Chapter 8

• Metallic

• Ionic

• Covalent.

Free electrons hold metals atoms together.

Electrostatic attraction between ions. Sometimes describe as the “transfer of electrons".

Describes: Na(s) + ½ Cl2 → NaCl(s); Ag + Cl^- → AgCl(s)

Sharing of electrons. Attractions of electrons and nuclei (p+). Repulsions between electrons and repulsions between nuclei (p+).

• Brittle

• High melting points

• Crystalline

• Cleave along smooth line.

Lewis symbols contain a number of dots to represent the number of valence electrons.

• Can be drawn to represent covalent bonding (none for ionic bonding).

• All valence electrons are shown; a line is a single bond and represents the sharing of two electrons (also called a “bonding pair”)

Atoms share, lose, or gain electrons in order to have 8 valence electrons to become like noble gases.

Energy releases to make ionic compounds. Related to the strength of an ionic bond.

smaller lattice energy, smaller attractive force.

greater attractive force and lattice energy.

The ability of an atom to pull an electron towards itself.

increase; decrease.

• Nonpolar covalent: 0 - 0.4

• Polar Covalent: 0.5 - 1.9

• Ionic: 2.0+

Electrons are unequally shared.

Electrons are equally shared.

Shows which element pulls under (more electronegative).

• Step 1: Draw skeleton structure “how are the atoms connected”. Usually the least electronegative atom are the center atom (not H).

• Step 2: Count valence electrons.

• Step 3: Add electrons so the total number equals step 2. Add electrons to the outside atoms first.

• Step 4: Check total number of electrons and check octet/duet rule.

4 electrons shared, 2 pairs.

6 electrons shared, 3 pairs.

• Have the highest bond enthalpy, hardest to break, strongest bonds, and shortest bonds.

Pretends all electrons are equally shared.

• All follow the octet rule and have the right amount of electrons. Formal charges will help pick which is best.

• Note: Formal charges are in circles.

Formal Charge = (valence e-) - ½(bonding e-) - (all nonbonding e-)

fewest; most; negative.

Needed when one Lewis Structures doesn’t accurately represent a molecule. Instead, the average of the structures best describes the molecule.

Electrons are moved, same number of electrons, same atom arrangements.

• Odd number of electrons.

• Fewer than eight electrons.

• More than eight electrons.

Second is better because the oxygen is more electronegative.

H and Be can have 2 valence electrons. B can have 6 valence electrons (use formal charges). F will not form more than one bond (no positive formal charges for F).

Elements on row 3 and below can have more than eight electrons because they have access to d orbitals (use formal charges).

Energy needed to break a bond. Takes energy to break bond.