Chapter 6 Sections 6.1 – 6.4. 6.1 – Chemical Bonds Chemical Bond = A link between atoms Why does...

Chapter 6

Sections 6.1 – 6.4

6.1 – Chemical Bonds

Chemical Bond = A link between atoms

Why does it occur?

The nucleus of one atom is attracted to the electrons of another.

Types of Bonds (an overview)(You will see all of these again later in the chapter!)

1. Ionic BondIon = Atom which has gained or lost

electron(s)Metal =

-LEFT side of Periodic Table-Weak nucleus / Low Electronegativity-LOSERS of electrons-Become + charged ions

Types of Bonds (an overview)

1. Ionic Bond

Nonmetal =

-RIGHT side of Periodic Table

-Strong nucleus / High Electronegativity

-GRABBERS of electrons

-Become - charged ions

Types of Bonds1. Ionic Bonds

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

Examples:

Na =

Na ION =

Cl =

Cl ION =

Types of Bonds

1. Ionic Bonds

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

Usually a metal + a nonmetal

Types of Bonds (an overview)(You will see all of these again later in the chapter!)

2. Covalent Bond = A bond caused by a SHARING of electrons

Usually a nonmetal + a nonmetal

Nonpolar Covalent = Equal sharing of the electrons. Atoms are close in strength

Polar Covalent = Unequal sharing of the electrons. One atom is a little bit stronger than the other World of Chemistry; #8

Chemical Bonds; End at 16:25 – formation of Hydrogen molecule

Types of BondsHow 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.

Types of BondsIf 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:

Types of Bonds (an overview)

3. Metallic bond

Usually metals only

-The metal gives up valence electrons.

-Electrons are free to move about.Atom

Electron Sea

More Detail on the Bond Types6.2 – Covalent Bonds

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

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

atomsMolecular Compound = Compound made of

moleculesMolecular Formula = The type and number

of atoms in a molecule (ex – H2O)

Formation of Covalent Bonds

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

The Octet Rule

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

Drawing Lewis Structures

-A picture of the covalent bonds in a molecule

-Connect valence electrons with LINES

-For Academic classes, atoms follow the octet rule unless stated

Examples:

Single Bond = 1 pair of electrons (2 e-s total) shared between two atoms

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

*Single bonds are the LONGEST in length; Triple are the SHORTEST

*Single bonds have the LOWEST bond energy; Triple have the HIGHEST

Ionic Bond = Bond formed by the attraction of a cation (lost electrons) to an anion (gained electrons)

Crystal Lattice = 3-Dimensional network of ions

Formula Unit = Simplest ratio of ions

More Detail on the Bond Types6.3 – Ionic Bonds

NaCl

Dot structures for Ionic Compounds:

-Will reach noble gas configuration

-Draw an ARROW to show the transfer of e-

-Draw as many of each ion as needed

Examples:

Comparison of Ionic and Molecular

Molecular Compounds IONIC Compounds

Bond Type Covalent Bonds Ionic

Structure Individual Molecules

Crystal Lattice

Strength of Bond

Strong Bonds Very Strong Bonds

mp/bp Low mp / bp High mp / bp

Drawing Connect Dots Arrows with Charges

Other Conduct electricity when melted or in water

Drawing the Pictures – when you’re not told the TYPE of substance

• Do electronegativity difference first!!

• Examples:

6. 4- Metallic Bonding

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

6.4 – Metallic BondingProperties of Metals:1. Good conductors of heat – e- sea shakes

2. Good conductors of electricity – e- in sea can move

3. Malleable – atoms can be pushed closer

4. Ductile – atoms can be pushed closer

5. Luster – light bounces off e- sea

Chapter 6

Section 6.5

Putting Partial Charges on Molecules

Properties of MOLECULAR Compounds

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 (bonds + lone pairs)4. Use chart info

*Academic will be given a WORD BANK and the option to use Model Kits*

VSEPRAreas Bonds Lone

pairsShape Bond

AnglesStructure Example

2 2 0 Linear 180o http://cheminfo.chem.ou.edu/~mra/jmol/jmol.php

BeCl2;CO2

3 3 0 Trigonal Planar 120o http://cheminfo.chem.ou.edu/~mra/jmol/jmol.php

BF3

3 2 1 Bent <120o http://cheminfo.chem.ou.edu/~mra/jmol/jmol.php

SO2

4 4 0 Tetrahedral 109.5o http://cheminfo.chem.ou.edu/~mra/jmol/jmol.php

CH4

4 3 1 TrigonalPyramidal

<109.5o http://cheminfo.chem.ou.edu/~mra/jmol/jmol.php

NH3

4 2 2 Bent <109.5o http://cheminfo.chem.ou.edu/~mra/jmol/jmol.php

H2O

VSEPR

Examples:

VSEPR

Examples:

VSEPR

Examples:

Hybridization

Combination of equal energy orbitals to form new orbitals which all have the same shape and energy

Carbon:

C

BECOMES

C

1s22s22p2 four sp3 hybrid

Types of Molecules

FIRST – Draw Lewis Structure & Include Partial Charges!

1. Dipole = Molecule with overall charge

2. NonPolar With Polar Sites = Molecules with area of charge which cancel out

3. Nonpolar = Molecule with no areas of charge

Types of Molecules

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

1. Is there charge on the molecule?

Yes2. Can it be sliced?

YES = Dipole NO = NPWPS

No = Nonpolar

Examples

Intermolecular ForcesAKA – EXTERNAL BONDS

The attraction BETWEEN Molecules

Types of External Bonds:

1. Dipole-Dipole Interactions

-Occur due to attraction between partial charges

-Occur between two dipoles (Fix notes) – the strongest external bond

Hydrogen Bond = External bond that involves a hydrogen atom

Intermolecular Forces2. London Force

-Occurs between Nonpolar (or Nonpolar With Polar Sites) molecules – CHANGE THIS IN THE NOTES!!

-Very weak connection (nonpolar to nonpolar is the weakest)

The Steps: (only needed for Honors)

A. Electrons in one molecule shift instantaneously to one side

B. Instantaneous charge results

C. Electrons in another molecule are repelled

D. Very weak attraction results

Properties Based on Number / Strength of External Bonds

1. State of Matters>l>g

2. Evaporation (*volatility)slow>fast

3. Thickness (*viscosity)thick>thin

4. Wetness (*adhesion)To feel wet the substance must bond to your skin (to the Na+Cl-)

5. DissolvingLIKE DISSOLVES LIKE

Demonstrations