Frederick A. Bettelheim
William H. Brown
Mary K. Campbell
Shawn O. Farrell
Omar J. Torres
www.cengage.com/chemistry/bettelheim
Chapter 3
Chemical Bonds
William H. Brown • Beloit College
Lewis Model of Bonding
In 1916, Gilbert N. Lewis pointed out that the lack of
chemical reactivity of the noble gases indicates a high
degree of stability of their electron configurations.
The Octet Rule
Octet rule: The tendency of group 1A-7A elements to
react in ways that achieve an electron configuration of
eight valence electrons.
• An atom that loses one or more electrons becomes a
positively charged ion called an cation.
• An atom that gains one or more electrons becomes a
negatively charged ion called a anion.
The Octet Rule
Example: In losing one electron, a sodium atom forms a
sodium ion, which has the same electron configuration
as neon.
Na (11 electrons): 1s2 2s2 2p6 3s1
Na+ (10 electrons): 1s2 2s2 2p6
The Octet Rule
Example: In gaining one electron, a chlorine atom forms
a chloride ion, which has the same electron
configuration as argon.
Chlorine atom (17 electrons): 1s2 2s2 2p6 3s2 3p5
Chloride ion (18 electrons): 1s2 2s2 2p6 3s2 3p6
The Octet Rule
The octet rule gives us a good way to understand why
Group 1A-7A elements form the ions they do; but it is not
perfect:
• Ions of Period 1 and 2 elements with charges greater
than +2 are unstable. For example, boron does not
lose its three valence electrons to become B3+, nor
does carbon lose its four valence electrons to become
C4+ .
• Ions of Period 1 and 2 elements with charges greater
than -2 are also unstable. For example, carbon does
not gain four valence electrons to become C4- .
• The octet rule does not apply to Group 1B-7B
(transition elements), most of which form ions with
two or more different positive charges.
Naming Cations
• Elements of Groups 1A, 2A, and 3A form only one type of
cation.
• The name of the cation is the name of the metal followed
by the word “ion”.
Naming Cations
For naming cations derived from transition and inner
transition elements, most of which form more than one
type of cation, there are two options:
• Use a Roman numeral enclosed in parentheses
following the name of the element to show the charge.
or
• Use the suffix -ous to show the lower positive charge
and the suffix -ic to show the higher positive charge.
Naming Cations
Naming Anions
For monatomic (containing only one atom) anions, add
“ide” to the stem name.
• Here are the monatomic anions we deal with most
often:
Naming Polyatomic Ions
Forming Chemical Bonds
According to the Lewis model:
• An atom may lose or gain enough electrons to acquire
a filled valence shell and become an ion. An ionic
bond is the result of the force of attraction between a
cation and an anion.
• An atom may share electrons with one or more other
atoms to acquire a filled valence shell. A covalent
bond is the result of the force of attraction between
two atoms that share one or more pairs of electrons.
Electronegativity
Electronegativity: A measure an atom’s attraction for the
electrons it shares in a chemical bond with another atom.
• On the Pauling scale, fluorine, the most
electronegative element is assigned a value of 4.0,
and all other elements are assigned values relative to
fluorine.
• Electronegativity is a periodic property and depends
on nuclear charge and the distance of the valence
electrons from the nucleus.
Electronegativity
Ionic Compounds
According to the Lewis model, an ionic bond is formed
by the transfer of one or more valence electrons from an
atom of lower electronegativity to an atom of higher
electronegativity.
• The more electronegative atom gains one or more
valence electrons and becomes an anion.
• The less electronegative atom loses one or more
valence electrons and becomes a cation.
• The compound formed by the combination of an anion
and a cation is called an ionic compound.
Forming an Ionic Compound
In forming sodium chloride, NaCl, one electron is
transferred from a sodium atom to a chlorine atom.
We use a single-headed curved arrow to
show this transfer of one electron.
Formulas of Ionic Compounds
The total number of positive charges must equal the total
number of negative charges.
• Lithium ion, Li+, and bromide ion, Br-, form LiBr.
• Barium ion, Ba2+, and iodide ion, I-, form BaI2.
• Aluminum ion, Al3+, and sulfide ion, S2-, form Al2S3.
• Sodium ion, Na+, and bicarbonate ion, HCO3-, form
NaHCO3.
• Potassium ion, K+, and phosphate ion, PO43-, form
K3PO4.
Naming Ionic Compounds
Binary ionic compounds
• The name of the metal from which the positive ion is
formed followed by the name of the negative ion;
subscripts are not directly specified in the name.
• AlCl3 is aluminum chloride.
• LiBr is lithium bromide.
• Ag2S is silver sulfide.
• MgO is magnesium oxide.
• KCl is potassium chloride.
Naming Ionic Compounds
Binary ionic compounds of metals that form two or more
different cations
• For systematic names, use a Roman numeral
enclosed in parentheses following the name to show
the charge on the metal ion; for common names, use
the appropriate -ous, -ic suffix.
• CuO is copper(II) oxide; cupric oxide.
• Cu2O is copper(I) oxide; cuprous oxide.
• FeO is iron(II) oxide; ferrous oxide.
• Fe2O3 is iron(III) oxide; ferric oxide.
Naming Ionic Compounds
Ionic compounds that contain polyatomic ions
• Name the positive ion first (using the appropriate
rules) followed by the name of the negative ion.
• NaNO3 is sodium nitrate.
• CaCO3 is calcium carbonate.
• NaH2PO4 is sodium dihydrogen phosphate.
• NH4OH is ammonium hydroxide.
• FeCO3 is iron(II) carbonate; ferrous carbonate.
• Fe2(CO3)3 is iron(III) carbonate; ferric carbonate.
• CuSO4 is copper(II) sulfate; cupric sulfate.
Forming a Covalent Bond
A covalent bond is formed by sharing one or more pairs
of electrons.
• The pair of electrons is shared by both atoms and, at
the same time, fills the valence shell of each atom.
• Example: in forming H2
Polarity of Covalent Bonds
Although all covalent bonds involve sharing of electron
pairs, they differ in the equality of the sharing:
• Nonpolar covalent bond: Electrons are shared equally.
• Polar covalent bond: Electron sharing is not equal.
• The equality of the sharing depends on the relative
electronegativities of the bonded atoms.
Polarity of Covalent Bonds
Examples:
Bond
H-Cl
O-H
N-H
Na-F
C-Mg
C-S
Difference in
Electronegativity Type of Bond
3.0 - 2.1 = 0.9
3.5 - 2.1 = 1.4
3.0 - 2.1 = 0.9
4.0 - 0.9 = 3.1
2.5 - 1.2 = 1.3
2.5 - 2.5 = 0.0
polar covalent
polar covalent
polar covalent
ionic
polar covalent
nonpolar covalent
Polarity of Covalent Bonds
In a polar covalent bond:
• the more electronegative atom attracts the shared
electrons more strongly and acquires a partial
negative charge; indicated by d- or the head of a
crossed arrow.
• the less electronegative atom attracts the shared
electrons less strongly and acquires a partial positive
charge; indicated by d+ or the tail of a crossed arrow.
Drawing a Lewis Structure
1. Determine the number of valence electrons in the
molecule. For a cation, subtract one electron for each
positive charge on it. For an anion, add one electron for
each negative charge on it.
2. Determine the connectivity of atoms.
3. Connect the atoms by single bonds.
4. Show bonding electrons as a single line; show
nonbonding electrons as a pair of Lewis dots.
5. In a single bond, atoms share one pair of electrons; in a
double bond, they share two pairs, and in a triple bond
they share three pairs.
Lewis Structures
Lewis Structures
Practice problems:
• Draw a Lewis structure for hydrogen peroxide, H2O2.
• Draw a Lewis structure for methanol, CH3OH.
• Draw a Lewis structure for acetic acid, CH3COOH.
Exceptions to the Octet Rule
Atoms of Period 2 elements use 2s and 2p orbitals for
bonding:
• These four orbitals can contain a maximum of 8
electrons; hence the octet rule.
Atoms of Period 3 elements have one 3s orbital, three 3p
orbitals, and five 3d orbitals:
• These nine orbitals can accommodate more than eight
electrons, by using 3d orbitals; period 3 atoms can
have more than eight electrons in their valence shells.
Exceptions to the Octet Rule
Phosphorus
Sulfur
Molecular Compounds
Molecular compound: A compound in which all bonds are
covalent.
Naming binary molecular compounds:
• The less electronegative element is named first (it is
generally written first in the formula).
• Prefixes “di-”, tri-”, etc. are used to show the number of
atoms of each element; the prefix “mono-” is omitted
when it refers to the first atom, and is rarely used with
the second atom. Exception: carbon monoxide
• NO is nitrogen oxide (common name: nitric oxide)
• SF2 is sulfur difluoride
• N2O is dinitrogen oxide (common name: laughing gas)
Resonance
For many molecules and ions, no single Lewis structure
provides a truly accurate representation.
Figure 3.3 Three Lewis strictures for the carbonateion.
Resonance
Linus Pauling - 1930s
• Many molecules and ions are best described by
writing two or more Lewis structures. These
molecules or ions are said to exhibit resonance.
• Individual Lewis structures are called contributing
structures.
• Double-headed (resonance) arrows are placed
between individual contributing structures.
• The molecule or ion is a hybrid of the various
contributing structures.
Resonance
Figure 3.4 The carbonate ion represented as a hybrid of
three equivalent contributing structures.
Resonance
All contributing structures must:
1. Have the same number of valence electrons.
2. Obey the rules of covalent bonding.
• No more than 2 electrons in the valence shell of H.
• No more than 8 electrons in the valence shell of a 2nd
period element.
• 3rd period elements, such as P and S, may have up to
12 electrons in their valence shells.
3. Differ only in distribution of valence electrons; the
position of all nuclei must be the same.
4. Have the same number of paired and unpaired electrons.
Curved Arrow
Curved arrow: A symbol used to show the redistribution
of valence electrons.
• the tail of the arrow identifies a pair of electrons
whose location is changing.
• the head of the arrow identifies the the new location of
the involved pair of electrons.
In using curved arrows, there are only two allowed types
of electron redistribution:
• from a bond to an adjacent atom.
• from an atom to an adjacent bond.
VSEPR Model
Valence-Shell Electron-Pair Repulsion (VSEPR)
• Valence electrons of an atom may be involved in
forming single, double, or triple bonds or they may be
unshared.
• Each arrangement of electrons creates a negatively
charged region of electron density around a nucleus.
• Because like charges repel each other, the various
regions of electron density around an atom spread so
that each is as far away as possible from the others.
VSEPR Model
Predict the shape of a
methane molecule, CH4.
• The Lewis structure shows
carbon surrounded by four
regions of electron density.
• According to VSEPR, the
four regions radiate from
carbon at angles of
109.5°, the electron pair
geometry is tetrahedral,
and the shape of the
molecule is tetrahedral.
• Figure 3.9 The shape of
methane molecule.
VSEPR Model
Figure 3.7 The shape of the
ammonia molecule, NH3.
• The Lewis structure
shows nitrogen
surrounded by four
regions of electron
density.
• According to VSEPR, the
four regions radiate from
nitrogen at angles of
109.5°, the electron pair
geometry is tetrahedral,
and the shape of the
molecule is pyramidal.
VSEPR Model
Figure 3.8 The shape of a
water molecule, H2O.
• The Lewis structure
shows oxygen
surrounded by four
regions of electron
density;
• The electron pair
geometry is tetrahedral,
and the shape of the
molecule is bent.
• The measured H-O-H
bond angle is 104.5°.
VSEPR Model
Figure 3.9 The shape of a formaldehyde molecule, CH2O.
• The Lewis structure shows carbon surrounded by
three regions of electron density.
• According to the VSEPR, the three regions radiate
from carbon at angles of 120°, the electron pair
geometry is trigonal, and the shape of the molecule is
trigonal planar.
VSEPR Model
Figure 3.9 The shape of an ethylene molecule, C2H4.
• The Lewis structure shows carbon surrounded by
three regions of electron density.
• According to VSEPR, the three regions radiate from
carbon at angles of 120°, each carbon is trigonal
planar, and the shape of the molecule is planar.
VSEPR Model
Figure 3.10 The shape of an acetylene molecule, C2H2
• The Lewis structure shows carbon surrounded by two
regions of electron density.
• According to VSEPR, the two regions radiate from
each carbon atom at an angle of 180° and the shape
of the molecule is linear.
• The measured H-C-C bond angle is 180°.
VSEPR Model
Molecular Polarity
A molecule will be polar if:
• it has polar bonds and
• its centers of partial positive and partial negative
charges do not coincide with the result that one
region of the molecule attains a net partial positive
charge and another attains a partial negative charge.
Carbon dioxide, CO2, has two polar C=O bonds but,
because of its geometry, CO2 is a nonpolar molecule.
Molecular Polarity
Ammonia, NH3, has three polar bonds and, because of its
geometry, is a polar molecule.
Molecular Polarity
Both dichloromethane, CH2Cl2, and formaldehyde, CH2O,
have polar bonds and are polar molecules. Acetylene,
C2H2, is a nonpolar molecule.
Molecular Polarity
Water, H2O, has two polar bonds and, because of its
geometry, is a polar molecule.
Chapter 3 Chemical Bonds
End
Chapter 3