I. Bond Formation
A. Electronegativity:
relative tendency of an atom to attract electrons to itself when it is
bonded to another atom.
1. Elements are assigned electronegativities based on many experimental
tests.
2. Electronegativities are influenced by the same factors that affect
ionization energies.
3. It follows the same trend as electron affinity and ionization
energy. It increases across a period and up a group.
4. The most active metals (lower left) have lowest electronegativities.
Francium has the lowest electronegativity.
5. Non-metals (upper right) have the highest electronegativities.
Flourine has the highest electronegativity of all
other elements.
B. Bond Strength:
a measure of the energy needed to break the bonds between atoms in molecules
of a compound.
1. The greater the distance in electronegativities, the greater the
bond strength.
2. Example: Arrange the following elements in order of increasing attraction for electrons in a bond (electronegativity).
a. Sb, F, In, Se
(In, Sb, Se, F)
b. Fr, Ga, Ge, P, Zn
(Fr, Zn, Ga, Ge, P)
C. Bond Character
1.
Difference
in electronegativity is high: electrons are transferred between
atoms which form ionic bonds.
2.
Difference
in electronegativity is low: electrons are shared between atoms
forming covalent bonds.
3.
Determining Bond Type
a. First, find the electronegativities of both elements.
b. Then, find the difference between the two electronegativities.
c. Use the chart to determine the % of ionic and covalent character.
d. Unless the two atoms are identical, (F-F), all bonds have some
ionic and covalent characteristics.
e. Example: Determine whether an ionic or covalent bond
will form between the atoms for each of the following pairs:
B-P : |2.01 - 2.06|
= 0.05
covalent
Be-Si : |1.47 - 1.74| =
0.27
covalent
C-Na : |2.50 - 1.01| =
1.49
covalent
Li-O : |0.97 - 3.50|
= 2.53
ionic
Mg-N : |1.23 - 3.07| =
1.84
ionic
D. Ionic Bonds
1. Characteristics of Ionic Compounds
a. transfer of electrons from one atom to another
b. normally a metal and a non-metal
c. high melting points
d. soluble in water (conduct electricity in solution)
e. crystallize as sharply defined particles
2.
Ionic
Bond: electrostatic force that holds two ions together due to
their differing charges.
Example: NaCl
E. Covalent Bonds
1.
Characteristics of Covalent Compounds
a. Electrons are shared between atoms
b. Most are formed between nonmetals.
c. low melting points
d. They are brittle.
e. don't conduct electricity well
2.
Example: Cl2
F. Metallic Bonds
1.
Characteristics of Metallic Bonds
a. When metals bond with one another, they don't share or transfer
electrons.
b. Metal crystals form when atoms crowd together and the outer level
orbitals from all those atoms overlap.
c. The electrons can move more easily from one atom to the next.
Thus the are called delocalized electrons.
d. Metallic bonds are extremely strong.
2. Alloys: metallic materials that consist of 2 or more
elements (usually metals).
a. Alloys are not true metals.
b. They are solid solutions.
c. Examples:
Brass - alloy of copper and zinc
Steel - alloy of iron and carbon
Bronze - alloy of copper and tin
G. Polyatomic Ions
-- groups of ions bonded covalently but they possess an overall charge
just as other ions.
II. Particle Size
A. Ionic Radii
1. internuclear distance: the sum of the radii of two
ions in a compound.
2. ionic radius: distance from the nucleus to the outermost
electron in an ion.
B. Covalent Radii
1. bond length: the sum of the covalent radii of both
atoms forming a covalent bond.
2. Predicting Bond Length
example: Predict the length of the bond formed between an atom
of arsenic and an atom of sulfur.
You must use a table of values for bond lengths.
As - covalent radius 120 pm (picometers)
S - covalent radius
103 pm
223 pm