step 5: If the central atom has less than 8 electrons, move
one of the lone pairs from a terminal atom and form a double
bond between that terminal atom and the central atom—same
idea if a triple bond is needed.
Lewis Structure Examples
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Draw a Lewis structure for each of the following.
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SO3
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SO42-
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C3H8
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CH3COOH
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CO2
4.10—More About Naming Compounds
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The pattern used to name binary covalent compounds is
similar to that used to name binary ionic compounds:
name = name of first element in the formula + stem of
second element + -ide
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A prefix is also included to indicate
the number of atoms of each element
in the molecule.
Note: The prefix mono- is not
used for the first element.
Naming Binary Covalent Compounds
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Name each of the following compounds:
SO2
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XeF6
N2O5
Write the formula for each of the following compounds:
phosphorus trichloride
silicon tetraiodide
diphosphorus pentaoxide
4.8—Shapes of Molecules and Polyatomic
Ions
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Most molecules and polyatomic ions have distinct threedimensional geometries and shapes.
The geometries and shapes of molecules or polyatomic ions
can be predicted using a theory called the valence-shell
electron-pair repulsion theory, or VSEPR theory (sometimes
pronounced "vesper" theory).
According to the VSEPR theory, electron pairs in the valence
shell of an atom will repel each other and get as far away from
each other as possible.
VSEPR Theory
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The first step in using VSEPR theory is to determine the
geometry of the molecule/ion.
• All areas of electron density (bonding and nonbonding
electron pairs) are considered when determining the
geometry.
• Double or triple bonds are considered one area of electron
density.
Electron Pair Arrangements
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According to the VSEPR theory, the arrangement of electron
pairs (also known as the electron geometry) around the
central atom (represented by E) depends on the number of
areas of electron density.
• Two areas form a linear geometry.
• Three areas produce a trigonal planar geometry.
• Four areas result in a tetrahedral geometry.
Molecular Shape
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Molecular shape focuses on how the bonded atoms arrange
themselves in space. This arrangement is based on the
geometry.
The shape is different from the geometry only when lone pairs
are located around the central atom.
For example, a molecule with three areas of electron density
(trigonal planar geometry), but one of the areas is a lone pair
results in a bent shape.
Summary of VSEPR Theory
Areas of
Electron
Density
Geometry
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Tetrahedral
3
2
Trigonal
Planar
Linear
Number of
Bond Angle
Lone Pairs
109.5°
120°
180°
Shape
0
Tetrahedral
1
Trigonal
Pyramidal
2
Bent
0
Trigonal
Planar
1
Bent
0
Linear
VSEPR Theory Examples
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Draw a Lewis structure and determine the geometry, shape, and
bond angles for each of the following.
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CO2
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NH3
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CO32-
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CHCl3
4.9—The Polarity of Covalent Molecules
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Sometimes there is an unequal sharing of electrons in a
covalent bond. This occurs because some atoms are more
“greedy” for electrons than other atoms.
The “greediness” of an atom is determined by the
electronegativity which is a measure of an atom’s attraction
for electrons.
Bond Polarity
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A polar covalent bond is the result of unequal sharing of
electrons.
**Note: Most bonds with a difference in EN of 0.4 or less can
be considered nonpolar.
In a polar covalent bond the more electronegative atom
acquires a partial negative charge (δ-) and the less
electronegative atom acquires a partial positive charge (δ+).
Molecular Polarity
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A molecule that has polar covalent bonds may be a polar or
nonpolar molecule depending on the distribution of the partial
charges.
A symmetrical distribution of charge in a molecule results in a
nonpolar molecule. A nonsymmetrical distribution of charge
results in a polar molecule.
Examples of Molecular Polarity
4.11—Other Interparticle Forces
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Ionic and covalent bonds represent two of the forces that
occur between ions or atoms which hold the compounds
together.
There are other forces that enable one molecule to be
attracted to another (intramolecular forces). These include:
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dipole forces,
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hydrogen bonding,
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dispersion forces.
Intramolecular Forces
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Intramolecular forces (IMFs) play a big roll in the boiling point
and melting point of substances. As well as solubility (one
substance dissolving in another).
Dipole forces are the attractive forces that exist between polar
covalent compounds (e.g. H2O and CO).
They are the attraction between the partial positive end of one
molecule and the partial negative end of another.
Dipole forces can exist between the same type of molecules
or between two different molecules.
Intramolecular Forces (Continued)
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Some polar covalent molecules (e.g. H2O) experience hydrogen
bonding, which is the result of a strong attractive dipole force
between molecules in which the hydrogen atoms covalently
bonded to very electronegative atoms (O, N, or F) are
attracted to an O, N, or F on another molecule.
Hydrogen bonding is the IMF that holds DNA in a double helix.
The last type of IMF (dispersion forces) occurs in all
compounds, but are the only force in nonpolar covalent
compounds. They work in much the same way as dipole
forces because they result from momentary nonsymmetric
electron distributions in molecules or atoms.
The Behavior of Selected Pure Substances
in Response to Heating