Chemistry
Lecture 8: Basic concepts of chemical
bonding
Chapter 8: Basic concepts of chemical bonding
Chemistry – The Central Science – Twelfth Edition (Theodore L. Brown et al.,)
1. Lewis symbols and the
octet rule
Lewis symbols
• The electrons in the outermost shell are called valence electrons
• Outermost shell is sometimes called the valence shell
• Lewis structures (dot structures) are representations of the valence
electrons
• These structures will be used extensively to discuss chemical bonds
• To draw the Lewis structure of an atom:
• Write the symbol for the element
• Draw the dots to represent the number of valence electrons
Images: Cengage Learning
Lewis symbols
The outer electrons are called valence electrons -
these convey the most important properties of any atom
• Noble gases = Group 8A à very stable, unreactive to other atoms (high ionization
energies, low affinity for additional electrons, and general lack of chemical reactivity)
• Full valence shell
• Eight electrons in the valence shell = “octet”
Octet Rule - representative elements tend to form
compounds that will fill their valence shells
• Normally results in compounds in which each atom has 8 electrons in
its outermost shell
• Representative elements undergo reactions that leave them with eight
valence electrons so that they have a noble gas electron configuration
• Nonmetals can achieve an electron octet by sharing an appropriate
number of electrons à a covalent bond (more on that later).
Image: Cengage Learning
2. Ionic bonding
Ions
• Metals tend to form ionic compounds with
nonmetals
• Metals have fewer than 4 valence electrons
• Cannot satisfy the octet rule by forming covalent bonds
• Atoms with fewer or more than 4 valence electrons satisfy the
octet rule by losing or gaining electrons to form ions, respectively
• These valence electrons are exchanged in order to best fill the empty
spaces in the valence shells, allowing them to get filled (octet)
• By losing or gaining electrons, an atom is converted into a
charged particle called an ion
CATIONS
• A sodium atom does not
satisfy the octet rule
• Na loses one electron
from the 3rd shell to leave
a filled 2nd shell octet
• The Na ion now has a
positive (+) charge since
it lost an electron
Positively charged ion - Cation
Image: Cengage Learning
ANIONS
• A Chlorine atom adds
one electron to 3rd shell
• Leaves a complete
octet
• The Cl ion now has a
negative (-) charge by
gaining an electron
Negatively charged ion - Anion
Image: Cengage Learning
Cations and Anions
• Cations: Ions that form when atoms lose electrons
• Positively charged because the number of protons exceed the number
of electrons
• In naming ionic compounds, these are listed first and are called by the
element name
• Anions: Ions that form when atoms gain electrons
• Negatively charged because the number of electrons exceed the
number of protons
• In naming ionic compounds, these are named by replacing the suffix
with “-ide”
The ion charge does not depend on the
other element in a compound
• When two elements form an ionic compound, the number of electrons
gained or lost by each atom is independent
• Each atom is going to gain or lose as many electrons needed to satisfy its octet
• This may necessitate more than one atom to absorb the excess electrons
Opposites attract - Ionic compounds
• Ions of opposite charge are attracted to each other
• Cations and anions are bound together into an ionic compound
• The cation-anion attraction is called an ionic bond
• Not a sharing of electrons like a covalent bond
• Not a molecule because it is not held together covalently
• In ionic compounds, ions settle into a pattern to
efficiently fill space & maximize ionic bonding
• Ions in an ionic solid are held rigidly by
attraction to its neighbors by intermolecular
forces
• Ionic compounds dissolve in water if they are
more attracted to water than each otherImage: Cengage Learning
Ionic bonding
Ionic bonding
Exothermic process
Sodium: low ionization energy – readily gives up an electron
Chlorine: high electron affinity – readily gains an electron
Lattice energy
• Lattice energy is the energy required to completely separate
one mole of solid ionic compound into its gaseous ions.
Highly endothermic process
• The reverse (sodium ion and chloride ion coming together)
would be highly exothermic (ΔH = -788 kJ/mol)
Lattice Energy
Lattice energy
• The magnitude of the lattice energy depends on the charges of
the ions, their sizes, and their arrangement in the solid.
• In lecture 5, the potential energy of two interacting charged
particles is given by
in which, Q1, Q2 are the charges of the particles
d is the distance between their centers
And k is a constant 8.99 x 109 J.m/C2
• For a given arrangement of ions, the lattice energy increases as
the charges on the ions increase and as their radii decrease.
Lattice energy
Lattice energy
Lattice energy
Electronic configurations of ions of the s- and
p- block elements
In ionic compounds:
• Group 1A, 2A ,3A elements will form 1+, 2+, 3+ cations,
respectively.
• Group 5A, 6A, 7A elements will form 3-, 2-, and 1- anions,
respectively
Born-Haber cycle: Calculation of lattice
energy
Transition metal ions
• In forming ions, transition metals lose the valence-shell s
electrons first, then as many d electrons as required to reach
the charge of the ion. The octet rule, therefore, is limited when
transition metal ions are involved.
• For instance, in forming Fe2+ from Fe, which has the electron
configuration [Ar]3d64s2, the two 4s electrons are lost, leading to
an [Ar]3d6 configuration. Removal of an additional electron
gives Fe3+, whose electron configuration is [Ar]3d5.
• Silver, for example, has a [Kr]4d105s1 electron configuration. In
forming Ag+, the 5s electron is lost, leaving a completely filled
4d subshell.
3. Covalent bonding
Octet Rule - representative elements tend to form
compounds that will fill their valence shells
• Normally results in compounds in which each atom has 8 electrons in
its outermost shell
• Representative elements undergo reactions that leave them with eight
valence electrons so that they have a noble gas electron configuration
• Nonmetals can achieve an electron octet by sharing an appropriate
number of electrons à a covalent bond
Image: Cengage Learning
Octet Rule - representative elements tend to form
compounds that will fill their valence shells
• Nonmetals can achieve an electron octet by sharing an appropriate
number of electrons à a covalent bond
• Fluorine has 7 valence electrons and is found as F2. In this form,
both Fluorine elements share an electron
• Lewis structures can be used to represent these molecules.
• Diatomic molecules
F
+
F
F F
Images: Cengage Learning
Formation of Chemical Bonds
• A chemical bond occurs when two atoms are
attracted enough to each other to stay together.
• A covalent bond occurs when electrons are
shared between two atoms.
• A pair of shared electrons is known as a bonding
electron pair.
• Lone pairs or non-bonding pairs of electrons are
the electrons not involved in the covalent bond.
The bonding pair
of electrons
F F
Nonbinding pairs
of electrons
Covalent Bonds with Hydrogen atoms
• Hydrogen atoms need only one electron to fill the valence shell to
achieve the electron arrangement of Helium.
H
H
H H
Image: Fundamentals of General, Organic, and Biological Chemistry, 7th ed. Pearson Education
Covalent Bonds with Hydrogen atoms
Atoms that need more than one electron form
more than one covalent bond to reach an octet
• The number of empty spaces in the valence shell equals the number
of covalent bonds that an atom can make
• Oxygen has six valence electrons - two empty spaces in the valence
shell
• Forms two bonds to share enough electrons to fill its outer shell
H
O
H
H O H
Image: Cengage Learning
Atoms that need more than one electron form
more than one covalent bond to reach an octet
• The number of empty spaces in the valence shell equals the number
of covalent bonds that an atom can make
• Nitrogen has five valence electrons - three empty spaces in the
valence shell
• Forms three bonds to share enough electrons to fill its outer shell
Image: Cengage Learning
Atoms that need more than one electron form
more than one covalent bond to reach an octet
• Carbon has four valence
electrons - four empty spaces
in valence shell
• Forms four bonds to share
enough electrons to fill its outer
shell
Image: Cengage Learning
• For each of these molecules, we can draw in single lines to represent
a shared pair of electrons
Image: Fundamentals of General, Organic, and Biological Chemistry, 7th ed. Pearson Education
The number of empty spaces in the valence shell equals the
number of covalent bonds that an atom can make
• Elements in Groups 4A-7A
can satisfy the octet rule
• Elements with 5 or more
empty spaces rarely form
covalent bonds to fill their
empty valence shell
• Exceptions to the octet rule:
• Not relevant to this
course
Determining the central atom in a Lewis
Structure
• To draw the structure of a molecule, you need to know which
atoms are bonded to one another
• Rule: The atom which forms the most bonds is normally at the
center of the molecule
• Ex: Chloramine has the chemical
formula NH2Cl. This compound is
used as a bactericide in drinking
water. Identify the central atom
and draw the Lewis Structure.
H
Cl
Cl
N
H
H
N
H
H
N
Multiple Covalent Bonds - Double and Triple
Bonds
• The bonding in some molecules cannot be explained by the
sharing of only two electrons between atoms
• The only way these molecules’ atoms can have outer-shell
electron octets is by sharing more than two electrons, resulting in
the formation of multiple covalent bonds.
Multiple Covalent Bonds - Double Bonds
• Double bonds form when atoms share two pairs of electrons
• Four electrons total are shared
• Represented by a double line (each line represents an electron pair)
• Oxygen (O2) is a diatomic molecule
Lewis Structure of
O2 molecule
Image: Cengage Learning
Multiple Covalent Bonds - Triple Bonds
• Triple bonds form when atoms share three pairs of electrons
• Six electrons total are shared
• Represented by a triple line (each line represents an electron pair)
• Nitrogen (N2) is a diatomic molecule
Image: Cengage Learning
Example: Draw a Lewis Structure for CSCl2
C
• Carbon atom - needs four valence electrons for
octet
• Will form 4 bonds
Cl
Cl
• Chlorine atoms - each atom needs one valence
electron
S
• Sulfur atom - needs two valence electrons for
octet
• Each atom will form 1 bond
• Will form two bonds
Example: Draw a Lewis Structure for CSCl2
•
•
•
•
Cl C ClCl
C
S
S
Cl
Carbon can form the most covalent bonds - at center of the molecule
There are only three other atoms, so there must be a double bond
Chlorines can only make one bond each
Sulfur can form two bonds; double bond must be between O and C
Common Bonding Patterns for Electrically
Neutral Atoms
• C, N, and O are elements
most often present in
multiple bonds
• Common in organic
molecules
• Exceptions to octet rule
exist
• Nonmetals in 3rd period (Bo)
• You will not be asked about
the molecules that don’t
adhere to the octet rule
Image: Cengage Learning
4. Bond polarity and
electronegativity
Electronegativity and Polar Bonds
• Two identical atoms have equal attraction to the electrons in a
bond, so they are shared equally
• This results in the formation of a nonpolar covalent bond
• Neither of the atoms has an electrical charge
Image: Cengage Learning
Electronegativity and Polar Bonds
• When atoms are not identical, bonding
electrons may be shared unequally
• A polar covalent bond is one in which the
electrons are attracted more strongly to one
atom than by the other
• Fluorine attracts the electrons more than
hydrogen in HF, so fluorine is partially
negative, and hydrogen is partially positive.
• δ is used to denote a small electrical charge
Electronegativity =
the ability of an atom to attract electrons
• Fluorine is the most
electronegative element
• Assigned a value of 4.0
• Lower values assigned to
less electronegative
elements
• In general,
electronegativity:
• Decreases going down a
group
• Decreases from right to left
Image: Cengage Learning
Eletronegativity
• The electronegativity of an atom in a molecule is related to the
atom’s ionization energy and electron affinity, which are
properties of isolated atoms.
• An atom with a very negative electron affinity and a high
ionization energy both attracts electrons from other atoms and
resists having its electrons attracted away; it is highly
electronegative.
Which atom will be positively-charged and which will be
negatively-charged in each of the following bonds?
C
O
Oxygen is more
electronegative than Carbon,
so O is negatively charged and
C is positively charged
δ+
C O
δ-
Cl
Cl
There is no difference in
electronegativity since both
atoms are identical, so neither
is positive or negative.
Cl Cl
Bond polarity
We can use the difference in electronegativity between two atoms
to gauge the polarity of the bond the atoms form. Consider these
three fluorine-containing compounds:
Dipole moments
• A molecule such as HF, in
which the centres of positive
and negative charge do not
coincide, is a polar molecule.
• The quantitative measure of the
magnitude of a dipole is called
its dipole moment, denoted µ.
• If two equal and opposite
charges Q+ and Q- are
separated by a distance r,
Dipole moments
• Dipole moments are usually reported in debyes (D), a unit that
equals coulomb-meters (C-m). For molecules,we usually
measure charge in units of the electronic charge e, 1.6 x 10-19
C, and distance in angstroms.
Dipole moments
Differentiating ionic and covalent bonding
• If the difference of the electronegativity between the two elements is
greater than 1.7 then the bond is ionic. The difference with a polar
covalent bond is 0.5 to 1.7 and a nonpolar covalent bond is from 0 to
0.4.
• In general, as the oxidation state of a metal increases, so does the
degree of covalent bonding. When the oxidation state of the metal is
highly positive (roughly speaking, +4 or larger), we should expect
significant covalency in the bonds it forms with nonmetals.
• For example, manganese(II) oxide, MnO, is a green solid that melts
at 1842 oC and has the same crystal structure as NaCl.
Manganese(II) oxide, Mn2O7, is a green liquid that freezes at 5.9 oC,
which indicates that covalent rather than ionic bonding dominates.
5. Drawing Lewis
structures
Lewis structure
To draw Lewis structure, use the following procedure:
1. Sum the valance electrons from all atoms. For an anion, add an
electron for each negative charge. For a cation, subtract one
electron from the total for each positive charge.
2. Write the symbols for the atoms, show which atoms are attached
to which, and connect them with a single bond (a dash,
representing two electrons). The central atom is generally less
electronegative than the surrounding atoms.
3. Complete the octets around all the atoms bonded to the central
atoms.
4. Place any leftover electrons on the central atom (proceed anyway
if exceeds octets).
5. If there are not enough electrons to give the central atom an octet,
try multiple bonds (double or triple bonds).
Formal charge
• In some instances, we can draw more than one Lewis structure
and have all of them obey the octet rule. How do we determine
which one is the most important? => Determine the formal
charge.
• The formal charge of any atom in a molecule is the charge the
atom would have if all the atoms in the molecule had the same
electronegativity (that is, if each bonding electron pair in the
molecule were shared equally between its two atoms).
Rules governing formal charges
• Calculate
Formal
Charge
Rules
Governing
Formal Charge
• Add up the lone pair electrons on the atom, and half of the shared
electrons around the atom (those in bonds)
• Subtract the assigned electrons from the element’s valence
electrons
• The sum of the formal charges on all atoms in a given molecule or
ion must equal the overall charge on the species
• Molecules = zero
• Ions = ionic charge
• Which is the preferred Lewis Structure?
• Formal charges closest to zero
• Negative formal charges on the most electronegative atoms
Example #1 – Formaldehyde, CH2O
Hydrogens
qNo unshared electrons
q½ of shared pair = 1 electron
q1 valence electron – 1 electron = FC of 0
0
H
0
0
H
C
O
0
Check: The sum of
the formal
charges is equal to
the charge (0)
Carbon
qNo unshared electrons
q½ of shared pairs = 4 electrons
q4 valence electrons – 4 electrons = FC of 0
Oxygen
q4 unshared electrons
q½ of shared pairs = 2 electrons
q6 valence electrons – 6 electrons = FC of 0
Example #2 – Ammonium ion
Hydrogens
qNo unshared electrons
H
+
q½ of shared pair = 1 electron
+1 0
q1 valence electron – 1 electron = FC of 0
N H
0
0
H
H
0
Nitrogen
qNo unshared electrons
q½ of shared pairs = 4 electrons
q5 valence electrons – 4 electrons = FC of +1
Check: The sum of the formal charges is equal to the ionic charge (1+)
Evaluating Competing Structures: PO43In drawing a Lewis structure, the first consideration is the
number of valence electrons available.
Ø
Ø
Ø
Ø
Phosphorus: 5 valence electrons
Oxygen: 6 valence electrons each
Charge: 3- (3 electrons in excess)
Total:
5e24e3e32e-
Potential Lewis Candidates
3-
O
O
P
O
O
3-
O
P
O
P
O
P
O
P
O
O
3-
O
O
3-
O
O
O
O
O
O
3-
O
O
All of these use
exactly our 32e“budget”
Potential Lewis Candidates
-1
O
-1
+1
O
P
-1
O
-1
O
P
-1
O
P
-2
O
0
O
P
0
O
-3
O
0
O
-1
P
-1 O
3-
O
3-
O
0
0
0
0
O
-1 O
3-
O
O
0
O
-1
3-
O
-1
0
0
-1
3-
O
0
Assign formal
charges
Evaluate Using Formal Charges
0
3-
O
0
O
P
0
O
-3
O
0
Oxygen is more electronegative than
phosphorus. Negative formal charge
should be on oxygen.
Evaluate Using Formal Charges
0
3-
O
0
O
P
-1
O
-2
O
0
Oxygen is more electronegative than
phosphorus. Negative formal charge
should be on oxygen.
Evaluate Using Formal Charges
-1
3-
O
0
O
-1
P
-1 O
O
0
Oxygen is more electronegative than
phosphorus. Negative formal charge
should be on oxygen.
Evaluate Using Formal Charges
-1
3-
O
-1
0
O
P
O
-1 O
0
-1
3-
O
-1
+1
O
P
-1
O
Oxygen is more electronegative than
phosphorus. Negative formal charge
IS on oxygen.
O
-1
§ Oxygen is more electronegative
than phosphorus. Negative formal
charge IS on oxygen.
§ However, this has a greater set of
formal charges assigned than the
structure above.
6. Resonance structure
Resonance
Resonance is invoked when more than one valid Lewis structure can
be written for a particular molecule.
Benzene, C6H6
H
H
H
H
H
H
H
H
H
H
H
H
The actual structure is an average of the resonance structures.
The bond lengths in the ring are identical, and between those of single and
double bonds.
Resonance bond length and bond energy
Resonance bonds are shorter and stronger than single bonds.
H
H
H
H
H
H
H
H
H
H
H
H
Resonance bonds are longer and weaker than double bonds.
Resonance in ozone, O3
O
O
O
O
O
O
Neither structure is correct.
Oxygen bond lengths are identical, and intermediate to single and double bonds
Resonance in polyatomic ions
Resonance in a carbonate ion:
Resonance in an acetate ion:
7. Exceptions to the octet
rule
Odd number of electrons
• It is impossible to complete the octet if the total number of
valance electrons is even.
Less than an octet of valence electrons
• A second type of exception occurs when there are fewer than
eight valence electrons around an atom in a molecule or
polyatomic ion.
• Most often counter with boron and beryllium
The chemical
behaviour of BF3 is
consistent
with this
representation
Less than an octet of valence electrons
More than an octet of valence electrons
• Molecules and ions with more than an octet of electrons around
the central atom are often called hypervalent.
• Examples of hypervalent species are PCl5, SF4, AsF6-,and ICl4• Hypervalent molecules are formed only for central atoms from
period 3 and below in the periodic table (larger atoms with
unfilled d orbitals).
8. Strength of covalent
bonds
Bond enthalpy
• The stability of a molecule is related to the strengths of its
covalent bonds. The strength of a covalent bond between two
atoms is determined by the energy required to break the bond.
It is easiest to relate bond strength to the enthalpy change in
reactions in which bonds are broken.
• The bond enthalpy is the enthalpy change, ΔH, for the
breaking of a particular bond in one mole of a gaseous
substance. For example, the bond enthalpy for the bond in Cl2
is the enthalpy change when 1 mol of Cl2(g) dissociates into
chlorine atoms:
Bond length and bond energy
• Bond enthalpy is labelled D(X-Y). E.g., D(C-H)
Atomization process
Bond enthalpies and the enthalpies of
reactions
Bond enthalpies and the enthalpies of
reactions
Bond enthalpies and the enthalpies of
reactions
Bond enthalpies and the enthalpies of
reactions
Bond enthalpies and bond length
Bonds between elements become shorter and stronger as multiplicity increases.
Bond enthalpies and bond length