Chapter 8: Basic Concept of
Chemical Bonding
Valence Electrons
• The electrons in the highest occupied level of an element’s
atoms.
• The number of valence electrons largely determines the
chemical properties of an element.
• Can be determined by looking at the electron configuration
or Lewis structure (dot diagram) for an element.
• Alkali metals have 1, Alkaline Earth metals have 2, Transition
metals have 2 or 3, Boron group has 3, Carbon group has 4,
Nitrogen group has 5, Chalcogens have 6, Halogens have 7,
and the Noble Gases have 8 (an octet).
• Chemist Gilbert Lewis explained that atoms form certain
kinds of bonds due to atoms tendency to achieve the
electron configuration of a noble gas (octet rule-Atoms tend
to gain lose, or share electrons until they are surrounded by
eight valence electrons).
Electronegativity
• The tendency
for the atoms
of the element
to attract
electrons
when they are
chemically
combined with
atoms of other
elements.
Ionic Bonding
• Two elements with very different electronegativities
(difference between 4.0 and 1.7).
• Elements with high electronegativity pull electron(s) more
strongly and therefore steal the electron(s) from the element
with a low electronegativity.
• The stronger elements therefore becomes negatively
charged (anion) since it has gained electrons and the
weaker elements becomes positively charged (cation) since
it has lost electrons.
• Since they have opposite charges, they attract one another
by electrostatic forces which creates the bond.
• Usually between one non-metal and one metal.
• Lewis Structure for Ionic Bonds:
Na
Cl
become Na+[ Cl ]-
Ionic Bond Characteristics
• Compounds are called formula units
• Most are crystalline solids at room temperature (pg 422 in regular chem
text has some great pictures)
• In solids NaCl, each sodium cation is surrounded by six chloride anions
and each chlorine anion is surrounded by six sodium cations. Giving
them both a coordination number (the number of ions of opposite charge
that surrounds the ion in a crystal) of 6.
• In this arrangement, each ion is attracted strongly to each of its
neighbors and repulsions are minimized, which results in a very stable
structure.
• High melting point (usually above 300 °C)
• High solubility (dissolving ability) in water since water is polar.
• Good conductor of electricity either when melted (the orderly crystal
structure breaks down allowing cations and anions to move freely
through the liquid) or when mixed in water (ions also become free to
move around allowing electrons to flow).
• Brittle, not very malleable because when struck with a hammer, the blow
tends to push ions of like charge into contact and they repel, shattering
the crystal.
Energetics of Ionic Bond Formation
• When atoms lose electrons to form ions, energy is required (taken in) so it
is an endothermic process. For example, when Na becomes Na+ it
requires 496 kJ/mol.
• When atoms gain electrons to form ions, energy is given off so it is an
exothermic reaction. For example, when Cl becomes Cl- it gives off 349
kJ/mol.
• Therefore, if the transfer of an electron from one atom to another were the
only factor in forming an ionic bond, the overall process would be an
endothermic process the requires 496-349= 147 kJ/mol. But this is only if
the ions formed far apart from each other.
• The principle reason that ionic compounds are stable is the attraction
between ions of unlike charge. This attraction draws the ions together,
releasing energy and causing ions to form a solid lattice.
• A measure of just how much stabilization results from arranging oppositely
charged ions in an ionic solid is given by the lattice energy, which is the
energy required to completely separate a mole of a solid ionic compound
into its gaseous state. For example, the lattice energy for NaCl is
788kJ/mol, which means that NaCl needs to take in that much energy to
be able to make the ions break apart. Therefore, when they come together
the opposite occurs…788 kJ/mol of energy is given off so the formation
process is actually very exothermic. (See picture on pg 300)
• Lattice energies of given in Table 8.2 on pg 301 actually increase as the
charges on the ions increase and as their radii decrease.
Covalent Bonding
• Two elements with similar electronegativities
(difference is less than 1.7).
• Since they have similar pull on the electron(s)
they end up both holding onto it and the
electron(s) and sharing them, which creates the
bond.
• Usually both non-metals.
• Lewis Structure for covalent bond:
H
H becomes
H H
Covalent Bonding Characteristics
•
•
•
•
Compounds are called molecules.
Can be solid, liquid or gas at room temperature.
Low melting point (usually below )
High to low solubility in water (depending on the
molecules’ polarity)
• Poor to nonconducting when dissolved in water
(aqueous).
• A greater range of physical properties is found due to the
widely varying intermolecular attractions (polar and
nonpolar bonding)
• Can have single, double or triple covalent bonds
between two atoms.
Polarity
• Polar bonds are when one atoms has a slightly
stronger pull on the electrons
H2 that are being
HF
shared
(difference between 1.7 and 0.3).
Therefore, one end of the bond is slightly more
negative (the one that is slightly stronger) and
one is slightly more positive (the one that is
slightly weaker).
• All the diatomic elements (Nitrogen, Oxygen,
Fluorine, Chlorine, Bromine, Iodine and
Hydrogen) have non-polar covalent bonds since
they share the electrons equally.
Polar Bonds vs. Polar Molecules
• The presence of a polar bond in a molecule
often makes the entire molecule polar, where
one end of the molecule has a slightly more
negative charge and the other end is slightly
more positive (called a dipole).
(Ex: H20 and NH3)
• However, if the polar bonds are arranged
symmetrically around the molecule the entire
molecule is actually non-polar.
(Ex: CO2 and CH4)
Drawing Lewis Structures for
Covalent Bonds
• Count the total valence electrons for the molecule: To do
this, find the number of valence electrons for each atom in the
molecule, and add them up.
• Determine the center atom(s). This will be the first atom
listed in the formula, unless, the first atom listed is Hydrogen.
• Put other elements in formula around central atom and
connect them with single bonds. If more than 4 outer atoms
they will have to be bonded to others outer atoms.
• Fill in eight electrons for each non-central atom. There are
a few exceptions to this rule we will learn later.
• Fill in eight electrons around central atom. Also some
exceptions we will learn later.
• Count how many total electrons you now have on your
drawing. If count is less the valence electrons determined
earlier add the extras to the central atom. If count is more
than the valence electrons determined earlier change some
unshared pairs into double or triple bonds.
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-if all electrons were shared equally.
• To calculate formal charge, we assign the electrons to the
atom as follows:
– All unshared (nonbonding) electrons are assigned to the atom on
which they are found.
– For any bond, single double or triple, half of the bonding electrons are
assigned to each atom in the bond.
• The formal charge of each atom is then calculated by
subtracting the number of electrons assigned to the atom
from the number of 4valence electrons in the isolated atom
originally.
• Practice website
• We generally choose the Lewis structure in which the atoms
bear formal charges closet to zero.
• We also generally choose the Lewis structure in which any
negative charges reside on the more electronegative atoms.
Resonance Structures
• We sometimes encounter molecules and ions in
which the experimentally determined
arrangement of atoms is not adequately
described by a single Lewis structure.
• Consider a molecule of ozone, O3, which is a
bent molecules with two equal O—O bond
lengths.
• See example on page 319 in book.
• See website example
Exceptions to the Octet Rule
• Three main types of exceptions for covalent
bonding:
– Molecules and polyatomic ions containing an
odd number of electrons
– Molecules and polyatomic ions in which an
atom has fewer than an octet of valence
electrons.
– Molecules and polyatomic ions in which an
atom has more than an octet of valence
electrons.
Molecules and polyatomic ions
containing an odd number of electrons
• In some molecules, the total number of
valence electrons is odd when two atoms
bond. For example, NO since N has 5 and
O has 6, the total is 11 so an octet could
never be achieved as you need to have an
even number for that.
Molecules and polyatomic ions in which
an atom has fewer than an octet of
valence electrons.
• Hydrogen, Beryllium and Boron have two few valence electrons
to ever obtain a full octet. See this.
• Hydrogen can have at most 2 valence electrons after it shares
its electron with another atom.
• Beryllium will have 4 valence electrons after it has finished
bonding.
• Boron will have 6 valence electrons after it shares its valence
electrons with other atoms.
Molecules and polyatomic ions in which
an atom has more than an octet of
valence electrons
• Elements in periods 3, 4, 5, 6 and 7 can
expand their octet to have 10, 12, or 14
valence electrons. See how this can happen.
Strengths of Covalent Bonds
• The stability of a molecule is related to the
strengths of the covalent bonds it contains.
• The strength of a covalent bond between two
atoms is determined by the energy required to
break that 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… which we will cover in
Chapter 5 (Section 5.3).
Metallic Bonding
• Metals are made up of closely packed cations
rather than neutral atoms. The cations are
surrounded by mobile valence electrons
(delocalized electrons), which drift freely from
one part of the metal to another.
• The attractions of the free-floating valence
electrons for the positively charged metal cations
creates the bond that holds the metal together.
Metallic Bond Characteristics
• Good conductors of electrical current because electrons can flow
freely in them. As electrons enter one end of a bar of metal, an equal
number leave the other end.
• Ductile and malleable due to the metals ability to be forced through
a narrow opening to produce a wire or shaped with a hammer
without shattering. A sea of drifting valence electrons insulates the
metal cations from one another. When a metal is subjected to
pressure, the metal cations easily slide past one another like ball
bearing immersed in oil.
• Also crystalline solids at room temperature (except for mercury).
Three arrangements of the atoms
– Body-centered cubic
– Face-centered cubic
– Hexagonal close-packed
Other Intermolecular Forces
• Molecules are often attracted to each other by a
variety of other forces, although weaker than ionic
or covalent bonds.
– Van der Waals forces:
• Dispersion forces-the weakest of all molecular interactions
caused by the motion of electrons.
• Dipole interactions-occurs when polar molecules are attracted
to one another
– Hydrogen bonds:
• Attractive forces in which a hydrogen covalently bonds to a very
electronegative atom is also weakly bonded to an unshared pair
of electrons of another electronegative atom. Since hydrogen
only has one electron when it bonds with a highly
electronegative atom, the nucleus of the hydrogen atom
becomes very electron deficient. Therefore, the sharing of a
nonbonding electron pair on a nearby electronegative atom
compensates for the deficiency.