COVALENT BONDINGChapter 8

Section Overview• 8.1: Molecular Compounds• 8.2: The Nature of Covalent Compounds• 8.3: Bonding Theories • 8.4: Polar Bonds and Molecules

MOLECULAR COMPOUNDSSection 8.1

Molecules and Molecular Compounds

• The atoms held together by sharing electrons are joined by a covalent bond which is like a “tug of war” for the electrons.

• A molecule is a neutral group of atoms joined together by covalent bonds (ex. Oxygen).

• A diatomic molecule is a molecule consisting of two atoms (ex. Oxygen).

• A compound composed of molecules is called a molecular compound (ex. Water).

• Molecular compounds tend to have relatively lower melting and boiling points than ionic compounds and are gases or liquid at room temperature.

Molecular Formulas• A molecular formula is a chemical formula of a molecular

formula.• A molecular formula shows how many atoms of each

element a molecule contains (ex. water = H20).

• However, a molecular formula doesn’t tell you about a molecule’s structure or how the atoms are arranged in the compound.

THE NATURE OF COVALENT COMPOUNDSSection 8.2

The Octet Rule in Covalent Bonding• In forming covalent bonds, electron sharing usually occurs

so that atoms attain the electron configuration of noble gases (eight valence electrons).

• Groups 4A, 5A,6A, and 7A are likely to form covalent bonds.

Single Covalent Bonds• Two toms held together by sharing a pair of electrons are

joined by a single covalent bond.• An electron dot structure can represent the shared pair of

electrons of the covalent bond by two dots.• A structural formula represents covalent bonds by dashes

and shows the arrangement of atoms.• Examples:• Hydrogen Gas =• Fluorine Molecule =

Single Covalent Bonds• A pair of valence electrons that is not shared between

atoms is called an unshared pair or a lone pair.• Water, for example, contains three atoms but only two

covalent bonds.• Other examples are ammonia and methane.

Single Covalent Bonds• Example Problem: Draw an electron dot structure for

hydrochloric acid (HCl)

Single Covalent Bonds• Example Problem: Draw an electron dot structure for

hydrochloric acid (HCl)• Solution:

First, draw separate electron dot structures.

We see that hydrogen has one valence and chlorine has seven. So, hydrogen will share its one with chlorine to form an octet.

Double and Triple Covalent Bonds• Atoms form double or triple covalent bonds if they can

attain a noble gas structure by sharing two pairs or three pairs of electrons.

• A bond that involves two shared pairs of electrons is a double covalent bond.

• A bond that involves three shared pairs of electrons is a triple covalent bond.

• An example of an element whose molecules contain triple bonds in nitrogen which has five valence electrons.

• Each nitrogen atom in the nitrogen molecule must gain three electrons to obey to octet rule.

• Another example is carbon dioxide, which has two double bonds.

Double and Triple Covalent Bonds

Coordinate Covalent Bonds• A coordinate covalent bond is a covalent bond in which

one atom contributes both bonding electrons.• An example is carbon monoxide.• A carbon atom needs to gain four electrons and an

oxygen atom needs two electrons.• A double covalent bond is formed when these atoms

come together.• However, carbon is unstable in that configuration.• So, oxygen donates a pair of unshared electrons.

Coordinate Covalent Bonds• Other examples can be found in polyatomic ions.• A polyatomic ion is a tightly bound group of atoms that

has a positive or negative charge and behaves as a unit.• An ammonium ion is an example.• The ammonium ion forms when a positively charged

hydrogen ion attaches to the unshared electron pair of an ammonia molecule.

Coordinate Covalent Bonds• Example Problem: Draw the electron dot structure for the

hydronium ion (H3O+).

Coordinate Covalent Bonds• Example Problem: Draw the electron dot structure for the

hydronium ion (H3O+).

• Solution:

Draw the electron dot structure of a water molecule.

Draw the electron dot structure for a single hydrogen ion.

Put them together into one structure.

Bond Dissociation Energies• The energy required to break the bond between two

covalently bonded atoms is known as the bond dissociation energy.

• This is expressed as the energy needed to break one mole of the bonds.

• A large bond dissociation energy corresponds to a strong covalent bond and thus a stable compound.

Resonance• A resonance structure is a structure that occurs when it is

possible to draw two or more valid electron dot structure that have the same number of electron pairs for a molecule or ion.

• Ozone is an example

• Earlier scientists thought that the electron pairs rapidly flipped back and forth between the different electron dot structures so a double headed arrow is used.

• However, today, we know that they do not actually resonate back and forth.

• The actual bonding of oxygen atoms in ozone is a hybrid, or mixture, of the extremes represented by the resonance forms.

Exceptions to the Octet Rule• The octet rule cannot be satisfied in molecules whose

total number of valence electrons is an odd number.• There are also molecules which an atom has fewer, or

more, than a complete octet of valence electrons.• Some examples include nitrogen dioxide (NO2), chlorine

dioxide (ClO2), nitric oxide (NO), boron triflouride (BF3), phosphorus trichloride (PCl3), and phosphorus pentachloride (PCl5).

BONDING THEORIESSection 8.3

Molecular Orbitals• This model assumes that when two atoms combine, their

atomic orbitals overlap and produce molecular orbitals, or orbitals that apply to the entire molecule.

• Just as an atomic orbital belongs to a particular atom, a molecular orbital belongs to a molecule as a whole.

• Two electrons are required to fill a molecular orbital.• A molecular orbital that can be occupied by two electrons

of a covalent bond is called a bonding orbital.

Molecular Orbitals• Sigma Bonds: When two atomic orbitals combine to form

a molecular orbital that is symmetrical around the axis connecting the two atomic nuclei.

• Pi Bonds: The bonding electrons are most likely to be found in sausage-shaped region above and below the bond axis of bonded atoms.

VSPER Theory • According to the VSPER theory, the repulsion between

electrons pairs causes molecular shapes to adjust so that the valence-electron pairs stay as far apart as possible which affects their shape.

• Unpaired electrons are also important in determining shape.• The unshared pair strongly repels the bonding pairs which

pushes them closer together.• For example, in methane, all electrons are shared and the

angle between the hydrogens and the central carbon is 109.5.

• In ammonia, one of the electron pairs is unshared and the angles between the hydrogen and the central nitrogen is 107.

VSPER Theory

VSPER Theory

Hybrid Orbitals• Orbital hybridization provides information on both

molecular bonding and shape.• In hybridization, several atomic orbitals mix to form the

same total number of equivalent hybrid orbitals.• Hybridization single bonds: Carbon’s outer electron

configuration is 2s22p2, but one 2s is promoted to 2p to form 2s12p3, allowing it to bond to four hydrogens in methane. These orbitals mix together to form four sp3 orbitals. In methane, each of the four sp3 hybrid orbitals overlaps with a 1s orbital of hydrogen.

Hybrid Orbitals• Hybridization double bonds: Ethene has one carbon-

carbon double bond and four carbon-carbon single bonds. In ethene, sp2 hybrid orbitals form from the combination of one 2s and two 2p atomic orbitals of carbon. Two sp2 hybrid orbitals of each carbon from sigma-bonding with the four available hydrogen 1s molecules. The third sp2 orbitals of each of the two carbons overlap to form a carbon-carbon sigma-bonding orbital. The non-hybridized 2p carbon orbitals overlap side-by-side to form a pi-bonding orbital.

Hybrid Orbitals• Hybridization triple bonds: Ethyne (aka acetylene) is a

linear molecule. The best hybrid orbital description is obtained if a 2s atomic orbital of carbon mixes with only one of three 2p atomic orbitals. The result is two sp hybrid orbitals for each carbon. One sp hybrid orbital from each carbon overlaps with a 1s orbital or hydrogen to form a sigma bond. The other sp hybrid orbital of each carbon overlaps to form a carbon-carbon sigma bond. The two p atomic orbitals from each carbon also overlap.

Hybrid Orbitals: Methane

Hybrid Orbitals: Ethene

Hybrid Orbitals: Ethyne

POLAR BONDS AND MOLECULESSection 8.4

Bond Polarity• The bonding pairs of electrons in covalent bonds are

pulled, kind of like in a tug of war, between the nuclei of the atoms sharing the electrons.

• When the atoms in the bond pull equally, the bonding electrons are shared equally, and the bond is called a nonpolar covalent bond.

• A polar covalent bond, or just polar bond, is a covalent bond between atoms in which the electrons are shared unequally.

• The more electronegative atom attracts electrons more strongly and gains a slightly negative charge.

• The less electronegative atom has a slightly positive charge.

Bond Polarity• For example, HCl is made of hydrogen and chlorine, to

form hydrogen chloride.• Hydrogen has an electronegativity of 2.1 and chlorine has

an electronegativity of 3.0.• These values are significantly different, so the covalent

bond in hydrogen chloride is polar.• So, chlorine gains a slight negative charge and hydrogen

gains a slight positive charge.

Bond Polarity

Bond Polarity• Example Problem: Which type of bond will form between

each of the following pairs of atoms?

a. N and H

b. F and F

c. Ca and Cl

d. Al and Cl

Bond Polarity• Example Problem: Which type of bond will form between

each of the following pairs of atoms?

a. N and H• Solution:

Find individual electronegativity values

Determine difference

Use chart to determine bond type

N (3.0) and H (2.1)

Difference = 0.9

Bond = covalent

Bond Polarity• Example Problem: Which type of bond will form between

each of the following pairs of atoms?

b. F and F

c. Ca and Cl

d. Al and Cl• Solution:

b. F (4.0) and F (4.0); 0.0; nonpolar covalent

c. Ca (1.0) and Cl (3.0); 2.0; ionic

d. Al (1.5), Cl (3.0); 1.5; polar covalent

Polar Molecules• In a polar molecule, one end of the molecule is slightly

more negative and the other end is slightly positive.• A molecule that has two poles is called a dipolar molecule,

or dipole. • When polar molecules are placed between oppositely

charged plates, they tend to become oriented with respect to the positive and negative poles.

• The effect of polar bonds on the polarity of an entire molecule depends on the shape and the orientation of the polar bonds.

• For example, CO2 has two polar bonds, but CO2 is linear and the bonds are in opposite directions making it overall nonpolar. Water also has two polar bonds, but it is bent, making it polar.

Attractions Between Molecules• Intermolecular attractions are weaker than either ionic or

covalent bonds.• However, these attractions are responsible for

determining whether a molecular compounds is a gas, liquid, or solid at a given temperature.

• Two types of attractions are van der Waals Forces and hydrogen bonds.

Attractions Between Molecules• Van der Waals Forces: The two weakest attractions

between molecules.• Dipole interactions occur when polar molecules are

attracted to one another. The attraction involved occurs between the oppositely charged regions of polar molecules. The slightly negative region is attracted to the slightly positive region.

• Dispersion forces are caused by the motion of electrons and are the weakest of all intermolecular interactions. When the moving electrons happen to be momentarily more on the side of a molecule closest to a neighboring molecule, their electric force influences the neighboring molecule’s electrons to be momentarily more on the opposite side.

Attractions Between Molecules• Hydrogen Bonds: Attractive forces in which a hydrogen

covalently bonded to a very electronegative atom is also weakly bonded to an unshared electron pair of another electronegative atom. This other atom can be in the same molecule or in a nearby molecule.

• Hydrogen bonds are present in water.• One hydrogen from a water

molecule, is attracted to an

oxygen of another water molecule.

Intermolecular Attractions and Molecular Properties• At room temperature, some compounds are gases while

others are liquids or solids.• The physical properties of a compound depend on the type

of bonding it displays.• A greater range occurs among covalent compounds due to

the varying intermolecular attractions.• In most solids formed from molecules, only the weak

interactions between molecules need to be broken in order for it to melt.

• However, there are network solids, in which all of the atoms are covalently bonded to each other and can only melt when all covalent bonds throughout are broken (ex. Diamond and silicon carbide).

Characteristics of Ionic and Covalent Compounds

Characteristic Ionic Compound Covalent Compound

Representative unit Formula unit Molecule

Bond formation Transfer of one or more electrons between atoms

Sharing of electron pairs between atoms

Type of elements Metallic and nonmetallic

Nonmetallic

Physical state Solid Solid, liquid, or gas

Melting point High Low

Solubility in water Usually high High to low

Electrical conductivity

Good conductor Poor to nonconducting