Lecture 3
Polar and non-polar covalent
bonds
Dr. A.K.M. Shafiqul Islam
21.07.08
Both lithium and sodium can lose a
single electron to form a cation.
• Lithium has a single electron in a 2s orbital. If the single
electron is lost, then a lithium cation is formed that has a
filled outer shell.
• Sodium has a single electron in a 3s orbital. Sodium
loses an electron to yield a sodium cation that has a
filled outer shell.
• Fluorine and chlorine contain seven electrons in
their outer shell. The addition of another
electron produces an octet.
Ionic bonds of sodium chloride
• Sodium gives up an electron to chlorine. The positively
charged sodium ions and the negatively charged
chloride ions are independent species held together by
the attraction of opposite charges, called electrostatic
attraction. This results in the formation of an ionic bond.
Structures of sodium chloride
(a) Crystalline sodium chloride.
(b) The electron-rich chloride ions are red and the electronpoor sodium ions are blue. Each chloride ion is
surrounded by six sodium ions, and each sodium ion is
surrounded by six chloride ions.
Covalent bonds in fluorine
• Two fluorine atoms can each attain a filled
second shell by sharing their unpaired valence
electrons, resulting in a covalent bond.
Covalent bonds of hydrogen
• Two hydrogen atoms can form a covalent bond
by sharing electrons. Each hydrogen acquires a
stable, filled first shell
Covalent bond of hydrogen chloride
• Hydrogen and chlorine can form a covalent bond
by sharing electrons. Hydrogen fills its only shell
and chlorine achieves an outer shell of eight
electrons
Formation of a proton and a
hydride
• A hydrogen atom loses its valence electron to
form a proton, a positively charged hydrogen
ion. A hydrogen atom gains an electron to form
a hydride, a negatively charged hydrogen ion.
The Electronegativities of Selected
Elements
Covalent bonds of oxygen, nitrogen,
and carbon
•
•
•
Because oxygen has six valence electrons, it needs to form two covalent
bonds to achieve an outer shell of eight electrons.
Nitrogen, with five valence electrons, must form three covalent bonds.
Carbon, with four valence electrons, must form four covalent bonds.
Use the symbols d- on the most electronegative atom
and the symbol d+ on the most electropositive atom
• Oxygen is the most electronegative and thus is
given a d• Carbon is the most electropositive and is given a
d+
Polar covalent bonds of hydrogen
chloride, water, and ammonia
• A polar covalent bond is the result of the differences in
electronegativities of the two atoms involved in the bond.
• A polar covalent bond has a slight positive charge on the
most electropositive atom and a slight negative charge
on the most electronegative atom.
Polar bond of hydrogen chloride
• The direction of bond polarity is indicated with
an arrow. The arrow is drawn along the bond,
toward the most electronegative atom.
Bonding types
The three major types of bonding found in most compounds are ionic,
polar covalent, and nonpolar covalent bonds.
• An ionic bond is the transfer of electrons.
• A covalent bond is the sharing of electrons.
• A polar covalent bond is between atoms of different electronegativity.
• A nonpolar covalent bond is between atoms of similar
electronegativity.
• Potassium fluoride and sodium chloride are examples of ionic bonds.
An O-H bond and an N-H bond are examples of polar covalent bonds.
A C-H bond and a C-C bond are examples of nonpolar covalent
bonds.
Electrostatic potential maps for
lithium hydride, hydrogen, and
hydrogen fluoride
• For lithium hydride, the hydrogen atom has the greater
electron density than the lithium atom.
• In hydrogen fluoride, the hydrogen has less electron
density than a hydrogen in a hydrogen molecule.
Explanation of colors on
electrostatic potential maps
• The colors indicate the distribution of charge in
the molecule. Red is for electron-rich areas and
blue is for electron-deficient areas.
Lewis structures of water, hydronium ion,
hydroxide ion, and hydrogen peroxide
• Lewis structures show the valence electrons represented
as dots. Lone-pair electrons or nonbonding electrons
are valence electrons not used in bonding. A formal
charge is the difference between the number of valence
electrons minus the number of lone-pair electrons minus
half the number of bonding electrons.
Electrostatic potential maps for the hydronium
ion, water, and the hydroxide anion
• In the hydronium ion, the oxygen has a slight negative
charge.
• In the hydronium ion, the oxygen has a negligible charge.
• The oxygen has the greater charge in the hydroxide ion.
Lewis dot structures of nitrogen
compounds.
• Nitrogen has five valence electrons. In ammonia, nitrogen
has five valence electrons. In the ammonium ion, nitrogen
has lost one valence electron. Nitrogen has gained one
valence electron in the amide anion. In hydrazine, one
valence electron is shared between the two nitrogens.
Lewis structures of carbon compounds
• Carbon has four valence electrons. Carbon has lost one
valence electron in the methyl cation. Carbon has
gained one valence electron in the methyl anion.
Carbon has four valence electrons in the methyl radical.
In ethane, one valence electron is shared between the
carbon atoms.
Lewis structures of hydrogen,
bromide, bromine, and chlorine
• Hydrogen has one valence electron. Hydrogen ion has lost one
electron. Hydride ion has gained one electron. Hydrogen radical has
the valence electron. Bromide ion has gained one electron.
• Bromide radical has the seven valence electrons. Each bromine
atom in the bromine molecule has seven electrons. Each chlorine
atom in the chlorine molecule has seven electrons.
Number of covalent bonds with hydrogen,
halides, oxygen, nitrogen, and carbon
• Generally, hydrogen forms one covalent bond, oxygen
forms two covalent bonds, nitrogen forms three covalent
bonds, halides form one covalent bond, and carbon
forms four covalent bonds.
Lewis dot structures of methyl bromide, dimethyl
ether, formic acid, methylamine, and nitrogen
• Hydrogen and bromide each forms one covalent bond. Oxygen forms
two covalent bonds. Nitrogen forms three covalent bonds. Carbon
forms four covalent bonds. A single bond contains two electrons, a
double bond contains four electrons, and a triple bond contains six
electrons.
Lewis structures of methyl bromide, dimethyl
ether, formic acid, methylamine, and nitrogen
• A pair of shared electrons can be shown as a
line between two atoms. In formic acid, two
electron pairs are illustrated with a double bond.
In nitrogen, three electron pairs are shown with
a triple bond.
Drawing Lewis structures
• For HNO2, 18 valence electrons are present. Place the
valence electrons around the atoms that supply them.
Check to see if all atoms have an octet. In the first
structure, nitrogen does not have an octet. A pair of
electrons from the oxygen is used to make a double
bond between the nitrogen and the oxygen to give
nitrogen an octet.
Kekulé structures of methyl bromide, dimethyl
ether, formic acid, methylamine, and nitrogen
• In Kekulé structures, the bonding electrons are
drawn as lines and the lone-pair electrons are
left out entirely.
Condensed structures of methyl bromide, dimethyl
ether, formic acid, methylamine, and nitrogen
• Condensed structures are simplified structures that
omit some of the covalent bonds and listing atoms
bonded to a particular carbon or other atom.
Kekulé and condensed structures of
2-bromo-5-chlorohexane
• In the Kekulé structure, all of the covalent bonds
are shown. In the condensed structure, the
atoms other than hydrogen are shown hanging
from the carbon.
Kekulé and condensed structures of
hexane
• In the Kekulé structure, all of the covalent bonds
are shown. In the condensed structure, the
repeating CH2 groups can be shown in
parentheses.
Kekulé and condensed structures of
4-methyl-2-hexanol
• In the Kekulé structure, all of the covalent bonds
are shown. In the condensed structure, groups
bonded to a carbon can be shown in
parentheses to the right of the carbon or
hanging from the carbon.
Kekulé and condensed structures of
4,4-dimethyl-1-pentanol
• In the Kekulé structure, all of the covalent bonds are
shown. In the condensed structure, groups bonded to a
carbon can be shown in parentheses to the right of the
carbon or hanging from the carbon. Groups bonded to
the far-right carbon are not put in parentheses.