Unit 1 - Principles of Chemistry: Chapter 8: Covalent Bonding

Covalent Bonding

Introduction

• Covalent bonding involves the sharing of electrons between atoms.
• Covalent compounds are more abundant than ionic compounds.
• Water ($H_2O$) is an example of a covalent compound, while salt dissolved in seawater is an ionic compound.

Extension Work

• In simple cases, each atom contributes one electron to the shared pair.
• However, both electrons can originate from the same atom.

Hint

• Electrons are represented as dots or crosses for clarity, but they are identical.

Reminder

• Atoms with 8 outer shell electrons are isoelectronic with noble gases.

Covalent Bonding in a Hydrogen Molecule ($H_2$)

• Covalent bonds are depicted using dot-and-cross diagrams.
• Hydrogen nuclei are strongly attracted to the shared electron pair.
• Hydrogen forms diatomic molecules ($H_2$) with a strong covalent bond.
• Molecules consist of a fixed number of atoms connected by covalent bonds.
• Hydrogen molecules are diatomic.
• Proteins and DNA can contain thousands of atoms joined by covalent bonds.

Significance of Noble Gas Structures

• In $H_2$, each hydrogen atom shares one electron, achieving a noble gas configuration (helium).
• The number of protons defines the element, not the number of electrons.
• Atoms share electrons to achieve a full outer shell: 2 for hydrogen, and 8 for most other atoms (octet rule).
• When a central atom is bonded to other atoms (e.g., $CH<em>4$ or $PCl</em>5$), the outer atoms usually have 8 electrons (or 2 for H).
• Exceptions exist where the central atom does not have 8 electrons.

Stability and Energy

• Chemical stability is related to energy: lower energy states are more stable.
• Bond formation releases energy, resulting in a more stable substance.

Covalent Bonding in Hydrogen Chloride (HCl)

• Chlorine has 7 outer shell electrons and shares 1 with hydrogen to achieve a noble gas configuration.
• Chlorine's electron configuration becomes [2, 8, 8], like argon.
• Hydrogen attains 2 electrons, like helium.
• Only outer shell electrons participate in bonding at this level.
• Covalent bonds can be represented by lines.

Why Hydrogen Forms Molecules

• Bond formation releases energy, increasing stability.
• More bonds lead to greater energy release and greater stability.
• $H_2$ molecule is more stable than two separate H atoms.

Covalent Bonding in a Chlorine Molecule ($Cl_2$)

• Each chlorine atom shares 1 electron to achieve 8 electrons in the outer shell.

Covalent Bonding in Methane ($CH_4$)

• Carbon has 4 outer shell electrons and shares 1 with each of 4 hydrogen atoms.
• Carbon forms 4 covalent bonds, one with each hydrogen atom.

Covalent Bonding in Ammonia ($NH_3$)

• Nitrogen has 5 outer shell electrons and shares 3 with three hydrogen atoms.
• Nitrogen forms 3 covalent bonds, one with each hydrogen atom.

Covalent Bonding in Water ($H_2O$)

• Oxygen has 6 outer shell electrons and shares 2 with two hydrogen atoms.
• Oxygen forms 2 covalent bonds, one with each hydrogen atom.
• Water has two bonding pairs and two lone pairs of electrons around the central oxygen atom.
• The four pairs of electrons are arranged tetrahedrally, resulting in a bent or V-shaped molecular geometry.
• Water is a polar molecule due to its bent shape and unequal electron distribution.

Covalent Bonding in Ethane ($C<em>2H</em>6$)

• Ethane has a carbon-carbon covalent bond and carbon-hydrogen bonds.

What is a Covalent Bond?

• Covalent bonds are electrostatic attractions between positively charged nuclei and negatively charged shared electrons.
• Particles are held together due to attraction between positive and negative charges.

Covalent Bonding in an Oxygen Molecule ($O_2$): Double Bonding

• Oxygen has 6 outer shell electrons.
• Two oxygen atoms share 2 electrons each, forming a double covalent bond.
• Each oxygen atom ends up with 8 electrons in its outer shell.

Triple Bond in a Nitrogen Molecule ($N_2$)

• Nitrogen has 5 outer shell electrons.
• Two nitrogen atoms share 3 electrons each, forming a triple covalent bond.
• Each nitrogen atom achieves 8 electrons in its outer shell.
• The triple bond is very strong, making nitrogen relatively unreactive.

Covalent Double Bonding in Carbon Dioxide ($CO_2$)

• Carbon has 4 outer shell electrons and oxygen has 6.
• Each oxygen atom shares 2 electrons with the carbon atom.
• Two double bonds are formed between the carbon and the two oxygens.
• All atoms have 8 electrons in their outer shells.

Double Bond in Ethene ($C<em>2H</em>4$)

• Ethene is similar to ethane but has a double bond between the carbon atoms.

Organic Molecules Containing Halogen Atoms

• Bromomethane ($CH_3Br$): Three H atoms and one Br atom are joined to the central C atom.
• Bromine (Br) is in Group 7 and has 7 outer shell electrons, forming 1 covalent bond.

More Difficult Molecules

• Outer atoms typically have 8 electrons (or 2 for hydrogen).
• Central atom may have more or fewer than 8 electrons.

Boron Trifluoride ($BF_3$)

• Boron has only 3 outer shell electrons and can share a maximum of 3 electrons.
• Each fluorine atom shares 1 electron, resulting in boron having 6 electrons in its outer shell.

Sulfur Dioxide ($SO_2$)

• Sulfur is the central atom, and oxygen atoms are the outer atoms.
• Each oxygen atom shares 2 electrons with the sulfur atom, forming two double bonds.
• Sulfur originally has 6 electrons and shares 4, resulting in sulfur having 10 electrons in its outer shell.

Intermolecular Forces and Giant Covalent Structures

Simple Molecular Structures

• Molecules contain fixed numbers of atoms joined by strong covalent bonds.
• Intermolecular forces exist between molecules, holding them in liquid or solid states.
• Intermolecular forces are weaker than covalent bonds.
• Boiling water breaks intermolecular forces, not covalent bonds.
• Simple molecular structures: substances with molecules attracted by intermolecular forces (e.g., $H<em>2O$, $CO</em>2$, $CH<em>4$, $NH</em>3$, $C<em>2H</em>6$).
• Substances with simple molecular structures are typically gases, liquids, or low-melting-point solids.
• Little energy is needed to overcome weak intermolecular forces.

Melting and Boiling Points

• The halogens (Group 7) have simple molecular structures with intermolecular forces.
• Melting and boiling points increase with relative molecular mass.

Table 8.1: Melting and Boiling Points of Halogens

• Boiling points increase down the group, indicating stronger intermolecular forces with increasing relative molecular mass.

Other Physical Properties

• Covalent molecular compounds do not conduct electricity due to the absence of ions and immobile electrons.
• They tend to be insoluble in water but soluble in organic solvents.

Giant Covalent Structures

• Diamond: a form of pure carbon, each carbon atom forms four covalent bonds in a tetrahedral arrangement.
• It's a giant covalent structure extending in three dimensions.
• Diamond has a very high melting and boiling point due to strong covalent bonds.
• Melting or boiling simple molecular structures only requires overcoming intermolecular forces, not breaking covalent bonds.

Properties of Diamond

*Hardness (the hardest naturally occurring substance).
*Non-conductivity of electricity (all electrons are tightly held in covalent bonds).
*Thermal conductivity (vibrations are quickly transmitted through the structure).
*Insolubility (strong covalent bonds would need to be broken to dissolve it).

Graphite

• Graphite: another form of carbon with a layer structure.

Properties of Graphite

*Softness (layers slide over each other easily).
*High melting and boiling points (covalent bonds must be broken to disrupt the structure).
*Electrical conductivity (delocalized electrons move throughout the layers).
*Insolubility (strong covalent bonds).
*Lower density than diamond (layers are relatively far apart).

Fullerenes

• Fullerenes (e.g., $C_{60}$ fullerene) consist of molecules with weak intermolecular forces.

Properties of $C_{60}$ Fullerene

*Lower melting and boiling points than diamond and graphite (only intermolecular forces need to be broken).
*Solubility in some solvents (weak intermolecular forces).
*Non-conductivity of electricity (electrons cannot jump from molecule to molecule).

Allotropes of Carbon

• Diamond and graphite are allotropes of carbon (different forms of the same element).
• $C_{60}$ fullerene is another allotrope.