Chemistry Notes: Chapter 6
Chemical Bond—mutual attraction between the nuclei and the valence electrons of different
atoms that binds the atoms together. Two main types of bonds:
1. Ionic bonding—results from the electrical attraction between large numbers of cations and
anions formed from the transfer of electrons.
2. Covalent Bonding—results from the sharing of electron pairs between two atoms.
a. Nonpolar covalent bond—when the bonding electrons are shared equally resulting in a
balanced distribution of electric charge.
b. polar covalent bonds—bonds in which the bonded atoms have unequal attraction for
the shared electrons resulting in an uneven distribution of charge.
Electronegativity values are used to estimate what type of bond is formed. Take the difference
between the two electronegativity values, divided by the larger value, times 100. If the value is
less than 5% (less than .3 difference), the bond is nonpolar covalent. If the value is between 5%
and 50% (.3 to 1.7 difference), the bond is polar covalent. If the value is greater than 50%
(greater than 1.7 difference) , the bond is ionic.
Section 6-2
Molecule—neutral group of atoms that are held together by covalent bonds. A single molecule
of a compound is capable of existing on its own.
Molecular Compound—a compound whose simplest units are molecules. Molecular
compounds are made from nonmetals.
Molecular formula—shows the types and numbers of atoms in a single molecule of a molecular
compound. (H2O, C6H12O6)
Chemical Formula—Represents the composition of a compound. It indicates the relative
numbers of atoms of each kind in the compound using the symbols and subscripts. A
molecular formula is one kind of chemical formula.
Diatomic Molecule—Molecule containing only two atoms. (read page 165)
Bond Length—average distance between two bonded atoms.
Bond Energy-energy required to break a chemical bond and form neutral isolated atoms.
Octet Rule—Chemical compounds tend to form so that each atom, whether gaining, losing or
sharing electrons, has an octet of electrons in its highest energy level. All compounds want to
achieve the stability of the noble gases.
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There are some exceptions to the octet rule such as the H-H bond, which makes H2 stable
because Hydrogen only has one energy level and the first energy level is full with two electrons.
Electron dot notation: electron configuration in which only the valence electrons of an atom
are shown and they are indicated by dots placed around the element’s symbol.
They can also be used to represent molecules. See page 170-171.
A pair of dots between two symbols represents a pair of shared electrons.
An unshared (lone) pair is a pair of electrons that is not involved in the bonding and belongs
exclusively to one atom.
Lewis Structures—formulas in which atomic symbols represent nuclei and all inner-shell
electrons, dot-pairs or dashes between two atomic symbols represent electron pairs in covalent
bonds, and dots adjacent to only one atomic symbol represent unshared electrons.
A structural formula indicates the kind, number, arrangement, and bonds but not the unshared
electron pairs of the atoms in a molecule. F-F or H-Cl.
Single bond—bond produced by the sharing of only one pair of electrons between atoms. H-H
Double bond—bond produced by the sharing of two pairs of electrons between atoms. O = O.
Triple Bond—Bond produced by the sharing of three pairs of electrons between atoms such as
in the N2 triple bond.
Double or triple bonds are called multiple bonds. (Sample problem p. 174.)
Resonance: when a single Lewis Structure cannot correctly represent the bonding in molecules
or ions. Sometimes two or more structures are drawn and the bond is assumed to be the
average of these two structures. P 175.
Covalent-Network Bonding—When covalent compounds such as the carbon atoms in
diamonds do not contain individual molecules but instead are continuous, 3-D networks of
bonded atoms.
Section 6-3
Ionic compounds are positive and negative ions that are combined so that the numbers of
positives equal the number of negatives.
Most exist as crystalline solids.
The chemical formula for an ionic compound does not represent a molecule. Ionic
compounds never form molecules. Instead, it represents the ratio of ions present in any sample
size.
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A formula unit is the simplest collection of atoms from which an ionic compound’s formula
can be established. The ratio of ions in the formula unit is determined by the charges of the ions
combining. (read p. 177)
Characteristics of ionic bonding:
Ionic bonds form crystalline structures, not molecules.
Bond strength in ionic compounds is determined by the lattice energy released when one
mole of the compound is formed from the gaseous ions.
Ionic Compounds
Molecular Compounds
High Melting points
low melting points
Many are solids at room temp.
Many are gases at room temp.
Most are water-soluble
Not water-soluble.
When melted, good conductors of el. Poor conductors of electricity
Strong attractive forces between ions Weak attractive forces between
Hard but brittle.
molecules which is why they have
low melting points.
Polyatomic Ions—a charged group of covalently charged atoms. The charge results from an
excess or shortage of electrons. Polyatomic ions combine in the same manner as individual
atoms.
6-4
Metallic Bonding: Bonding that results from the attraction between metal atoms and the
surrounding sea of electrons. In a metallic bond, the outer electrons of the atoms area free to
roam throughout the entire metal. The electrons are delocalized which means they don’t belong
to any one atom. Instead, they move about the metal’s network of empty atomic orbitals.
Electrons form a seas of electrons. This gives metals some distinctive properties.
1. Good conductors of heat and electricity.
2. Can absorb a wide variety of light and then emit that light which makes them shiny.
3. Metals are malleable—the can be hammered into thin sheets.
4. Metals are ductile—can be drawn in to a thin wire.
Heat of Vaporization—amt. of heat needed to convert the metal to a gas. The higher the heat
of vaporization, the stronger the bonds that make up the metal.
6-5
The polarity of the bonds along with the geometric shape of the molecule determines the
polarity of the molecule or uneven distribution of charge. The predict the shape of a molecule,
one needs to use the VSEPR theory.
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VSEPR theory (valence shell electron pair repulsion) repulsion between the sets of valencelevel electrons surrounding an atom causes these sets to be oriented as far apart as possible.
(read page 183-186)
Hybridization—mixing of two or more atomic orbitals of similar energies on the same atom to
produce new orbitals of equal energies. The orbitals are called hybrid orbitals. Example sp3
orbitals in the Methane molecule. P 188
Intermolecular forces—forces of attraction between molecules (Table 6-7)
Dipole—created by equal but opposite charges that are separated by a short distance. Dipoles
go from the positive pole to the negative pole and are represented with an arrow.
Dipole-dipole forces hold polar molecules together. They are short-range forces acting only on
nearby molecules. These forces create high boiling points such as with water boiling at 100 
C. The polarity of the bonds and the geometric shape of the molecule determine molecular
polarity necessary to produce dipole-dipole bonds.
Water is polar because of its bent shape. Carbon dioxide is nonpolar because of its linear shape.
(p. 191)
hydrogen bonding—(a special kind of dipole) intermolecular force that results when an
hydrogen atom that is bonded to a highly electronegative atom of one molecule is attracted to
the unshared pair of electrons of an electronegative atom in a nearby molecule. Gives water its
high boiling pt. and other unique properties.
London Dispersion forces—intermolecular attractions resulting from the constant motion of
electrons and the creation of instantaneous dipoles. Act between all atoms and molecules but
they are the only intermolecular forces acting among noble gas atoms and nonpolar molecules.
They increase with increasing atomic mass, which causes increases in boiling pts.
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