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Unit 3 - Part 1:
Bonding Objective - to be able to understand and name the forces
that create chemical bonds.
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Bonding: Key Terms to Know
1. Chemical formula
2. Molecular formula
3. Bond Energy
4. Bond length
5. Octet Rule
6. Electron dot notation
7. Lewis Structures
8. Covalent Bonds
9. Ionic Bonds
10.Polyatomic ions
11.Resonance structures
12.Metallic bonding
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Molecular Compounds
• A molecule is a neutral
group of atoms that are
held together by covalent
bonds.
Caffeine
• A molecular formula
shows the number of
atoms of each element in
one molecule.
C6H12O6
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Ionic bonds -
created by
transferring
electrons
Covalent
bonds -
creating by
sharing
electrons
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Characteristics of the Covalent Bond • Bond energy is the
energy required to
break a chemical
bond and form neutral
isolated atoms.
• The distance between
two bonded atoms is
the bond length.
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The Octet Rule • Atoms can fill their outermost orbitals by sharing
electrons through covalent bonding. • Octet rule: Chemical compounds tend to form so
that each atom, by gaining, losing, or sharing electrons, has an octet of electrons (8) in its highest energy level.
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The Octet Rule, continuedExceptions to the Octet Rule
• Exceptions to the octet rule include those for atoms that
cannot fit eight electrons, and for those that can fit more
than eight electrons, into their outermost orbital.
• Hydrogen forms bonds in which it is surrounded by only
two electrons.
• Boron has just three valence electrons, so it tends to
form bonds in which it is surrounded by six electrons.
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Electron-Dot Notation
• Electron-dot notation shows the valence electrons of an atom, indicated by dots placed around the element’s symbol.
• ONLY the valence electrons are shown.
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Lewis Structures
• The pair of dots between the two symbols represents
the shared pair of a covalent bond.
• An unshared pair, also called a lone pair, is a pair of
electrons that is not involved in bonding and that
belongs exclusively to one atom.
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Lewis Structures
• The pair of dots representing a shared pair of
electrons in a covalent bond is often replaced by a
long dash.• A single covalent bond, or single bond, is a covalent
bond in which one pair of electrons is shared between two atoms.
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Lewis Structures, continuedSample Problem C
Draw the Lewis structure of iodomethane, CH3I.
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Sample Problem C Solution
1. Determine the type and number of atoms in the molecule.
The formula shows one carbon atom, one iodine atom, and three
hydrogen atoms.
2. Write the electron-dot notation for each type of atom in the
molecule.
Carbon has four valence electrons.
Iodine has seven valence electrons.
Hydrogen has one valence electron.
Lewis Structures
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Sample Problem C Solution, continued
3. Determine the total number of valence electrons
available in the atoms to be combined.
C 1 × 4e– = 4e–
I 1 × 7e– = 7e–
3H 3 × 1e– = 3e–
14e–
Lewis Structures, continued
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Sample Problem C Solution, continued
4. If carbon is present, it is the central atom. Otherwise, the least-
electronegative atom is central. Hydrogen, is never central.
Lewis Structures, continued
5. Make sure each atom is surrounded by eight electrons (except
hydrogen).
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Multiple Covalent Bonds
• A double covalent bond, or simply a double bond, is a covalent bond in which two pairs of electrons are shared between two atoms.
• A double bond is shown either by two side-by-side pairs of dots or by two parallel dashes.
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Multiple Covalent Bonds
• A triple covalent bond, or simply a triple bond, is a covalent bond in which three pairs of electrons are shared between two atoms.
• example 1—diatomic nitrogen:
• example 2—ethyne, C2H2:
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Multiple Covalent Bonds, continuedSample Problem D
Draw the Lewis structure for methanal, CH2O, which is also known as
formaldehyde.
1. Determine the type of atoms in the molecule.
2. Write the electron-dot notation for each atom.
3. Determine the total number of valence electrons in the whole
compound.
4. If carbon is present, it is in the middle. Otherwise, it is the least
electronegative atom.
5. Make sure each atom has 8 valence electrons.
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Ionic Compounds • The rocks and minerals that make up Earth’s crust
consist of positive and negative ions held together by
ionic bonding.• example: table salt, NaCl, consists of sodium and chloride ions
combined in a one-to-one ratio—Na+Cl–—so that each positive
charge is balanced by a negative charge.
• An ionic compound is composed of positive and
negative ions that are combined so that the numbers
of positive and negative charges are equal.
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Formation of Ionic Compounds
• The sodium atom has one valence electrons and the
chlorine atom has seven valence electrons.
• Atoms of alkali metals easily lose one electron to form
cations.
• Atoms of halogens easily gain one electron to form
anions.
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• Crystal lattice - an orderly arrangement found in ionic crystal.
Formation of Ionic Compounds, continued
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The Metallic-Bond Model, continued
• The chemical bonding that results from the attraction
between metal atoms and the surrounding sea of
electrons is called metallic bonding.
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Metallic Bonding• The unique characteristics of metallic bonding gives
metals their characteristic properties, listed below.
• electrical conductivity
• thermal conductivity
• malleability
• ductility
• shiny appearance
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Polyatomic Ions
• A charged group of
covalently bonded
atoms is known as a
polyatomic ion.
• Like other ions,
polyatomic ions have a
charge that results
from either a shortage
or excess of electrons.
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• Some examples of Lewis structures of polyatomic
ions are shown below.
Polyatomic Ions, continued
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Polyatomic Ions, continued • The charge of the ammonium ion is determined
as follows, continued:• When nitrogen and hydrogen atoms combine to
form an ammonium ion, one of their electrons is lost, giving the polyatomic ion a total negative charge of 10–.
• The total charge is therefore (11+) + (10–) = 1+.
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Resonance Structures - when molecules have
more than one Lewis Structure.
NO3-