Basic Concepts of Chemical
Bonding
Bonding
• Ionic – Electrostatic forces that exist
between two ions of opposite charges
transfer of electrons ( metal by non-metal)
• Covalent – sharing of e’s between two atoms
(two non-metals)
• Metallic – found in metals where the atoms
are bound to their neighbors but allow free
movement of electrons (chapter 23)
Lewis Symbols and the Octet Rule
• G.N. Lewis
• Valence e’s – available
to form bonds – reside
in incomplete outer
shell
• Lewis – developed
system for noting val
e’s – one dot for each
val electron
• Place dots next to
element following
Hund’s rule.
Ionic Bonding and Energy
Relationships
• Electrostatic attraction
of Cations + by
Anions –
• In order to maximize
the attraction ionic
solids exist in lattices
• Lattice energy –
energy required to
separate the crystalline
solid into a gas (a mole
of solid into gasious
ions)
• Halite NaCl Rock Salt
Energy Relationships
• removing an e from a metal is endothermic
• when a non-metal gains an e it is exothermic
• If e transfer was the only factor all ionic cmps
would be endothermic
• Because of the stability of the lattice
structure (going to a lower energy state)
releasing energy, all ionic cmps are
exothermic.
Energy Relationships
• Lattice Energy – tends
to increase as the
charge of the ions
increases and the size
of the ions decreases
• E = k Q1Q2
d
Q1Q2 = charge of part
d = distance between
k = constant
• Bauxite
Ionic Bond Formation
• The representative elements form ions that
have the noble gas configurations. Metals
lose electrons and non-metals gain
electrons.
• The transitional metals do not always form
electron configurations of the noble gases.
The s sublevel electrons are the first to be
transferred followed by the d sublevel
electrons if necessary.
Ionic Bond Formation
• Many stable ions are formed by emptying the
s orbital or by leaving the d sub-shell full
d10, half empty d5, or empty d0.
• Polyatomic Ins – form when molecules gain
or lose electrons. It is often unclear how
(which atom gains or loses electrons) but the
overall charge of the Ion is greater than the
molecule. (has more of less electrons)
Size of Ions
• Atoms increase in size
going from right to left
in a period.
• Cations are smaller
than their parent
atoms. Due to the +
charge and fewer
screening electrons
• Anions are larger than
their parent atoms .
Size of Ions
• The effective nuclear
charge decreases
because of the
increase of screening
electrons.
Isoelectric Series
• Is a series of ions or
atoms that have the
same electron
configuration. Within a
series the greater the
atomic # (greater the
number of protonseffective nuclear
charge) the smaller the
species.
Covalent Bonding
• Electron pair sharing between two atoms –
electrons attracted to both nuclei at the same
time – covalent bonding results
• Lewis described bonding patterns using
electron dot symbols (Lewis structures)
• Multiple bonds are formed when more than
one pair of electrons is shared between
atoms.
Single, Double, and Triple Bonds
• Single bonds – covalent
bond one pair of shared
electrons – ex: Cl-Cl
structural, Cl2 molecular
formula
• Double bonds – 2 pair of
shared electrons,
O=O structural, O2
molecular
• Triple bonds – 3 pair of
shared electrons N N,
Structural, N3 molecular
Bond Length and Strength
• In general the distance between bonded
atoms decreases as the number of shared
electron pairs increases.
• A triple bond is stronger than a single bond
but not three times stronger. The second and
third bonds are weaker than the first bond.
The first bond in a triple bond is stronger
than a “normal” single bond.
Strength of Covalent Bonds
• Bond dissociation energy – Bond energy –
enthalpy required to break a covalent bond is
the average of the bond enthalpies of
different molecules
• Delta H (rxn) = Sum (bond energies of bonds
broken) – Sum (bond energies of bonds
formed) The aprox enthalpy of a reaction
can be predicted by bond energies. This
relates well to the heat of formations data
that is more accurate.
Strength of Covalent Bonds
• Between atoms of comparable size the
greater the bond strength the shorter the
bond length.
Bond Polarity and
Electronegativity
• Covalent bonds are the
result of electron
sharing – types • Non-polar – when e’s
are shared equally
between nuclei –
diatomic elements are
an example
• Polar covalent – when
e”s are not shared
equally
Electronegativity
• If the electronegativities
are equal (i.e. if the
electronegativity
difference is 0), the bond
is non-polar covalent
• If the difference in
electronegativities
between the two atoms is
greater than 0, but less
than 2.0, the bond is
polar covalent
• If the difference in
electronegativities
between the two atoms is
2.0, or greater, the bond
is ionic
• Electronegativity – is a
quantity that describes an
elements ability to compete
for electrons in a covalent
bond. The greater the
number the better it
competes for e’s.
Ways of noting partial +,charges
Drawing Lewis Structures of
Covalent molecules
• Sum the valence e’s from all atoms in
species
• Write the atomic symbols showing how
atoms are connected – draw a single bond
between atoms
• Complete the octets of the atoms bonded to
the central atom (peripheral atoms)
• Place left over electrons on the central atom
even if it results in the central atom having
Drawing Lewis Structures
more than an octet.
• If there are not enough e’s to give the central
atom an octet, form multiple bonds by pulling
terminal e’s from the peripheral atoms and
placing them into the bond with the central
atom.
• Examples
Formal Charges
• Formal charges are a way to assign all the
valence elcectrons in a molecule to a “parent
atom”
• Rules
1. all bonding e’s are divided equally
between atoms that form bonds.
2. all non-bonding e’s are assigned to the
atom on which they reside.
Formal Charges
• The formal charge is the number of valence e’s in
the isolated atom (usually the group number in the
periodic table) minus the number of electrons
assigned by the rules
• When several different Lewis structures are
plausible, the one in which the formal charges are
minimized is generally the preferred one.
• If the same choose the structure with negative
charges residing on the more electronegative
atoms
Calculating Formal Charges
Examples
Practice Questions
• Identify formal charges that are not zero
Answers
Exceptions to the Octet Rule
• Most of the second period elements (C,O)
are always observed with octets. Other
elements do not easily achieve or ever
achieve octets.
• Molecules that contain odd numbers of e’s
although they are uncommon and tend to be
reactive are exceptions.
• Light elements (H,Li,Be,B) tend to be
surrounded by less than and octet of
electrons.
Exceptions to Octet Rule
• Third period elements and below in the
periodic table are capable of expanding their
octets because of the unfilled d orbitals thus
having greater than eight electrons.
• Examples on page 285-288