Sections 6.1 – 6.4

Chemical Bond = A link between atoms

Why does it occur?The nucleus of one atom is attracted to the electrons of another.

IONIC BONDIon = Atom which has gained or lost electron(s)Metal =

-LEFT side of Periodic TableWeak nucleus / Low Electronegativity

-LOSERS of electrons Become positively (+) charged ions

(Cations)

Nonmetal =-RIGHT side of Periodic Table Strong nucleus / High Electronegativity-GRABBERS of electrons Become negatively (-) charged ions

(Anions)

Atoms gain or lose valence electrons to become a NOBLE GAS CONFIGURATION

Right/Left Side?

Metal/ Nonmetal?

Lose/Gain? Noble Gas it resembles?

Mg

Li

Cl

O

Ionic bond = A chemical bond between a cation (+) and an anion (-). Caused by a TRANSFER of electron(s).

Usually a metal + a nonmetal

COVALENT BOND = A bond caused by a SHARING of electrons

Usually a nonmetal + a nonmetalNonpolar Covalent = Equal sharing of the

electrons. Atoms are close in strengthPolar Covalent = Unequal sharing of the

electrons. One atom is a little bit stronger than the other

How do you tell which type of bond it is?-By ELECTRONEGATIVITY

A chart of electronegativity will be provided to you.

-The greater the difference in electronegativity – the more ionic the bond.

-Electrons spend more time closer to the element with higher electronegativity.

If the ABSOLUTE VALUE of the electronegativity difference is:

GREATER THAN 1.7 = IONIC Bond LESS THAN 0.3 = NONPOLAR COVALENT

Bond 0.3 – 1.7 = POLAR COVALENT Bond

Examples:

METALLIC BONDUsually metals only-The metal gives up valence electrons.-Electrons are free to move about.

Atom

Electron Sea

Covalent Bond = A sharing of electronsMolecule = A group of atoms held by

covalent bonds (ex – water)Diatomic Molecule = Molecule with only 2

atoms (7 naturally occurring ones)Molecular Compound = Compound made of

moleculesMolecular Formula = The type and number of

atoms in a molecule (ex – H2O)

Sharing electrons in a covalent bond makes the atoms more stable and decreases the energy of the atoms. Energy is released when a bond is FORMED.

Overlapping of Orbitals – Example H2:

H H

+

H2

Atoms in a compound obtain the electron configuration of a NOBLE GAS to gain stability

Exceptions to the Octet Rule: Hydrogen Incomplete octets Expanded octets

Surrounded by only 2 electrons◦ Can only form single bonds!!

Too few valence electrons Rather rare:

◦ Boron, aluminum, beryllium Examples:

Some central atoms form 5 or 6 bonds Elements with atomic #s greater than 10 Examples:

-A picture of the covalent bonds in a molecule

-Basic Rules:◦ Determine a central atom

If C is present – always central H is never central Group 17 is never central

◦ Arrange to form skeleton (like a plus sign)◦ Place dots around each element and connect

dots

Examples:

More Examples:

Single Bond = 1 pair of electrons (2 e-s total) shared between two atoms (longest length; lowest bond energy)

Double Bond = 2 pairs of electrons (4 total e-s) shared between two atoms

Triple Bond = 3 pairs of electrons (6 total e-s) shared between two atoms (shortest length; highest bond energy)

More Examples:

Same as non-charged molecules, except:◦ For positive ions – subtract electrons to total # of

valence electrons◦ For negative ions – add electrons to total # of

valence electrons Examples:

Ionic Bond = Bond formed by the attraction of a cation to an anion

Crystal Lattice = 3-Dimensional network of ions

Formula Unit = Simplest ratio of ions

NaCl

Dot structures for Ionic Compounds:-Want to reach noble gas configuration“Math equation”-Draw an ARROW to show the transfer of e--Draw as many of each ion as neededExamples:

Molecular Ionic

Bond Type Covalent Ionic

Structure Individual Molecules Crystal Lattice

Strength of Bond Strong VERY strong

Mp/bp Low High

Drawing Lewis Structures “Math Eqn”

Other --- Conducts electricity when in water

Do electronegativity difference first!! Examples:

Metals have LOW electronegativity – Will LOSE electrons

The steps:-Donate valence electrons to electron sea-Electrons free to move about-All electrons in sea are shared by all atoms

Properties of Metals:1. Good conductors of heat – e- sea shakes2. Good conductors of electricity – e- in sea

can move3. Malleable – atoms can be pushed closer4. Ductile – atoms can be pushed closer5. Luster – light bounces off e- sea

Section 6.5

VSEPR = Valence Shell Electron Pair Repulsion Theory

Valence electrons move as far away from each other as possible

1. Draw Lewis Structure2. Look at Central Atom3. Count electron areas (shared areas & lone

pairs)4. Use chart info

Name Shape Shared Areas

Lone Pairs Bond Angles

Linear 2 0 180°

Bent 2 1 120°

Trigonal Planar

3 0 120°

Tetrahedral 4 0 109.5°

Trigonal pyramidal

3 1 107°

Bent 2 2 104.5°

Trigonal bipyramidal

5 0 90°, 120°

Octahedral 6 0 90°

Examples:

Examples:

Examples:

Additional Handout!! Not in note packet!! Partial Charges

◦ In a polar bond ONLY!!◦ A tug of war occurs – one atom is “stronger”

than the other! δ+ and δ-

◦ Greek letter delta

◦ Compare EN: Higher EN value of the two = δ-

Lower EN value = δ+

Examples:◦ H2O

◦ CH4

◦ NH3

◦ CO2

◦ And the hardest – CH3Cl

Mixing a set of atomic orbitals◦ forms a new set of atomic orbitals with the same total electron capacity◦ properties and energies intermediate between those of the original

unhybridized orbitals.

Three types: sp (triple bonds), sp2(double bonds), sp3(single bonds)

Carbon:

C

BECOMES

C

1s22s22p2 four sp3 hybrid

1. DipoleDipole = Molecule with overall charge2. NonPolar With Polar Sites (NPWPS)NonPolar With Polar Sites (NPWPS) =

Molecules with area of charge which cancel out

3. Nonpolar Nonpolar = Molecule with no areas of charge

How do you tell the difference?-Ask yourself these questions…

Is the molecule polar or nonpolar (Difference in EN)

Polar (∆EN = 0.3-1.7) Can charge be sliced?

YES = Dipole NO = NPWPS

Nonpolar (∆EN < 0.3)

One straight line so all positive charge is separated from all

negative charge; through BONDS only!!

AKA – EXTERNAL BONDSThe attraction BETWEENBETWEEN MoleculesTypes of External Bonds:1.1. Dipole-Dipole InteractionsDipole-Dipole Interactions

-Occur due to attraction between partial charges-Occur between:

• Two dipoles (strongest)• Dipole to NPWPS• Two NPWPS (weakest)

Hydrogen Bond = Special case of a dipole-dipole external bond that involves a hydrogen atom

2.2. London ForceLondon Force-Occurs between nonpolar molecules-Very weak connectionThe Steps:

A. Electrons in one molecule shift instantaneously to one side

B. Instantaneous charge resultsC. Electrons in another molecule are repelledD. Very weak attraction results

1. State of Matters > l > g

**This means that solids have strongest external bonds; gases have weakest bonds

2. Evaporation (Volatility – tendency of a substance to vaporize)slow > fast

**Those compounds that evaporate very slow have stronger bonds than those that evaporate quickly

3. Thickness (Viscosity – a measurement of resistance to flow)thick > thin

**Substances that are “thicker” have stronger bonds

4. Wetness (Adhesion – the force of attraction)To feel wet the substance must bond to your skin (to the Na+Cl-

**If you feel wetness, the substance is bonding to your skin

5. DissolvingLIKE DISSOLVES LIKE

**Polar dissolves in polar; nonpolar dissolves in nonpolar