This is one of the chapters you must

read….chapter 6…bonding

Student will learn:

1. three types of bonding

ionic, covalent, metallic

2. two categories of bonding

polar, non-polar

3. how to draw Lewis structure

4. how to calculate electronegative

5. bond characteristics

6. VSPER theory

Chemical Bonds ch6 p.161

?What holds chemicals together? Chemical Bonds:

electrical attraction between +nuclei and

–valence electron of different atoms

• By bonding together the atoms are more stable, and have a lower level of energy arrangement

3 types of Bonding:

l. Ionic Bonding: lose or gain –e

metal + Non-metal = ionic bonding

= makes ions

2. Covalent Bonding: share –e

Non-metal + non-metal=covalent bonds

= makes molecules

3. Metallic Bonding: -e flow free in a sea of –e

Transition metals

Pick the bonding

NaCl, CH4, HCl, K2S, FeSO4, LiF, H20, Cu, Zn, Mg(OH)2

2 catagories for the bonding

polar non-polarUnequal equal

attraction for electrons balanced attraction

Ionic bonding (Metals+nonMetals) is always polar

Covalent (nonMetals +nonMetals) maybe either polar/nonpolar

2 ways to figure out

Draw Lewis structure calculate and use chart

Lewis Structure for iodomethane CH3I1. Lewis Dot for each element

C H I2. Arrange to form skeleton

If a carbon then always in middleIf no carbon then least electronegative atom

in middleHydrogen never in middle

See that it is lopsided…… polar covalent molecule

How about individual bonds?

Lewis for ammonia NH3

1. Draw Lewis dot for each element

NH

2. Do skeleton : hydrogen never in middle

Does it look lopsided….

polar covalent molecule

Lewis for formaldehyde CH20l. Lewis dot for each C H O

2. Do skeleton: carbon always in middle

Notice left out –e….move to make a double bond

Notice lopsided: polar……..covalent

Single bonds

Double bonds

Triple bonds

Lewis struture for : CCl4

l. Draw Lewis dot for each: C CL

2. Draw skeleton: carbon always in middle

Does it look lopsided?........No…..

non-polar covalent molecule What about each bond?

Use table on p.151 and chart on page 162

Calculating Polarity of Bonds: using differences in electronegativity

Remember: electronegativity = ability to gain electrons

Bonding is rarely purely ionic or covalent………most of time somewhere in between

Use table on p.151 and chart on page 162 (overhead 31)

Subtract the two electronegativity numbers then ? is it less than 1.7

= polar covalent?

Calculate bond type and polarity

KCl, MgCl2, H2, H2S Cs2S, SCL2,

Comparing Characteristics Ionic bonds (vs) Covalent bonds

Metals + nonmetalsGain or lose electrons so….+ or – ends ……very polarWill form a crystalline lattice,

look on page 177 (model)

Stronger bondsMost are solidsHigher melting pointHigher boiling pointMany dissolve in water, +ion and -ion break apart in water so will conductelectricity in water. Some do not dissolve becausethe pull between the charges are greater thanthe attraction of H2O molecule

Hard but brittle----why?A shift of one row of ions causes a large build up of repulsive forces. And --do not like-- so if one layer moves that forces the other layers to move so they are brittle.

Non-metals + Non-metals

Share electrons

Exist as individual molecules

Weaker bonds

Most are gases, some liquids

Very low melting point

Very low boiling point

Will evaporate at room temperature

Overhead 70

Ionic compounds form

Crystalline lattice

How ionic compounds dissolve

Why Ionic Compounds are brittle

Comparing Characteristics Ionic bonds (vs) Covalent bonds

Metals + nonmetalsGain or lose electrons so….+ or – ends ……very polarWill form a crystalline lattice,

look on page 177 (model)

Stronger bondsMost are solidsHigher melting pointHigher boiling pointMany dissolve in water, +ion and -ion break apart in water so will conductelectricity in water. Some do not dissolve becausethe pull between the charges are greater thanthe attraction of H2O molecule

Hard but brittle----why?A shift of one row of ions causes a large build up of repulsive forces. And --do not like-- so if one layer moves that forces the other layers to move so they are brittle.

Non-metals + Non-metals

Share electrons

Exist as individual molecules

Weaker bonds

Most are gases, some liquids

Very low melting point

Very low boiling point

Will evaporate at room temperature

Overhead 70

Why most covalents are liquids or gases and evaporate easy.

Metallic Bonding p.181

Transitional Metals: vacant outer p orbitals

because filling up d orbitals first.

4s2, 3d10, 4p… they overlap

This overlapping lets –e roam freely about the metal network of empty atomic orbitals.

These mobile –e form a sea of electrons which are packed in a lattice form.

Overhead 68

Characteristics of Metallic bonding “Cu,Au,Ag, Fe”l. Conduct electricity

Conduct heat :::::: due to the “sea of electrons” ability to move freely

2. Reflect light, Shiny, Polish :::::: Contain many orbitals (d10)

separated by extremely small energy differences, metals can absorb a wide range of light frequencies. This absorption of light energy accounts for the ability to reflect light and be shiny. …..p. 181

3. Malleable: hammer into a thin sheet.::::::possible because the metallic bonding is same in all directions throughout the

solid because of “sea of electrons”

Ductile: ability to be drawn into a thin wire.

::::::Because the metallic bonding is same in all directions throughout the solid because of

“sea of electrons”

VSEPR THEORY

Valence Shell Electron Pair Repulsion

Theory : replusion between Valence Shell Electrons Pairs surrounding an atom causes these sets to be oriented as “far apart as possible”.

“AS FAR APART AS POSSIBLE”

Lewis dot, VSEPR TO PREDICT GEOMETRY OF MOLECULE,

Intermolecular force: the attraction between molecules

3 types: dipole-dipole

hydrogen bonding

London dispersion forces

Dipole-dipole: strongest intermolecular force

created when by equal but opposite charges of the molecule come within close distance of each other.

Hydrogen bonding: hydrogen atom that is bonded to a highly electronegative atom is attracted to an unshared pair of electrons of an electronegative atom in a near-by molecule.

“this is why water expands when frozen”

London Dispersion force: results from the constant motion of electrons and the creation of instantaneous dipoles.

Because London dispersion force depends on the motion of electrons, the strength increases with increasing atomic masses.