Ionic and Covalent Bonding
Ionic Bond is a chemical bond formed by the electrostatic attraction between
positive and negative ions.
Covalent Bond is a chemical bond that is sharing electrons between two atoms.
Metallic Bonds are bonds where electrons are bonded to several neighboring
atoms.
Valence electrons are bonded electrons residing in the incompletely filled outer
electron shell.
Lewis electron dot structure - simple way of showing valence electrons
of an atom or ion are represented by dots placed around the letter symbol of the
element.
Octet Rule - Atoms tend to gain, lose, or share electrons until they are surrounded
by 8 valence electrons
Na(s) + Cl2(g)
Na+ +
[ Cl ]-
Na+ ∆Hf = -240 kJ/mol
Cl- ∆Hf = -167.2 kJ/mol
Lattice energy is the change in energy that occurs when one ionic solid is
separated into isolated ions in the gas phase.
NaCl(s)
Na+(g) + Cl-(g)
∆Hlattice = +786 kj/mol
The magnitude of the lattice energy depends on the charges of the ions, size and
arrangement of the atoms.
Ionic Radii
Ionic radius is a measure of the size of the spherical region around the nucleus of
an ion within which the electrons are most likely to be found.
Example: From X ray diffraction experiments an iodine- iodine distance is 426 pm.
Therefore the radius of an iodine atom = 1/2 x 426pm = 213 pm
Ionic radii increase down any column because of the addition of electron shells
Isoelectric - refers to the different species having the same number and
configuration of electrons
Cation
Na+
Mg+2
Al+3
Radius (pm)
95
65
50
Covalent Bonds
Lewis electron dot formula - is a formula using dots to represent valence electrons.
Bonding pairs vs nonbonding pairs
Single covalent bond -a bond sharing 1 pair of electrons from two atoms.
Coordinate covalent bond - a bond formed when both electrons of the bond are
donated by one electron.
Multiple Bonds
Single bond- a covalent bond in which a single pair of electons is shared by two
atoms.
Double bond - a covalent bond in which two pairs of electrons are shared by two
atoms.
Triple bond - a covalent bond in which three pairs of electrons are shared by two
atoms.
Electron pairs are not necessarily shared equally.
H : H
Na+ Cl-
H :Cl
Electronegativity
A measure of the ability of an atom in a molecule to draw bonding electrons
to itself.
X=
I.E. is the ionization energy
E.A. is the electron affinity
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Use electronegativity values to arrange the following bonds in order of increasing
polarity: P - H , H - O, C - Cl.
You can use an electronegativity scale to predict the direction in which the electron
shift during bond formation; the electrons are pulled toward the more
electronegative atom.
δ+
H Cl δ-
Writing Lewis Dot structures
1. Calculate the total number of valence electrons for the molecule.
2. Write the skeleton structure of the molecule or ion
3. Distribute electrons to the atoms surrounding the central atom(or atoms)
to satisfy the octet rule.
4. Distribute the remaining electrons as pairs to the central atom (or atoms)
Atoms that often form multiple bonds are: C, N, O, and S.
Sulfur dichloride, SCl2, is a red fuming liquid used in the manufacture of
insecticides. Write the Lewis formula for the molecule
Example: Write the electron dot structure for CO2 and the chlorite ion, ClO2-
Exceptions
Less than an octet: BF3
More than an octet: PCl5, SF4
Formal charge - is an aid in deciding between two or more possible Lewis dot
structures are equally shared between bonded atoms and that the electrons of each
lone pair belongs completely to one atom.
To determine formal charges you begin by writing the possible Lewis dot
structures.
Rules for formal charges are:
1. Half of the electrons of a bond are assigned to each atom in the bond
2. Both electrons of a lone pair are assigned to the atom to which the lone
pair belongs
Formal charge = valence electrons on free atom - 1/2 (number of electrons in a
bond) - (number of lone pair electrons)
Example cyanide ion, CN-.
Rules for Determining the correct Lewis structure
Rule A: Whenever you can write several Lewis formulas for a molecule, choose
the one having the lowest magnitudes of formal charges
Rule B: When two proposed Lewis structures for a molecule have the same
magnitudes of formal charges, choose the one having the negative formal charge
on the more electronegative atom.
RuleC: When possible, choose Lewis structures that do not have like charges on
adjacent atoms
Example: Determine the correct Lewis structure for COCl2
Resonance
Delocalized bonding - a type of bonding in which a bonding pair of electrons is
spread over a number of atoms rather than localized between the two.
The most stable resonance structure is the one where the atoms bear the smallest
formal charge and any negative charge resides on the most electronegative atom.
Example: Describe the electron structure of the carbonate ion, CO3-2, in terms of
electron dot formulas.
Bond Length - is the distance between the nuclei in a bond.
Covalent radii are values
assigned to atoms in such a
way that the sum of covalent
radii of atoms A and B
predicts an approximate
A-B bond length
Bond Order - is defined in Lewis structure terms, is the number of pairs of
electrons in a bond.
Example: consider the molecules N2H4, N2, and N2F2. Which molecule has the
shortest nitrogen-nitrogen bond? Which has the longest nitrogen-nitrogen bond?
Bond Energy is a measure of the strength of a bond.
Bond dissociation energy - is the energy required to break a particular bond in a
molecule.
A B bond energy (BE) is the average enthalpy change for the breaking of an A B
bond in a molecule in the gas phase.
Calculate the bond energy for a C H in methane, CH4.
CH4(g)
C(g) + 4H(g);
∆H= 1662 kJ
In general, the enthalpy of a reaction is (approximately) equal to the sum of the
bond energies for bonds broken minus the sum of the bond energies for bonds
formed.
Bond energies are perhaps of the
greatest value when you try to
explain heats of reactions or to
understand the relative stabilities
of compounds
Example: Use bond energies to estimate the enthalpy change for the combustion of
ethylene, C2H4, according to the equation
C2H4(g) + 3O2(g)
2CO2(g) + 2H2O(g)