Purpose of Experiment 12
become familiar with:
1. Lewis structures
2. principles of the VSEPR model
3. the three-dimensional structures of covalent molecules
chemical bond
when atoms or ions are strongly attached to one another
Three types of chemical bonds
1. ionic
2. covalent
3. metallic
Ionic bond
- electrostatic forces btwn ions of opposite charge
- formed by transfer of one or more electrons
- from atom of low ionization energy to atom of high ionization energy
usu metals with nonmetals (except noble gases)
Covalent bond
sharing of electrons between two atoms
More familiar exs= bonds among nonmetallic elements
Metallic bonds
found in metals like gold, iron, magnesium
each atom is bonded to several neighboring atoms
give rise to high electrical and thermal conductivity
b/c in m bonds electrons are relatively free to move throughout the 3D shapes of mc's
Which electrons are involved in bonding
valence electrons (reside in incomplete outer shell of an atom)
creator of Lewis Dot structure
G.N. Lewis, American chemist
Lewis symbol consists of:
chemical abbreviation for the element
and a dot for each valence electron
each side accommodates up to two electrons
What can atoms do to achieve noble gas config?
gain, lose, or share electrons
octet rule
atoms tend to gain, lose, or share electrons until they are surrounded by eight valence electrons
octet of electrons= full s and p subshells
Properties of ionic substances
(most substances do not have these characteristics)
usually brittle
high melting point
crystalline solids w/ well-formed faces
- characteristics result from electrostatic forces that maintain ions in a rigid, well-defined, 3D arrangement
Example of how H2 forms a bond
- the nuclei and electrons cause electron density to concentrate btwn the nuclei
- the shared pair of electrons acts as a kind of "glue" holding the atoms together
single bond
double bond
sharing of one pair of electrons
sharing of two pairs of electrons
- triple bond is the sharing of three pairs of electrons
Steps for drawing Lewis dot structures
1. Sum the number of valence e-
a. add electrons for negative charges
b. subtract electrons for positive charges
2. Decide the central atom (least electronegative excluding H)
3. Assign leftover electrons to the terminal atoms until all octet
4. Assign any leftover e- to the central atom (if central atom doesn't have octet, create multiple bonds)
Why formal charge?
sometimes can draw several diff Lewis structures that all obey the octet rule
formal charge helps decide the structure that is the most reasonable
formal charge (of an atom)
the charge that an atom in a molecule or ion would have if all atoms had the same elcectronegativity
equation for formal charge
(regular valence e-) - (valence e- found in the molecule)
* remember to count only one of the two bonded molecules for each bonded atom
Most bonds btwn two diff kinds of atoms are usually...
polar
Why are most bonds of two diff atoms usu polar?
b/c diff atoms have diff electronegativities (so electrons are not shared equally by the two bonded atoms)