A. Ionic Bonding1. attraction between large numbers of (+) ions and (-) ions2. results when there is large electronegativity differences3. generally involves a metal and a nonmetal

B. Covalent Bonding1. sharing of electron pairs between two atoms2. results when small electronegativity difference3. generally involves two nonmetal atoms4. nonpolar covalent bond

a. Electrons are shared equallyb. Same two nonmetals or very small

electronegativity difference5. polar covalent bond

a. The more electronegative atom attracts the shared electrons more strongly

b. Results in partial charges on the atoms

C. Greater electronegativity difference (see p.151) results in a more polar bond or greater ionic character

A. Molecule – a neutral group of nonmetal atoms joined together by covalent bonds

B. Molecular Compound1. generally contains only nonmetals2. has covalent bonds3. particles are called molecules4. represented by a molecular formula which gives the actual number of atoms or each element

C. The formation of a covalent bond results from the attraction of the shared electrons by nuclei of both atoms

D. Characteristics of a covalent bond1. bond length – the distance between the nuclei of the two atoms2. bond energy – the energy needed to break a bond

E. Octet Rule1. atoms lose, gain, or share electrons to have 8 in the outer level; (except H=2)2. most compounds follow the octet rule3. exceptions to the octet rule

a. Odd number of valence electrons for the compound

b. BF3 (also some other halides of B or Be)

c. Halogen atoms attached to P or S; more than 8 electrons around the central atom

F. Electron Dot Symbola. Use dots to represent the valence electronsb.

K Mg Al Si

N S Br Ar

G. Lewis Structure (electron dot structures) can be used to show how atoms share electrons in a compound

H. Single, Double, and Triple Covalent Bonds

a. single bond – sharing one pair of electrons

b. double bond – sharing 2 pairs of electrons

c. Triple bond – sharing 3 pairs of electrons

I. Rules for Lewis Structures1. count the total number of valence electrons (write them down)2. determine the central atom (usually the element closest to the center of periodic table)3. draw skeleton structure using only single bonds4. distribute the remaining electrons to have 8 around each atom (except H=2)5. carbon always has 4 bonds (except when bonded to only 1 atom)6. only one bond to F, Cl, Br, I

J. Draw Lewis Structure

water methane

ammonia carbon dioxide

K. Resonance Structures – when more than one Lewis structure can be written for the molecule

1. ozone (O3)

2. sulfur trioxide (SO3)

A. Ionic Compoundsa. Generally contain a metal and a nonmetalb. (+) and (-) ions combined so that charges are equalc. Most form crystalline solidsd. Represented by a formula unit; simplest ratio of the ions

B. Electrons Dot Symbols and Ionic Bonding1. sodium and fluorine

2. magnesium and chlorine

C. Ionic CompoundsMetal and nonmetal

Ionic bonds

High melting points

All are solids

Hard, but brittle

Nonconductors as solids

Conduct when melted or dissolved

Molecular Compounds

Contain only nonmetals

Covalent bonds

Low melting and boiling points

Solids, liquids, and gases

Nonconductors

D. Polyatomic Ion – a group of covalently bonded atoms with a (+) or (-) charge

Draw Lewis Structuresammonium ion sulfate ion

A. Outer electrons of a metal move freely from atom to atom

B. Metallic bonding is the result of the attraction between metal atoms and the surrounding “sea of electrons”

C. Metallic Properties1. good conductors of electricity and heat2. shiny appearance3. malleable – can be hammered into different shapes4. ductile – can be drawn into a wire

D. Higher heat of vaporization results from stronger metallic bonding

A. VSEPR Theory (Valence Shell Electron Pair Repulsion)1. electron pairs repel and are as far apart as possible2. determines the shape of the molecule (molecular geometry)

B. Atoms Attached Unshared Pairs Bond Shape

to Central Atom On Central Atom Angles

4 0 109.5 tetrahedral

3 1 109.5 pyramid

3 0 120 triangle

2 2 109.5 bent

2 1 120 bent

2 0 180 linear

C. Draw Lewis Structure and determine bond angle and shape

methane ammonia

sulfur trioxide water

sulfur dioxide carbon dioxide

D. Hybridization1. s and p atomic orbitals mix to form hybrid orbitals2.

Type of Hybrid Orbital Found if atom has

sp3 only single bonds

sp2 one double bond

sp one triple bond (or two double)

D. Hybridization3. Show type of hybridization for

each carbon atom

E. Polar Molecule (dipole)1. lines up a certain way in an electric field2. molecule with one end slightly positive and the other end slightly negative3.

hydrochloric acid carbon dioxide

water carbon tetrachloride

chloroform

F. Intermolecular Forces (attraction between molecules)1. Stronger attraction between molecules results in higher melting/boiling point; determines whether the substance is solid, liquid or gas2. Types of Intermolecular Forces

a. London Dispersion Forces (weakest)

b. Dipole Forcesc. Hydrogen Bonding (strongest)

G. London Dispersion Forces1. Weakest intermolecular force2. Caused by the motion of electrons; having a majority of electron on one side of molecule causes a temporary unbalanced charge3. All molecules have dispersion forces, but they are important only if it does not have dipole forces or hydrogen bonding4. Strength increases with greater molar mass

H. Dipole Forces1. attraction between polar molecules

2. H – Cl --------H – Cl

I. Hydrogen Bonding1. strongest intermolecular force2. hydrogen attached to N, O, F is also attracted to an unshared pair on a nearby molecule3. water has hydrogen bonding