COVALENT COMPOUNDS

Formed when two atoms share electrons

o Usually involve 2 or more

non-metal atoms

o Similar electronegativity

values (electronegativity

difference is less than 1.7).

Can be pure covalent –

equal sharing of e (<.3)

Polar covalent – uneven

sharing of electrons (.3-

1.7)

o Called molecules

o There are more covalent

compounds than ionic

compounds.

Binary covalent – include 2

elements only

Organic molecules (which

include carbon) are

covalent

Glucose C6H12O6

Sucrose C12H22O11

Hydrocarbon C6H6 etc.

The shared valence electrons

are found in a molecular orbital, formed by overlapping

atomic orbitals of the atoms

involved in bonding.

Recall the properties of

covalent molecules listed

previously in chapter 5 notes

(flow chart).

NAMING COVALENT COMPOUNDS

Covalent molecules are NOT

named like ionic compounds!!!

Steps to naming binary covalent

molecules: (only include 2

elements)

1. the first element (least

electronegative) is named first

2. the second element has the

ending –ide

3. Prefixes are used before the

element name to indicate the

number of each atom in the

molecule (the subscript)

# of atoms Prefix1 Mono

2 Di3 Tri4 Tetra5 Penta6 Hexa7 Septa8 Octa

9 Nona10 Deca

Do NOT use a prefix for the first

atom when there is only one

present.

When the element name begins

with a vowel, the (a) and (o) are

dropped from the end of the

prefix

Examples:

CCl4

carbon tetrachloride

CO

carbon monoxide

(eliminate the o from the prefix)

P2Cl5

diphosphorus pentachloride

(more than one of the first

element)

N2O3

dinitrogen trioxide

Si3N4

trisilicon tetranitride

H2O

dihydrogen monoxide

Steps for writing covalent

formulas from the name.

1. the elements appear in the

same order as in the name

2. the prefix indicates the

subscript in the chemical

formula

Examples:

Boron trifluoride BF3

Dinitrogen monoxide N2O

Dinitrogen tetroxide N2O4

ENERGY and STABILITY of COVALENT BONDS

Most atoms have relatively LOW STABILITY and HIGH POTENTIAL ENERGY.

When a compound forms, the atoms become MORE STABLE and the potential energy is at a MINIMUM.

Example:

The potential energy curve for hydrogen.

http://www.usm.maine.edu/~newton/Chy251_253/Lectures/LewisStructures/Dihydrogen.html

When the nuclei are farthest apart, the potential energy is ZERO.

As they get closer, the energy DECREASES.

When the potential energy is LOWEST (at -436 kJ//mole), the atoms bond.

The distance between the two nuclei at this lowest energy is called the BOND LENGTH and is 75 pm for the H2 molecule.

When the REPULSION of the two atoms perfectly balance the ATTRACTIVE FORCES between the two nuclei, a COVALENT BOND forms.

Since the potential energy DECREASED, energy has been RELEASED. For H2, the energy released is -436 kj/mol.

The energy released when forming the bond is the exact amount of energy that would be needed to break the bond. Energy needed = +436 kj/mol

The energy required to break a bond is known as the BOND ENERGY.

BOND ENERGY and

BOND LENGTH

Bond energ

Bond lengt

Electronegativity Difference

y(kj/

mol)

h (pm)

HF 570 92 1.8CF 552 138OO 498 121HH 436 75HC

l432 127 1.0

CCl

397 177

HBr

366 141 0.8

HI 299 161 0.5

BOND ENERGY (STRENGTH) is inversely related to the BOND LENGTH. The higher the bond energy, the SHORTER the bond.

Bond length is actually an AVERAGE DISTANCE between the two nuclei since the distance is constantly changing due to the bond being able to vibrate and bend.

BOND STRENGTH and POLARITY

The HIGHER the electronegativity difference, the STRONGER the bond.

o Ex. HF has a much stronger bond than HI.

o HF is nearly IONIC and HI is almost PURE COVALENT.

LEWIS DOT STRUCTURES

RESONANCE

when 2 or more equivalent lewis structures exist for a

molecule or compound

all structures should be represented with a double

arrow between

the actual bond strength is an average of all the

bonds.

Examples:

(see back of notes)

MOLECULARGEOMETRY: ARRANGEMENT and

SHAPE

The three dimensional shape of molecules can be predicted using the

VALENCE SHELL ELCTRON PAIR REPULSION THEORY.

omaximizes space between electron pairs to predict shape.

oVSEPR

Arrangement of electrons – based on the number of regions of

electron density around the central atom.

Shape of the molecule – can be predicted by counting the number of

bonding pair electrons and lone pair electrons.

Treat double and triple bonds as single electron pairs when

determining VSEPR shape.

once the shape is determined, certain bond ANGLES can be

predicted. Bond angles are altered by the present of LONE PAIR

ELECTRONS. When present, the bond angles are compressed due

to the REPULSION of the.

o lone pairs require more space since they are only controlled

by one nucleus.

The arrangement of electrons and the shape of the molecule with be

the same when there are NO LONE PAIR ELECTRONS.

VSEPR and MOLECULAR GEOMETRY

BONDING PAIRS of ELECTRONS

LONE PAIR ELECTRONS

TOTAL NUMBER OF ELECTRON PAIRS

ARRANGEMENT

SHAPE SKETCH and BOND ANGLES

1 0 1

2 0 2

3 0 3

2 1 3

4 0 4

3 1 4

2 2 4

5 0 5

4 1 5

3 2 5

2 3 5

BONDING PAIRS of ELECTRONS

LONE PAIR ELECTRONS

TOTAL NUMBER OF ELECTRON PAIRS

ARRANGEMENT

SHAPE SKETCH and BOND ANGLES

6 0 6

5 1 6

4 2 6

3 3 6

2 4 6

HYBRIDIZATIONIn order for bonds to have equivalent energy, mixing of orbitals must occur.

Ex. Methane (CH4) predicts 4 equivalent bonds. Which orbitals of C are used?

mix the orbitals used into a new “hybrid orbital” with a new name.

Total number of e pairs around a

central atom

Arrangement Hybridization

2 Linear sp3 Trigonal

planarsp2

4 Tetrahedral sp3

5 Trigonal bipyramidal

sp3d

6 octahedral sp3d2

POLARITY and DIPOLE MOMENT

POLAR BONDS:

Result from high ELECTRONEGATIVITY DIFFERENCES between atoms of a molecule.

Causes PARTIAL POSITIVE and PARTIAL NEGATIVE ends of the molecule (called DIPOLES).

SHAPE AFFECTS POLARITY:

DIPOLE MOMENT – overall direction of electron “pull” within a molecule. Show using a molecular model. Sometimes dipoles will CANCEL each other, and the result will be a molecule with NO NET DIPOLE due to the shape.

Molecul Sketch Direction of

e dipole or no net dipole

HBrBeCl2BBr3

SeCl2CO2

H2O

POLARITY AFFECTS PROPERTIES:

For example: o CO2 is non-polar, so the ATTRACTIVE FORCE

between CO2 atoms is VERY WEAK.

o This results in a lower MELTING POINT and BOILING POINT.

H2O is POLAR, the molecules interact with each other and attractive forces are greater.

This results in a higher MELTING POINT and BOILING POINT.

A very strong INTERMOLECULAR force exists between water molecules due to its polarity.

o This force is called HYDROGEN BONDING

It is not the bond within water….it is a bond between water molecules.

HYBRIDIZATION