Chapter 2 - Covalent Substances

Covalent Substances: Covalent Bonding (Part 1)

• Covalent Bonding: Involves the sharing of electrons between non-metal atoms to form molecules.

• Molecule: Two or more non-metal atoms chemically bonded together.

• Diatomic Molecule: Two non-metal atoms covalently bonded, can be the same or different elements.

• Octet Rule: Atoms are more stable with 8 electrons in their valence shell.

• Intramolecular Bond: Covalent bond within a molecule.

• Covalent bonds can be single, double, or triple, depending on the number of electron pairs shared.

Sharing Electrons

• Three main types of bonding between atoms:

  1. Covalent

  2. Ionic

  3. Metallic

• Covalent bonding involves sharing electrons in diatomic and larger molecules.

• Ionic bonding involves the transfer of electrons between atoms.

• Metallic bonding involves two or more metal atoms bonded together.

• Carbon (C) has four valence electrons.

• Nitrogen (N) has three valence electrons.

• Oxygen (O) has two valence electrons.

• Hydrogen (H) and halogens have one valence electron.

Single Covalent Bond

• Sharing of one pair of electrons.

• Example: Hydrogen gas (H₂), represented as H-H.

Double Covalent Bond

• Sharing of two pairs of electrons.

• Example: Oxygen gas (O₂).

Triple Covalent Bond

• Sharing of three pairs of electrons.

• Example: Nitrogen gas (N₂).

Molecular Formula

• Represents the number and type of atoms in a molecule.

• Does not show orientation or shape of the molecule.

• Example: Serotonin, C10H12N2O

Electron Dot (Lewis) Structures

• Show valence electrons as dots around the element symbol.

• Valence electrons less than four are shown distributed around the symbol.

• Valence electrons more than five are represented in pairs.

• A pair of valence electrons not involved in covalent bonding is a lone pair.

Structural Formula

• Lines represent a pair of shared electrons.

• Double lines indicate a double covalent bond.

• Triple lines indicate a triple covalent bond.

• Example: Serotonin, C{10}H{12}N_2O

Other Representations

• Skeletal Formula: Vertices are C atoms, H atoms are assumed to give each C atom 4 bonds.

• Space-Filling Model: Demonstrates the three-dimensional arrangement, shape, and size of atoms.

• Ball and Stick Model: Demonstrates the three-dimensional arrangement of atoms.

Examples of Molecules and their Representations

• Hydrogen gas (H₂): Electron dot and structural formulas.

• Oxygen (O₂): Electron dot and structural formulas.

• Chlorine gas (Cl₂): Electron dot and structural formulas.

• Nitrogen (N₂): Electron dot and structural formulas.

• Hydrogen chloride (HCl): Electron dot and structural formulas.

• Carbon dioxide (CO₂): Electron dot and structural formulas.

• Water (H₂O): Electron dot and structural formulas.

• Ammonia (NH₃): Electron dot and structural formulas.

• Methane (CH₄): Electron dot and structural, ball and stick, and space-filling models.

• Ethane (CH₃CH₃): Electron dot, structural, skeletal, and space-filling models.

• Ethene (CH₂CH₂): Electron dot, structural, skeletal, and space-filling models.

Worked Example: Ammonia

• (a) Molecular formula: NH₃

• (b) Structural formula: Illustrate single bonds between N and 3 H atoms.

• (c) Electron dot structure: Show 5 valence electrons around N, including one lone pair.

Covalent Substances: Covalent Bonding (Part 2)

Shapes of Molecules

• Molecules are three-dimensional (3D).

• The actual shape depends on the presence of lone pairs of electrons.

• Valence Shell Electron Pair Repulsion (VSEPR) Theory: Like charges repel.

• VSEPR theory is used to predict the shapes of molecules based on the repulsion of electron pairs.

Common Molecular Shapes:

• Linear: Two atoms share a pair of electrons.

• Bent: Central atom with two single bonds and two lone pairs.

• Pyramidal: Central atom with three single bonds and one lone pair.

• Tetrahedral: Central atom with four single bonds and no lone pairs.

Tetrahedral Shape (e.g., Methane, CH₄)

• Carbon has four valence electrons and forms four single bonds.

• No lone pairs around the C atom.

• Each bond angle is 109.5^\circ.

• Wedge-dash notation represents a 3D molecule in 2D.

Pyramidal Shape (e.g., Ammonia, NH₃)

• Nitrogen has five valence electrons and forms three single bonds.

• One lone pair around the N atom.

• Each bond angle is 107^\circ.

• Lone pairs are not included in the overall shape of a molecule.

Bent Shape (e.g., Water, H₂O)

• Oxygen has six valence electrons and forms two single bonds.

• Two lone pairs around the O atom.

• Each bond angle is 104.5^\circ.

Linear Shape (e.g., Hydrogen Chloride, HCl)

• Hydrogen has one valence electron and forms one single bond.

• Halogens have seven valence electrons and form one single bond.

• Each bond angle is 180^\circ.

Other Linear Molecules (e.g., Carbon Dioxide, CO₂)

• Electrons in double and triple bonds are considered in the same way as electrons in a single covalent bond.

Polarity

• Both the shape of a molecule and the types of atoms involved determine its polarity.

• Non-metals differ in their electronegativity.

• Electronegativity: How strongly an atom attracts bonding electrons towards itself.

• Polar Covalent Bond: Covalent bond between atoms with an unequal distribution of electrons.

Non-Polar Covalent Bonds

• Occur when two atoms of the same type are covalently bonded (e.g., Cl and Cl).

• Occur when two atoms have similar electronegativity (e.g., C and H).

• Electrons are shared equally in a non-polar bond.

Determining Polarity of Bonds

• Pauling scale of electronegativity is used.

• Difference in electronegativity:

  • Less than 0.4: Non-polar covalent bond

  • Between 0.4 and 2.0: Polar covalent bond

  • More than 2.0: Ionic bond

Molecule Polarity

• If a molecule contains only non-polar bonds, it is non-polar overall.

• If a molecule is perfectly symmetrical and contains polar bonds, it is non-polar overall.

• If a molecule contains at least one polar bond and is asymmetric, it is polar.

• Polar Molecule: Molecule with a partially positively charged end and a partially negatively charged end.

• Permanent Dipole Moment: Covalent molecules with a permanent positive and negative charge due to differences in electronegativity.

Covalent Substances: Intramolecular Bonding and Intermolecular Forces (Part 1)

Intramolecular vs. Intermolecular Forces

• Intramolecular Bonding: Very strong covalent bonding within molecules.

• Intermolecular Forces: Forces between molecules, weaker than intramolecular bonding.

• Types of Intermolecular Forces:

  • Dispersion Forces

  • Dipole-Dipole Attraction

  • Hydrogen Bonding

Dispersion Forces

• Weakest type of intermolecular force.

• Occur because electrons are constantly moving.

• Instantaneous dipole moment is created when electrons spontaneously move to one side of an atom.

• A temporary negative side of a molecule is attracted to a temporary positive side of another molecule.

• Strength increases with the number of electrons in atoms and molecules.

Dipole-Dipole Attraction

• Stronger than dispersion forces.

• Occurs between polar molecules due to permanent dipole-dipole moments.

Hydrogen Bonding

• Strongest form of intermolecular force and dipole-dipole attraction.

• Occurs between H on one molecule bound to F, O, or N and F, O, or N on another molecule.

• Due to the presence of lone pairs of electrons on F, O, or N.

Relative Strength of Intermolecular Forces

• Dispersion Forces: 1

• Dipole-Dipole Attraction: 10

• Hydrogen Bonds: 50

Covalent Substances: Intramolecular Bonding and Intermolecular Forces (Part 2)

Properties of Molecular Substances

• Determined by intermolecular forces.

• Intermolecular forces are responsible for melting point, boiling point, and hardness.

• The strength of intermolecular forces reflects the amount of energy it takes to break them.

• It takes more energy to break hydrogen bonds than dipole-dipole and dispersion forces.

Kinetic Energy and States of Matter

• Gases have more kinetic energy than liquids.

• Liquids have more kinetic energy than solids.

• Melting Point: The temperature at which a substance changes state from a solid to a liquid; depends on the type of intermolecular forces present.

• Boiling Point: The temperature at which a substance boils and changes state from a liquid to a gas.

Molecular Size and Shape

• As molecular size increases, so does the number of dispersion forces, leading to higher melting and boiling points.

• The more linear a molecule, the more adjacent molecules that can pack together, leading to higher melting and boiling points.

Hardness

• Harder substances have greater intermolecular forces between adjacent molecules.

Covalent Substances: Macromolecules

Carbon

• Non-metal with 4 valence electrons.

• Has three naturally occurring isotopes: ^{12}C, ^{13}C, and ^{14}C.

• The building block of life due to its ability to form complex, stable molecules.

Allotropes of Carbon

• Different physical forms of an element due to different structural arrangements of atoms.

• Examples: Diamond, graphite, charcoal, fullerene, graphene.

Diamond

• Bonded in a 3-dimensional (3D) covalent network lattice.

• Each carbon is covalently bonded to four other carbon atoms.

• Properties: High melting point, very hard, brittle, does not conduct electricity, high thermal conductivity, insoluble.

Applications of Diamond

• Jewellery

• Cutting tools (hardness resists wear)

• Thermal conductor in electrical components

• Optical components (diamond lasers)

• Abrasive (able to induce friction)

Graphite

• Structured as a covalent layer lattice with layers of 2-dimensional (2D) carbon lattices held together by weak dispersion forces.

• Each carbon is bonded to three other carbon atoms.

• The lattice has delocalized electrons that are free to move.

• Properties: High melting point, High thermal conductivity, Soft and slippery, less dense than diamond, insoluble, conducts electricity.

Applications of Graphite

• Pencils

• Electrical conductor (carbon brushes in electric motors, electrodes in batteries)

• Industrial lubricant (layers can slide over each other)