General Chemistry Chapter 6 Notes
Atoms rarely exist in nature as independent particles. They tend to bond together with other atoms to form
compounds. These compounds are held together by CHEMICAL BONDS. A chemical bond is a mutual electrical
attraction between the nucleus of one atom and the valence electrons of another.
Why does this happen? Independent atoms have relatively high potential energy. Nature favors lower potential energy
and bonding lowers the potential energy to create a more stable compound.
Main group elements tend to either gain or loss electrons to create ions. These ions, cations and anions, are electrically
attracted to one another creating a bond called an IONIC BOND. However, the definition of valence electrons indicates
that electrons can be lost, gained, or shared. The sharing of electrons is
referred to as COVALENT BONDING.
These chemical bonds are rarely purely ionic or covalent, but tend to have
characteristics of both. The degree of ionic to covalent bond can be
determined by the difference between the electronegativities of the
elements involved. If the difference is 1.7 or greater than bond is
considered to be ionic. If the difference is between 0.3 and 1.7, the bond is
considered to be POLAR COVALENT meaning the covalent bond where the
shared electrons are unequally attracted to the atoms. One atom has a
greater electrical attraction to than the other to the electrons.
Electronegativity differences below 0.3 indicate a NONPOLAR-COVALENT
bond. The electrons are shared equally between the bonded atoms. An
easier way to determine the bonds type is to identify the elements involved. Metals bonding to non-metals are IONIC
BONDS. Non-metals bonding to non-metals are COVALENT BONDS.
Covalent bonds make most of the chemicals that are in living things or produced by living things. The covalent bonded
chemicals are referred to as MOLECULES, a neutral group of atoms that are covalently bonded. The compound in its
simplest units is called the MOLECULAR FORMULA. Its composition can be defined by its CHEMICAL FORMULA by
indicating the relative numbers of atoms of each kind by using atomic symbols and numerical subscripts. These
formulas for covalently bonded compounds are often called MOLECULAR FORMULAS.
It is important to remember that bonds occur in order to lower potential energy and achieve stability. This process is
related to the number of valence electrons in a particular atom. Noble gases have a total of 8 valence electrons because
the s and p orbitals of the highest energy level are completely filled. The concept of having or obtaining 8 valence
electrons is known as the OCTET RULE.
Covalent bonds involve only the valence electrons. It is helpful to use ELECTRON-DOT NOTATION to visualize the
valence electrons and how they bond. Electron-dot notation is an electron configuration in which only the valence
electrons are shown as dots surrounding the element symbol.
6th
2nd
3rd
5th
Symbol
7th
1st
4th
8th
The Roman numerals above the
groups on the periodic table
indicate the number of valence
electrons for an element.
The electron-dot notations can
then be used in LEWIS DOT
STRUCTURES to visualize how
covalent bonds are formed. Lewis
dot structures are formulas in
which dot pairs or dashes between
two element symbols represent
shared electrons in covalent bonds and dots adjacent to one of the element symbols represent unshared electrons.
In Lewis dot structures, one pair of shared electrons represents a SINGLE BOND. MULTIPLE BONDS consisting of more
than one pair of shared electrons can exist. These multiple bonds are DOUBLE BONDS (2 shared pairs) and TRIPLE
BONDS (3 shared pairs).
An IONIC COMPOUND is composed of a positive ion (CATION) and a negative ion (ANION) combined so that the
positive and negative charges are equal. The arrangement of the positive and negative ions form an orderly rigid
structure called a CRYSTAL LATTICE.
Comparing the characteristics of ionic versus covalent bonds,
Ionic Bonds
Covalent bonds
Strength
stronger
weaker
Melting points
higher
lower
Boiling points
higher
lower
Physical characteristic
hard brittle solids
many are gases
POLYATOMIC IONS are charged groups of covalently bonded atoms. They combine with atoms or other polyatomic ions
of the opposite charge to form compounds. Compounds containing polyatomic ions will have both covalent and ionic
bonds associated with the compound.
METALLIC BONDING is the type of bonding found in metals. It is different form ionic or covalent bonding. The outer
electron orbitals of metals are occupied by very few electrons. Once metals atoms come close enough the outer orbitals
overlap the electrons are attracted to the nucleus of adjacent metal atoms. This attraction is referred to as the sea of
electrons. As a result, metals are capable of absorbing various frequencies of light and electrons are easily excited, but
immediately fall back to ground state. This characteristic explains the luster associated with metals. The loose
arrangement of electrons allow the atoms to slide past one another making
them malleable and ductile.
INTERMOLCULAR FORCES are the attractive forces between molecules of a
compound. These forces of attraction vary in strength and are responsible of
the states of matter (solids, liquids, gas). The first intermolecular force is
DIPOLE –DIPOLE FORCES. It is the strongest of all intermolecular forces and
involves compounds that are POLAR or having polar covalent bonds. Each
side of a polar compound has a charge associated with it or a dipole.
Opposing dipoles are attracted to each other just as if they are in chemical
bonds. The second intermolecular force is HYDROGEN BONDING. Hydrogen
bonding is when hydrogen is bonded to a highly electronegative atom that
has an unshared pair of electrons. The hydrogen from one molecule has an
attraction to the unshared pair of electrons from another molecule. These
intermolecular forces are found in water. Lastly, LONDON DISPERSION
FORCES are the attraction between molecules because of the temporary
charge implied in an electron orbital as the electron moves in the orbital.
London forces are the weakest of the three types of intermolecular forces,
but the strength does vary as the number of electrons involved increases.
They are also the only intermolecular forces acting upon noble gases and
non-polar compounds.