Covalent Bonding
Name_______________________
In ionic bonding, electrons are completely transferred from one atom to another. In the process of
either losing or gaining negatively charged electrons, the reacting atoms form ions. The oppositely
charged ions are attracted to each other by electrostatic forces, which are the basis of the ionic
bond.
For example, during the reaction of sodium with chlorine:
sodium (on the left) loses its
one valence electron to chlorine
(on the right),
resulting in
a positively charged sodium ion
(left) and a negatively charged
chlorine ion (right).
Notice that when sodium loses its one valence electron it gets smaller in size, while chlorine grows larger
when it gains an additional valence electron. This is typical of the relative sizes of ions to atoms. Positive
ions tend to be smaller than their parent atoms while negative ions tend to be larger than their parent. After
the reaction takes place, the charged Na+ and Cl- ions are held together by electrostatic forces, thus forming
an ionic bond. Ionic compounds share many features in common:
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Ionic bonds form between metals and nonmetals.
In naming simple ionic compounds, the metal is always first, the nonmetal second (e.g.,
sodium chloride).
Ionic compounds dissolve easily in water and other polar solvents.
In solution, ionic compounds easily conduct electricity.
Ionic compounds tend to form crystalline solids with high melting temperatures.
This last feature, the fact that ionic compounds are solids, results from the intermolecular forces
(forces between molecules) in ionic solids. If we consider a solid crystal of sodium chloride, the
solid is made up of many positively charged sodium ions (pictured below as small gray spheres)
and an equal number of negatively charged chlorine ions (green spheres). Due to the interaction
of the charged ions, the sodium and chlorine ions are arranged in an alternating fashion as
demonstrated in the schematic. Each sodium ion is attracted equally to all of its neighboring
chlorine ions, and likewise for the chlorine to sodium attraction. The concept of a single molecule
becomes blurred in ionic crystals because the solid exists as one continuous system. Forces
between molecules are comparable to the forces within the molecule, and ionic compounds tend
to form crystal solids with high melting points as a result.
Covalent Bonding
The second major type of atomic bonding occurs when atoms share electrons. As opposed to
ionic bonding in which a complete transfer of electrons occurs, covalent bonding occurs when two
(or more) elements share electrons. Covalent bonding occurs because the atoms in the
compound have a similar tendency for electrons (generally to gain electrons). This most
commonly occurs when two nonmetals bond together. Because both of the nonmetals will want to
gain electrons, the elements involved will share electrons in an effort to fill their valence shells. A
good example of a covalent bond is that which occurs between two hydrogen atoms. Atoms of
hydrogen (H) have one valence electron in their first electron shell. Since the capacity of this shell
is two electrons, each hydrogen atom will "want" to pick up a second electron. In an effort to pick
up a second electron, hydrogen atoms will react with nearby hydrogen (H) atoms to form the
compound H2. Because the hydrogen compound is a combination of equally matched atoms, the
atoms will share each other's single electron, forming one covalent bond. In this way, both atoms
share the stability of a full valence shell.
Covalent bonding between hydrogen atoms
Concept simulation - Recreates covalent bonding between hydrogen atoms.
(Flash required)
Unlike ionic compounds, covalent molecules exist as true molecules. Because electrons are
shared in covalent molecules, no full ionic charges are formed. Thus covalent molecules are
not strongly attracted to one another. As a result, covalent molecules move about freely and tend
to exist as liquids or gases at room temperature.
Multiple Bonds: For every pair of electrons shared between two atoms, a single covalent bond is
formed. Some atoms can share multiple pairs of electrons, forming multiple covalent bonds. For
example, oxygen (which has six valence electrons) needs two electrons to complete its valence
shell. When two oxygen atoms form the compound O2, they share two pairs of electrons, forming
two covalent bonds.
Lewis Dot Structures: Lewis dot structures are a shorthand to represent the valence electrons of
an atom. The structures are written as the element symbol surrounded by dots that represent the
valence electrons. The Lewis structures for the elements in the first two periods of the periodic
table are shown below.
Lewis Dot Structures
Lewis structures can also be used to show bonding between atoms. The bonding electrons are
placed between the atoms and can be represented by a pair of dots or a dash (each dash
represents one pair of electrons, or one bond). Lewis structures for H2 and O2 are shown below.
H2
O2
H:H
or
H-H
Polar and Nonpolar Covalent Bonding
There are, in fact, two subtypes of covalent bonds. The H2 molecule is a good example of the first
type of covalent bond, the nonpolar bond. Because both atoms in the H2 molecule have an equal
attraction (or affinity) for electrons, the bonding electrons are equally shared by the two atoms,
and a nonpolar covalent bond is formed. Whenever two atoms of the same element bond
together, a nonpolar bond is formed.
A polar bond is formed when electrons are unequally shared between two atoms. Polar covalent
bonding occurs because one atom has a stronger affinity for electrons than the other (yet not
enough to pull the electrons away completely and form an ion). In a polar covalent bond, the
bonding electrons will spend a greater amount of time around the atom that has the stronger
affinity for electrons. A good example of a polar covalent bond is the hydrogen-oxygen bond in the
water molecule.
Water molecules contain two hydrogen atoms
(pictured in red) bonded to one oxygen atom
(blue). Oxygen, with six valence electrons, needs
two additional electrons to complete its valence
shell. Each hydrogen contains one electron. Thus
oxygen shares the electrons from two hydrogen
atoms to complete its own valence shell, and in
return shares two of its own electrons with each hydrogen,
completing the H valence shells.
Polar covalent bonding simulated in water
The primary difference between the H-O bond in water and the H-H bond is the degree of electron
sharing. The large oxygen atom has a stronger affinity for electrons than the small hydrogen
atoms. Because oxygen has a stronger pull on the bonding electrons, it preoccupies their time,
and this leads to unequal sharing and the formation of a polar covalent bond.
Activity #1 - Use the above reading to anser the following:
As opposed to ____________________ bonding in which a complete transfer of electrons occur,
____________________ bonding occurs when two or more elements __________________ electrons.
Covalent bonding occurs because the atoms in the compound have a similar tendency for electrons
(generally to _____________________ electrons). This most commonly occurs when two
____________________ bond together. Because both of the nonmetals will want to
___________________ electrons, the elements involved will share electrons in an effort to
_______________ their valence shells.
2. How many valence electrons are in one atom of hydrogen?
3. How many valence electrons does hydrogen need to have a full first shell?
4. How does the hydrogen atom “pick up” another electron?
5. What compound does hydrogen form?
6. How do hydrogen atoms make a covalent bond?
8. Label the substances below as “ionic” or “covalent”.
H2O
NaCl
Element Lewis structure Element Lewis structure Element Lewis structure Element Lewis structure
Barium
Sulfur
Silicon
Arsenic
Xenon
Rubidium
Activity #2- Introduction to Lewis Structures
1. Go to http://www.ausetute.com.au/lewisstr.html Fill in the chart.
Number of
valence
Electrons
1
2
3
Example
Family 1
Hydrogen
Family 2
Beryllium
Family 13 Family 14 Family 15
Boron
Carbon
Nitrogen
4
5
6
7
8
Family 16
Oxygen
Family 17
(Halogens)
Fluorine
Family 18 Noble
gases (not He)
Lewis dot
diagram
2. Write the Lewis structure for an atom of each of the following elements. NOTE: You must look
up the group # on the periodic table to do this.
Activity #3 - Naming Covalent Compounds
Open http://dl.clackamas.edu/ch104/lesson8naming_covalentcompounds.html to answer these
questions.
Simple covalent compounds are generally named by using prefixes to indicate how many atoms of each
element are shown in the formula. Also, the ending of the last (most negative) element is changed to -ide.
1. When is the mono- prefix not used to show one atom of an element?
2. When do you drop the “o” and “a” endings of these prefixes?
3. How do you know which element to put first in the name?
4. Name the following compounds.
PH3______________________________
CO______________________________
HI_______________________________
N2O3_____________________________
5. Open Nomenclature. What are the common names of:
H2O______________________________
NH3______________________________
CH4______________________________
For the following use what you have learned to answer the following:
1. Write the formulas for the following covalent compounds.
antimony tribromide __________________________________
hexaboron monosilicide __________________________________
chlorine dioxide __________________________________
hydrogen iodide __________________________________
1
2
3
4
5
6
iodine pentafluoride __________________________________
dinitrogen trioxide __________________________________
ammonia __________________________________
phosphorus triiodide __________________________________
7. Write the names for the following covalent compounds.
P4S5 __________________________________
O2 __________________________________
Si2Br6_________________________________
SeF6__________________________________
SCl4__________________________________
CH4__________________________________
B2Si__________________________________
NF3___________________________________