Bonding

Bonding

I.
Introduction

A. Why?

Atoms are more stable when their valence is complete
(stable octet = 8). They achieve this through bonding.

As bonding occurs the net (overall) energy decreases because some of it is stored in the bond. This increases the stability.

B. Vocabulary
1. Ionization Energy: The measure of the ease in which an atom loses an electron to form a positive ion.
	- The less energy needed, the more easily the electron is lost - Called 1st ionization energy because it is the energy required to remove the 1st electron. - Listed in Reference Table S - Measured in KJ/mol
2. Electron Affinity: The ease with which atoms gain electrons gain electrons to form negative ions.
3. Isoelectronic: When the ion(s) of 2 different elements have the same electron configuration Example: K+ and Ar
4. Electronegativity: Measures the relative attraction of an atom for shared electrons. Listed in Reference Table S
Electronegative difference (the difference between the electronegativities of the two atoms in a bond) can be used to indicate bond type.
x >1.7 = Ionic Bond 1.7 > x < 0.4 = Polar Covalent x < 0.4 = Non-polar Covalent

II.
Bond Types

A. Ionic: Ions are TRANSFERRED between atoms creating a pair of oppositely charged ions. These ions are attracted to form an ion. · Created by the Transfer of electrons from a Metal to a Nonmetal · Strong bond · Ionic substances conduct electricity in solution (aq) · Called salts or crystals
	[K]+ [Cl]-
In order for this to happen one atom must readily lose an electron (metal) and one must readily gain (nonmetal).
B. Covalent: atoms SHARE pairs of electrons (can be 1, 2 or 3 pairs). These electrons belong to both atoms. · Created by the Sharing of electrons between two Nonmetals · Typically weak bonds (except Networks) · Covalent substances don’t conduct electricity · Called molecular substances
1. Non-polar -Evenly Shared electrons -Diatomic molecules -Example F2:
	2. Polar -Unevenly Shared electrons -Example HCl
3. Coordinate -One member of the bond brings both electrons (XX), the other has none. -H+ needs to do this to bond -Look for a lone pair (XX) -Example NH4:
4. Network -Strongest Bond -3-Dimensional solid, no individual molecules -Diamonds (C); Silicone Dioxide (SiO2), Silicone Carbide (SiC)

5. Multiple (double or triple bonds) -Two elements share more than 1 bond
-Double (2 bonds) Carbon compounds or Oxygen molecule (O=O)	-Triple (3 bonds) Nitrogen Molecule (N=N)

C. Metallic · Solid pure metals or alloys · Conduct electricity in Solid or Liquid state · Low number of valence electrons DELOCALIZE and move over entire crystal not just a single atom “Sea of Mobile Electrons”

III.
Intermolecular Attractions:
an attraction between molecules

A. Dipole Dipole Attractions
The side of the molecule that has a greater electronegativity will have more electrons. This can result in a momentary negative charge on that side.		The side of the molecule that has a lower electronegativity will have less electrons. This can result in a momentary positve charge on that side.
	These charges are called induced dipoles

In a polar covalent molecule the + and – ends attract the oppositely charged end of the other molecule. The bigger the atom the stronger the attraction. These attractions are responsible for the liquid and solid phase. High temperature can overcome these attractions and gas phase can be reached.

B. Hydrogen Bonding
	Hydrogen bonding is actually an intermolecular attraction that forms between the hydrogen of one molecule and a small, highly electronegative element in an adjacent molecule. Typically this is Fluorine (HF), Oxygen (H20) or Nitrogen (NH3).

C. London Dispersion Forces (a.k.a. van der Waals Forces)
These are relatively weak forces of attraction that exist between nonpolar molecules. The larger the molecule the stronger the attraction. The attraction will increase as the distance between the molecules decreases.

D. Molecule-Ion Attractions
An ionic salt breaks down to provide the ions in a solution (aq) and the water is the molecule. In a Molecule-Ion attraction the positive ion in the salt is attracted by the negative dipole of the water molecule.		We also see that the negative ion is attracted by the positive dipole of the water molecule. The water separates the ions, destroying the crystal structure. Salt (aq)

IV.
Polar vs. Nonpolar MOLECULES
(different than polar and nonpolar bond)

A. Symmetrical charge distribution is nonpolar molecule Examples CO2, CH4 and diatomics

B. Asymmetrical charge distribution is a polar molecule Examples HCl, NH3 and H2O