Chemical Bonding
The attraction between particles is called bonding. The underlying principles of
chemical bonding are (1) the need for all particles to obtain a state of lesser energy and
(2) the attempt of atoms to achieve a noble gas configuration, s2p6 or octet. There can
be bonding between individual atoms and ions (characterized by strong attractions) or
between/among groups of particles (characterized by weak attractions often called van
der Waals forces). The three strong bonds between atoms or ions are the ionic, covalent
and metallic bonds and are formed by gaining, losing or sharing electrons. The three
weak attractions are dipole attractions, hydrogen bonding and London dispersion forces
and are formed from unsymmetrical distribution of electronic charge.
I Overview of the Three Strong Bonds
A. Ionic bonds
1. attraction between a positively charged ion (cation) and a negatively
charged ion (anion) called electrostatic attraction
2. often forms between a metal and nonmetal(s)
3. is characterized by a large electronegativity difference (greater than
1.67)
4. produce solid compounds and have high melting points
5. produce compounds that are brittle
6. produce compounds that lack electrical conductivity in the solid state
7. produce compounds that are often soluble in water
8. water solutions of ionic compounds or the liquid state are electrical
conductors
9. the smallest descriptive unit of an ionically bonded substance is the
unit cell and the formula that represents the ratio of ions is call the
formula unit
B. Covalent bonds
1. the attachment that results from the sharing of electrons between two
or more atoms
2. forms between nonmetals
3. is characterized by small electronegativity differences
4. are termed nonpolar if the electronegativity differences is 0
(some texts define nonpolar as a differences of 0 to 0.2)
5. are termed polar if the electronegativity is greater than 0 but smaller
than 1.67 or between two non-identical nonmetals (some texts
define polar covalent bonds as an electronegativity difference of
between 0.2 and 1.67)
6. produce compounds that have low melting points and/or boiling points
7. produce compounds that are gases, liquids or easily melted solids
8. produce compounds that usually do not dissolve in water
9. produce compounds that are soft and frequently have an odor
10. the same nonmetal atoms may share electrons in more than one way
11. the simplest unit of a covalently bonded system is a molecule, if
neutral, or a polyatomic ion if charged
12. can be double or triples bonds between atoms
C. Metallic bonds
1. exist as system of metal ions in a “sea” of electrons
2. are characterized by delocalized electrons
3. explain the characteristics of metals
a. malleability
b. ductility
c. electrical conductivity in the solid and liquid states
d. thermal conductivity
e. luster
4. produce materials that have generally high melting points
5. produce materials that are not soluble in water
II. Closer Looks at the Strong Bonds—compare and contrast
A. Lewis structures and shapes of the particles that result form bonding
1.
monatomic ions
K∙
→
K+ + e
..
..
: Cl : + e → [ : Cl : ] ¯
˙
˙˙
Dot structure of potassium chloride
K+ ,
..
[ : Cl : ] ¯
˙˙
a. Ions are written as separate particles.
b. There is no “shape” to a formula unit; however repeated
formula units establish a unique crystalline structure or
geometry called a lattice.
2. molecules and polyatomic ions
Dot structure for water
..
H:O:
˙˙
H
a. rules for writing Lewis structures of covalent groups:
i. count the valence electrons; include the charge on
polyatomic ions, PAI
ii. determine the central atom: carbon is usually a central
atom and can form chains; hydrogen can make only
bond and is never a central atom; halogens form
only a single covalent bond when oxygen is not
present and will not generally be a central atom;
oxygen forms only two covalent bonds and is rarely
a central atom but can link two carbon atoms; in
simple molecules, the atom that appears only
once in the formula is the central atom
iii. two electrons, an electron pair, are placed between each
two atoms to form a skeleton structure
iv. remaining electrons are used to complete the octet
structure on the outer atoms of the skeleton
structure.
v. if any electrons remain, they are added to the central
atom
vi. the central atom may have fewer or more than eight
electrons in some situations
b. shapes of molecules and PAI
.
# electron pairs
on central atom
2
3
4
4
4
4
5
5
5
5
6
6
# of unbonded electron
on central atom
0
0
0
1
2
3
0
1
2
3
0
1
shape
_____
linear
trigonal planar
tetrahedral
trigonal pyramid
angular or bent
linear
trigonal bipyramid
seesaw
T-shaped
linear
octahedral
square based pyramid
B. Molecular polarity is the unsymmetrical arrangement of electronic charge in a
molecule. Molecules that do not have a symmetrical arrangement of charge in their
structure are called dipoles.
..
EX: H : O : , uneven distribution of electrons, polar molecule, dipole
˙˙
H
..
..
O : : C : : O , even distribution of electrons, nonpolar molecule, nondipole
˙˙
˙˙
Good sources of additional notes and extension information:
http://www.sciencebyjones.com/chemical_bonding.htm
This is a very good visual powerpoint and is very extensive if you get to
it.
http://www.champaignschools.org/staffwebsites/reidda/1st%20Year%20Chemistry%20Po
werPoints/Ch.%2016%20Covalent%20Bonds%20(teacher)2008.ppt#256,1,Ch. 16 Notes--Covalent Bonds
This is a list of various sites that have information on bonding at all
different levels.
http://river.vansd.org/14544/chemwebtutorial/tutbond.html
This is a site for more info on metallic bonding.
http://www.chemguide.co.uk/atoms/bonding/metallic.html
C. Nomenclature
1. Monatomic ions have a characteristic pattern of nomenclature.
a. Cations of the active metals are named like the atoms from which they
were formed.
EX: a potassium atom, K, forms a potassium ion, K+
Cations of transition metals are also named as the atoms from which
they are made but also carry a Roman numeral in parenthesis that
indicates the charge on the ion (because transition metals may have
more than one oxidation state)
EX: Fe2+, iron (II) ion
Fe3+, iron (III) ion
Often the older patterns of nomenclature are encountered; the ion with
the smaller charge is indicated with an “ous” suffix and the ion with
greater charge is indicated with an “ic” suffix.
EX: Fe2+, ferrous ion
Fe3 , ferric ion
Three ions are exceptions to these rules: Ag+, Cd2+, Zn2+ because they
do not have multiple oxidation numbers and act like the active metals
b. Anions are named with the route of the atom from which it is made
plus an “ide” ending
EX: a chlorine atom, Cl, forms a chloride ion, Cl ¯
2. Polyatomic ions have characteristic names that are usually derived from
the central atom; these names do not always follow a consistent
pattern and must be memorized.
C2H3O2¯ , acetate
HCO3¯, bicarbonate
ClO3¯,, chlorate
CN¯, cyanide
OH¯, hydroxide
MnO4¯, permanganate
NO3¯, nitrate
CO3¯ 2, carbonate
CrO4¯4, chromate
CrO7¯4, dichromate
C2O4¯2, oxalate
O2 ¯2, peroxide
SO4¯ 2, sulfate
BO3¯3, borate
PO4¯3, phosphate
S2O3¯2, thiosulfate
NH4+, ammonium
H3O+, hydronium
3. Ionic compounds are named by identifying each ion in the formula unit.
EX: NaBr, sodium bromide
CuS, copper (II) sulfide
4. Molecules can be named by either the newer method which follows the
pattern of the transition metals and uses a Roman numeral to indicate the
oxidation number of the first nonmetal in the formula or the traditional
method of using numerical prefixes to indicate the number of each atom in
the molecule.
EX: CO, carbon monoxide or carbon (II) oxide
CO2, carbon dioxide or carbon (IV) oxide
P4O10, tetraphosphorus decoxide or phosphorus (V) oxide
Note that the numerical prefix is often eliminated on the first element if it is
“mon” .
Numerical prefixes: mon- = one
hexa- = six
di= two
hepta- = seven
tri= three
octa- = eight
tetra- = four
nona- = nine
penta- = five
deca- = ten
5. Acids are special molecules that generally have their formulas begin with “H-“
or end with “–COOH”. They have a particular set of properties that are
studied in other units but the most important characteristic is that they
produce H+ ions in solution. The extent to which these ions are produced is
indicated by a terminology call pH. The pH scale runs from 0 to 14 and the
lower the number, the more acidic the sample.
a. binary acids ( those containing two elements):
prefix, hydro + root of nonmetal + ic
HCl, hydro- chloric acid
H2S, hydro-
sulfur-
ic
acid
b. ternary acids (those containing three or more elements)
HNO3, (made from the nitrate ion) nitr- ic acid
HNO2, (made from the nitrite ion) nitrous scid
H3PO4, (made from the phosphate ion) phosphor- ic acid
H3PO3, (made from the phosphate ion) phosphor- ous acid
NOTE the variations in the use of the roots sulfur and phosphorus in acid
nomenclature.
6. molecular formulas to recognize:
NH3, ammonia
H2O, water
C6H12O6, glucose or dextrose
CH2COOH, acetic acid (the ending, -COOH, is the organic naming method
for an acid)
CH4, methane
C2H5OH, ethanol
Good source for additional rules and examples of nomenclature.
http://chemed.chem.purdue.edu/genchem/topicreview/bp/ch2/names.html#ionic
Source for organic nomenclature.
http://www.cem.msu.edu/~reusch/VirtualText/nomen1.htm
Shows a stepwise approach to choosing the rules for naming compounds.
http://www2.pvc.maricopa.edu/tutor/chem/chem130/nomenclature/ncrules.html
This is a site of powerpoints on a variety of chemistry topics or which
bonding and nomenclature are part.
http://www.chalkbored.com/lessons/chemistry-11.htm