Covalent Bonding and Naming

I. Types of Covalent Bonds

• l. Nonpolar covalent bond-a covalent bond in which the bonding electrons are shared equally

• 2. Polar covalent bond-a bond is which the bonded atoms do NOT share the bonding electrons equally.

Types of Covalent Bonds

• If the difference in electronegativity between two bonded atoms is from 0.3 to 1.7, a polar bond will exist

• If the difference in EN is less than 0.3 then the bond is nonpolar covalent.

Practice

a. H and Cl

b. F and Br

c. S and I

d. O and H

The Octet Rule

• chemical compounds tend to form so that each atom, by gaining, losing, or sharing electrons, has an octet of electrons in its highest energy level.

Multiple Bonding

• single bond -two atoms share one pair of electrons

• double bond -two atoms share two pair of electrons

• triple bond -two atoms share three pair of electrons

Drawing Lewis Dot Structures

Add up the TOTAL number of valence electrons in the substance

Decide what is the central atom. The central atom is the one that is least represented. (or the least electronegative)

Arrange the dots so that each atom has an octet

Practice 1. Br2

2. CCl4

3. H2O

4. NH3

5. O2

6. N2

Exceptions to the Octet Rule

• Some atoms have less than an octet. Example: Hydrogen only needs 2 electrons surrounding it and boron only needs 6.

• H2

• BF3

Exceptions to the Octet Rule

Some atoms have more than an octet (One reason is because of bonding d orbitals as well as s and p orbitals.) Example: Sulfur can have up to 12 electrons surrounding it.

SF6

Resonance

• – a concept in which two or more Lewis structures for the same arrangement of atoms (resonance structures) are used to describe the bonding in a molecule or ion. To show resonance, a double-headed arrow is placed between a molecule’s resonance structures. Example: O3

Molecular ShapesThe Valence Shell Electron Pair Repulsion

Theory (VSEPR) is used to predict the shape of molecules.

The VSEPR theory states that the bonding and nonbonding pairs of electrons will arrange themselves so that the repulsive forces between them are at a minimum. Molecules will adjust their shapes so that the nonbonding electrons are as far apart as possible.

Molecular Shapes

• The number of shared and unshared pairs of electrons determines the shape of the molecule.

Linear

• The bond angle is 180º .

• 2 atoms bonded to the central atom

• 0 lone pairs

• Ex: CO2

Bent

The bond angle is 90º

2 atoms bonded to the central atom

2 lone pairs

Write the Lewis structure for SF2

Trigonal Planar

The bond angle is 120°

3 atoms bonded to the central atom

0 lone pairs

Write the Lewis structure for BCl3

Trigonal Pyramidal

The bond angle is 107°

3 atoms bonded to the central atom

1 lone pairs

Write the Lewis structure for NH3

Tetrahedral

The bond angle is 109.5°

4 atoms bonded to the central atom

0 lone pairs

Write the Lewis structure for CH4

Trigonal bipyramidal

The bond angle is 120°

5 atoms bonded to the central atom

0 lone pairs

Write the Lewis structure for PCl5

Octahedral

The bond angle is 120°

6 atoms bonded to the central atom

0 lone pairs

Write the Lewis structure for SF6

Polarity

A. nonpolar covalent bonds – EN difference is less than 0.3

There is equal sharing of electrons

B. polar covalent bonds – EN is between 0.3 and 1.7

There is unequal sharing of electrons

Practice

• Determine whether each of the following is a polar covalent bond, nonpolar covalent, or ionic bond.

• 1. N-O bond 2. Cl-Cl bond

• 3. C-S bond4. S-O bond

• 5. P-O bond6. Na-Cl bond

Determining Polarity In a polar molecule there is an uneven

distribution of charge. One end of the molecule is more negative or positive than the other

In a nonpolar molecule there is an even distribution of charge.

Determining Polarity

Molecules with nonbonding pairs of electrons on the central atom are polar.

H2O

NH3

Determining Polarity• The polarity of molecules with no

nonbonding pairs depends on the atoms bound to the central atom.

• If the surrounding atoms are identical, the molecule is nonpolar. This is because the bond dipoles cancel.

• Example: CH4

• When the atoms surrounding the central atom are different, the molecule is polar.

• Example: BBrI2

Practice

SO2

H2S

CO2

Intermolecular Forces of Attraction (IMFs)

Van der Waals forces The attractive forces between molecules are

collectively referred to as van der Waals forces.1. There are very strong intermolecular forces of attraction between polar molecules and weak intermolecular forces of attraction between nonpolar molecules.2. Many physical properties of a substance, such as freezing point and melting point, depend on the strength of the intermolecular forces of attraction.

Dipole-Dipole Forces

• 1. Dipole-dipole forces exist between polar molecules. The negative dipole of one molecule attracts the positive dipole of another molecule.

• Example: HCl

• 2. Dipole-dipole forces are weaker than chemical bonds.

Hydrogen Bonding• 1. Hydrogen bonding is a special type of

dipole-dipole force. Since no electrons are shared or transferred, hydrogen bonding is not a chemical bond.

• 2. Hydrogen bonding always involves molecules containing hydrogen that is chemically bonded to a highly electronegative atom of small atomic size, specifically nitrogen, oxygen or fluorine Example: H2O

London Dispersion Forces• 1. London dispersion forces are most noticeable

in nonpolar molecules• 2. London dispersion forces arise from the

motion of electron clouds . From the probability distributions of orbitals, it is concluded that the electrons are evenly distributed around the nucleus. However, at any one instant, the electron cloud may become distorted as the electrons shift to an unequal distribution. It is during this instant that a molecule develops a temporary dipole.This temporary dipole introduces a similar response in neighboring molecules, thus producing a short-lived attraction between molecules

Strength of Forces of Attraction from Weakest to Strongest

• 1. Chemical Bonds– a. Nonpolar Covalent– b. Polar Covalent– c. Ionic

• 2. Van der Waals Forces– a. London Dispersion Forces – b. Dipole-Dipole Forces– c. Hydrogen Bonding

Comparing Boiling Points• 1. When molar masses are similar, substances

with stronger intermolecular forces of attraction have higher boiling and freezing points.

• 2. In order to compare boiling and freezing points of a substances with similar molar masses you must:

• a. Identify the structural formula• b. Determine the polarity of the molecule.• c. Determine the dominant type of Van der

Waals forces• 3. Which of the following molecules has the

higher boiling point: CH4 or NH3? Explain your answer.

Naming and Writing Covalent Compounds

• Writing Formulas for Binary Molecular Compounds those containing 2 nonmetals. Use the prefix naming system - know theses prefixes

Naming and Writing Covalent Compounds

• mono – one • di – two• tri – three• tetra – four• penta - five• hexa – six• hepta – seven• octa - eight• nona - nine• deca – ten

Practice • nitrogen tetrasulfide

• carbon dioxide

• oxygen monofluoride

• sulfur hexachloride

• trioxygen decanitride

• tetrafluorine monophosphide

• hexafluorine nonasulfide

• heptabromine octanitride

Naming Binary Molecular Compounds

• The less electronegative element is given first.

• The second element is named by combining a prefix indicating the number of atoms contributed by the element to the root of the name of the second element and then adding –ide to the end.

Practice • CCl4 _________________________

• NF3 _________________________

• PBr5 _________________________

SF6 _________________________

• SO3 _________________________

• PCl5 _________________________

• N2O _________________________ PF6

_________________________

Naming Acids

Binary or hydrohalic acids

Name “hydro____ic acid”

Ex: HF ____________________

HCl ____________________

HBr ____________________

HI ____________________

H2S ____________________

Oxyacids – contain a polyatomic ion If the polyatomic ion ends “ate” the acid will

end in “-ic”

HNO3 ____________________

H3PO4 ____________________

H2SO4 ____________________

H2CO3 ____________________

HC2H3O2 ____________________

Oxyacids – contain a polyatomic ion

If the polyatomic ion ends “ite” the acid will end in “-ous”

HNO2 ____________________

H3PO3 ____________________

H2SO3 ____________________

Metallic Bonds

• Each metal donates its valence electron(s) to form an electron cloud

• This leaves positive particles which are "cemented" together with the negative electron cloud, often called a “sea of electrons.”