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Bonding

Ionic bond (formula units)

Between metal and a nonmetal

Transfer electrons

Covalent bond (molecules)

Between 2 nonmetals

Share valence electrons

Alloy (metallic “bond”)

Two metals just mix

Don’t chemically bond or react

Properties of Ionic Compounds

Ionic compounds exist as crystalline solids.

A crystal is a regular, repeating, three-dimensional arrangement

of positive and negative ions known as a crystal lattice.

Held together by strong electrostatic forces (opposites attract).

Identifiable properties:

• very high melting points

• hard but brittle (shatters when hammered)

• ions cannot move in the solid state

• when dissolved in water or melted to liquid state, ions

dissociate and form electrolytes

The energy required to separate one mole of the ions of an ionic

compound is called lattice energy.

Characteristics of the Covalent Bond

The forces of attraction are much weaker than

those of ionic bonds.

Molecules melt at low temperatures.

Can not conduct electricity in solution.

Diatomic molecules

Elements with strong electronegativities that bond with

themselves

“Super 7”

H2 , N2 , O2 , F2 , Cl2 , Br2 , I2

“Super 7”

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Ionic Formulas A general rule is to “criss-cross” the absolute value of the charges to

balance out the transfer of electrons

example: aluminum bound to oxygen

1) Write the symbols for the ions.

Al3+ O2–

2) Cross over the numerical value of each ion’s charge (not the charge itself) to the opposite element to form the subscripts Al2 O3 3) Check the combined positive and negative charges to see if they are equal. The overall net charge on an ionic compound is equal to zero.

(2)(+3) + (3)(-2) = 0

Naming Ionic Compounds

(with representative elements)

Binary Compound (only contains a total of two

types of elements)

a) Cation (metal) is identified simply by the element’s

name off the periodic table

b) Anion (non-metal) is named using the root name of the

element with an –ide ending.

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Naming Ionic Compounds

(with representative elements)

Compound containing a polyatomic acid (more

than two types of elements)

a) Cation (metal) is identified simply by the element’s

name off the periodic table

(exception: NH4+1 = ammonium)

b) Anion is named using the name of the polyatomic ion

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a) Cation (metal) is identified simply by the element’s name

off the periodic table

* if the metal can have multiple charges (those metals that

have “d” electrons available to give away), a Roman numeral is

added after the cation’s name to indicate the ion’s

charge

• examples: Fe2+ iron(II)

Fe3+ iron(III)

b) Anion (non-metal) is named using the root name of the

element with an –ide ending or the name of the

polyatomic ion

Naming Ionic Compounds

(with transition metals)

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Acids Compounds that contain Hydrogen ion (H+) when

dissolved in water (aqueous solution = aq)

Two categories:

Binary acids: hydrogen bound to one other element

Named:

a) use the prefix “hydro-”

b) take the root name of the 2nd element and add –ic

ending

c) add the word “acid” to the end

ex. HCl

H2S

HF

Hydrochloric acid

Hydrosulfuric acid

Hydrofluoric acid

Oxyacids: hydrogen bound to a polyatomic ion

NO “hydro” a) name is based on the name of the polyatomic ion

• if it ends in –ate change to –ic ending

• if it ends in –ite change to –ous ending

b) add the word “acid” to the end

ex. HNO2

H2CO3

HClO3

H2SO3

nitrous acid

carbonic acid

chloric acid

sulfurous acid

Covalent Bonds The forces of attraction are much weaker than those of

ionic bonds.

Molecules melt at low temperatures.

Can not conduct electricity in solution.

Some have the ability to share more than one pair of

electrons, forming multiple bonds.

Double bond: two pair of electrons (4 total) are shared

between the two atoms (ex. O2)

Triple bond: three pair of electrons (6 total) are shared

between the two atoms (ex. N2)

Covalent Compounds (molecules)

Bond between 2 non-metals that share their valence electrons

Naming binary covalent compounds (molecules):

a) Name the 1st element in the formula directly off the periodic

table

b) Name the 2nd element using the root name off the periodic

table with an –ide ending

c) Prefixes are used to indicate the number of atoms of each

element that are present in the compound *

1 – mono** 4 – tetra 7 – hepta 9 – nona

2 – di 5 – penta 8 – octa 10 – deca

3 – tri 6 – hexa

* mono- is never used for the 1st element

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The distance between the nuclei of two bonded atoms is

called the bond length.

Energy is released (exothermic) when a bond forms and is

absorbed (endothermic) when a bond breaks

The amount of energy required to break a covalent bond is

called the bond dissociation energy.

The stronger the bond, the greater the bond dissociation

energy and, therefore, the more difficult it is to break the

bond.

Shorter bonds have greater bond dissociation energies than

longer bonds.

Single bonds < double bonds < triple bonds

Lewis Structures: Uses electron dots to show the

arrangement of electrons in a molecule

Steps:

1. Predict the location of the atoms a) Hydrogen is always a terminal atom

b) Center atom is least electronegative element

c) Carbon (if present) is always a center atom

2. Count the total number of valence electrons in the elements to

be combined

3. Determine the number of “pairs” of electrons in the molecule

by dividing the total number of valence electrons by 2

4. Place a single line (“bonding pair”) between the center atom

and the terminal atoms

5. Subtract the number of pairs used from the total number of

pairs available

6. Starting with the terminal atoms, add unshared pairs so that

each atom is surrounded by eight electrons (remember

hydrogen only shares one pair)

7. If the center atom does not have an octet, one or two of the

lone pairs around the terminal atoms must be converted to

form multiple bonds

In general, carbon, nitrogen, oxygen and sulfur can form double or triple

bonds.

Resonance Structure:

• Occurs when more than one valid Lewis Structure can be

written for a molecule or ion.

• Only differ in the position of electron pairs, never in the

atoms position.

ex. O3 (ozone)

NO2-1 (nitrate)

CO3-2 (carbonate)

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Exceptions to the Octet Rule

1) Fewer than eight electrons around the atom (hydrogen and boron

containing compounds such as BH3)

2) Odd number of total valence electrons (These compounds usually

form polyatomic ions to “make-up” the difference)

ex. ClO2

2) Expanded Octets: central atom contains more than 8 electrons

Usually occur with non-metals beyond period 3 when bound to

highly electronegative elements fluorine, oxygen, and chlorine.

ex. SF6

Molecular Geometry (Shape)

VSEPR Theory: “valence-shell electron-pair repulsion.”

The electron pairs are oriented as far away from each other as possible

to minimize the repulsion around the center atom.

The shape of a molecule refers to the positions of atoms only.

5 main shapes (based on the octet rule)

1) Linear

2) Bent

3) Trigonal planar

4) Trigonal pyramid

5) Tetrahedral

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6) Trigonal bipyramid

7) Octahedral

Electronegativity and Polarity

Electronegativity: a measure of the tendency of an atom to

attract electrons in a chemical bond

Polarity: the uneven distribution of electrons (molecule is

asymmetrical around the center atom)

For polar covalent bonds, a dipole is established.

• The forces of attraction within a compound are known as

intramolecular forces. (holds together the atoms making

up a compound)

• Ionic

• Covalent

• Metallic

• The forces of attraction between molecules are known as

intermolecular forces.

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1) London Dispersion Forces

Very weak forces of attraction between non-polar

molecules

Result from the temporary dipole occurring as

molecules approach one another

The more electrons that are present, the stronger the

dispersion forces will be.

Polar molecules have dipoles (partial positive and partial

negative regions.)

2) Dipole-dipole Forces

Occur between polar molecules,

The partial negative region in one molecule attracts the

partial positive region in a neighboring molecules.

There is an electrostatic attraction between the

molecules.

3) Hydrogen Bonding

Hydrogen bound to an atom that has lone pairs of

electrons

The hydrogen atom is attracted to an unshared pair of

electrons on a neighboring molecule.

Are the strongest intermolecular force