Chem. 200 Dr. Saidane

Dr. Saidane
Chem. 200
Lecture Notes
CHAPTER 8, Part 1
Chemical Bonding
There are three major types of chemical bonds: Ionic, covalent and metallic.
Metallic bonds are bonds between metallic atoms. A metallic solid consists of cations
surrounded by a sea of electrons. The nucleus and the inner shells (core) electrons
represent the cation.
The valence electrons represent the sea of electrons. The
valence electrons are constantly moving from one atom to the other. The charge of
the electron sea cancels the charges of the cations. The mobility of the sea of
electrons makes metals conductor of electricity.
Ionic bonds are bonds between metals or nonmetals. An ionic bond is the attraction
between the opposite charges of cations and anions. When an ionic bond is formed,
the valence electrons of the metal are transferred to the valence shell of the nonmetal.
This transfer of electrons results in a complete (filled) valence shell for each of the
atoms in the bond.
Covalent bonds are bonds between two nonmetals. A covalent bond is a pair of
electrons shared between two atoms.
Ionic and covalent bonds are two extreme descriptions. All ionic bonds have some
covalent character.
Actual bonds lie somewhere between purely ionic and purely
covalent. The electronegativity difference between two elements is an indication of the
ionic or covalent character of a chemical bond.
The electronegativity of an element is its electron pulling power when it is part of a
Electronegativity tends to be high towards the upper right corner of the periodic table
and low on the lower left.
Elements with low electronegativities are often called electropositive.
An atom with a high electronegativity has a strong pulling power on electrons,
particularly for the electron pair it shares with its neighbor. The outcome of the tugof-war is that the more electronegative atom has a greater share in the electron pair of
the covalent bond.
Bonds formed between elements with a large difference in electronegativity tend to
have significant ionic character in their bonding.
Bond Polarizability and Polarity
In an ionic bond, cations are close to anions. The electron cloud of the anion is
distorted by the pull of the cation. This distortion process is called polarization.
Small highly charged cations are highly polarizing. Large, electron rich anions are
highly polarizable.
In a covalent bond formed between two atoms of different electronegativities, the
electron cloud of the less electronegative element will be distorted by the pull effect
of the most electronegative element. Atoms with high electronegativity are highly
polarizing. Electropositive atoms are highly polarizable. And the bond is polar.
Bond polarity is related to electronegativity difference. When the electronegativity
difference is large (greater than about 2), the bond is considered ionic. When the
difference is zero, the bond is considered nonpolar and is purely covalent. For all
the values between zero and two the bond is considered to be polar covalent.
Chemical bonds are represented using Lewis structures.
Lewis Symbols for Atoms
A Lewis symbol consists of the chemical symbol of the element and a dot for each of
 
its valence electrons. H , K , Mg
A single dot represents an electron alone in an orbital. A pair of dots represents two
paired electrons in the same orbital.
B. Lewis Structures for Ionic Compounds
We need to follow these steps:
1. Write the formula of the compound. Ex: Fe2O3.
2. Write the Lewis symbols of each element.
3. Calculate the total number of valence electrons in the compound. Ex: Fe has 3
and O has 6  Total number = (3x2) + (6x3) = 24.
4. Place the elements by alternating metal, nonmetal starting with the element that
has the highest number of atom in the formula. Ex: In Fe2O3 we start with
oxygen and we place the atoms as follows: O
Fe O Fe O.
5. Draw the arrows showing the transfer of valence electrons from the metal to the
6. Write the final structure using brackets for anions. Ex: [O]2- Fe3+[O]2-Fe3+[O]2(each atom must be surrounded by dots). The total number of dots should be
equal to the total number of valence electrons calculated in step 3.
When a main group metal atom forms a cation, it loses its s and p valence electrons
and acquires the configuration of the preceding noble-gas atom.
Ex: Al: [Ne] 2s2 2p1 and Al3+: [Ne]
The heavier metals in groups 13 and 14 behave similarly, but the resulting core
consists of the noble-gas configuration and an additional complete subshell of delectrons.
Ex: Ga: [Ar] 3d10 4s24p1 and Ga3+: [Ar] 3d10
The electrons in the filled d-subshell are not represented by dots in the Lewis
When p-block elements acquire electrons and form anions, they do so until they have
reached the electron configuration of the following noble gas.
Ex: O: [He] 2s22p4 and O2-: [Ar]
The formation of ionic bonds is represented in terms of Lewis symbols by the loss or
gain of electrons (dots transfer) until both species have reached an octet of electrons.
C. Lewis Structures for Covalent Compounds
Nonmetals share electrons with one another until they have completed their octet (or
duplet for hydrogen). A Lewis structure of a covalent compound shows the arrangements
of electrons as shared pairs (or lines) and lone pairs (pairs of dots).
To write the Lewis structure of a covalent compound we need to follow these steps:
Calculate the total number of valence electrons in the compound. Ex: CH4 has 8
valence electrons (4 electrons for C and one electron for each H).
2. Arrange the atoms in the compound. Choose the atom with the lowest ionization
energy for the central atom. Place the remaining atoms symmetrically around the
central atom. Ex: SO2 is OSO not SOO.
3. Place one electron pair (or a line) between each pair of bonded atoms. Each
bonded pair of atoms must have at least a single bond between them.
4. Complete the octet of each atom (duplet for H) by placing the remaining valence
electron as electron pairs around the atoms. If there are not enough electron pairs
form multiple bonds.
Exception to the Octet Rule
Some elements do not have an octet in their Lewis structures. They are exceptions to
the octet rule.
1. Atoms with less than an octet. Hydrogen and boron are the elements than can
have a valence shell with less than eight electrons. H has two electrons (duplet).
Boron tends to form bonds in which six electrons surround it, such as in BF3.
2. Atoms with an expanded valence shell. More than eight electrons can surround
elements in periods 3 and later when they combine with the highly
electronegative elements F, Cl, and O. In these cases we say that these atoms
have an expanded valence shell in which bonding involves electrons in the d
orbitals. Ex:
Phosphorous is surrounded by eight electrons in PCl4+, it is
surrounded by 10 electrons in PCl5, and is surrounded by 12 electrons in PCl6-.
3. Atoms with an odd number of valence electrons.
These atoms are called
radicals. They are generally highly reactive. Ex. NO has 11 valence electrons
(O has an octet but N has seven electrons.)
Resonance Structures
In some Lewis structures of covalent compounds with multiple bonds, the multiple
bonds can be written in several equivalent locations. For each Lewis structure of the
compound with a different location of the multiple bond we have a resonance structure
and each one of them is called resonance hybrid.
Lewis Acids and Bases
A Lewis acid is an element that has an empty valence orbital. It is an electron pair
A Lewis base is an element that donates a lone pair (non bonded pair).
When a Lewis base shares its lone pair with the Lewis acid, a Lewis acid-base
complex is formed. The bonding electrons are transferred from the base to the acid
(similar to an electron transfer in an ionic bond). After being transferred, the bonding
electron pair is shared by the two atoms (similar to a covalent bond). This type of
bond is called a polar coordinate bond.
Compounds such of boron and aluminum are Lewis acids, and form coordinate polar
bonds with halogens. Ex. Al2Cl6 is a compound that has 6 covalent bonds and 2
coordinate bonds.
Skills you should have mastered
1. Distinguish between ionic, covalent, and metallic bonds.
2. Explain the significance of electronegativity and what the ionic or covalent
character of a bond means.
3. Predict and explain periodic trends in the polarizability of anions and the
polarizing power of cations.
4. Explain how the polarity of a bond depends on the electronegativities of the two
atoms in the bond.
Define resonance hybrid and explain its relationship to the individual Lewis
structures that contribute.
6. Explain the characteristics of Lewis acids and bases and how they form bonds.
Use a table of electronegativities to predict which of two bonds has greater ionic
or covalent character.
Predict the polarity of a bond.
Predict the chemical formula of a binary ionic compound and write its formula
with Lewis structures.
Draw the Lewis structures of a molecule or ion.
Write the resonance structures for a molecule.
1. Describe the general characteristics of radical species.
2. Predict which atoms are likely to form molecules in which they have an expanded
3. Predict whether a compound can act as a Lewis acid or a Lewis base.