2. Covalent Substances

Definition: Covalent bonds form between non-metal atoms through the sharing of valence electrons, resulting in each atom achieving a full outer shell.
Molecule: A molecule is defined as a group of two or more non-metal atoms that are covalently bonded.
Intramolecular Bonds: The bonds between atoms within a molecule are known as intramolecular bonds.

Formation of Single Bond: Occurs when two atoms share one electron each.
Example: Chlorine molecule (Cl₂) where two chlorine atoms share electrons to complete their valence shells.
Electrons Required: Atoms share the number of electrons necessary to fill their respective valence shells; e.g., hydrogen (H) and chlorine (Cl) each need one electron, leading to the formation of H₂.

Definition: A double covalent bond is formed when two atoms share two pairs of electrons.
Characteristics: Double bonds are shorter and stronger compared to single bonds.
Example: Oxygen molecule (O₂), where two oxygen atoms share two pairs of electrons.

Definition: A triple covalent bond is where each atom shares three pairs of electrons.
Characteristics: These bonds are shorter and stronger than double bonds.
Example: Nitrogen molecule (N₂) formed by the sharing of three pairs of electrons between nitrogen atoms.

Composition: Atoms of different elements can combine to form molecular compounds.
Example: In hydrogen chloride (HCl), one electron from hydrogen and one from chlorine are shared, resulting in a single covalent bond.

Definition: Molecules composed of two or more atoms are referred to as polyatomic molecules.
Example 1: Water (H₂O), containing two single covalent bonds and two lone pairs on the oxygen atom.
Example 2: Methane (CH₄), which features four single covalent bonds and no lone pairs.

Definition: Lewis structures, or electron dot structures, represent molecules showing the arrangement of valence electrons.
Representation:
- Bonding Electrons: Shown as dots or lines between the atoms to indicate bonding pairs.
- Non-bonding Electrons: Represent non-bonding electron pairs (lone pairs) as dots around the atoms.
Flexibility: Crosses can be used alongside dots to signify shared electrons.

Valence Electrons: Determine the number of valence electrons in each atom (e.g., H has 1, O has 6).
Shared Electrons: Identify how many electrons each atom will share to fill their outer shell.
Symbol Representation: Write the element symbols and draw the valence electrons accordingly, unpaired for sharing.
Arrangement: Arrange the atoms to fulfill the octet rule (8 electrons for most, 2 for H).

Definition: Valence Shell Electron Pair Repulsion Theory (VSEPR) is used to predict molecular shapes based on electron pair repulsion.
Key Concepts:
1. Valence electrons are grouped into electron groups (bonding pairs, lone pairs).
2. Repulsion occurs between these electron groups.
3. Molecular shape allows groups to maximize distance from each other.

Counting Groups: Determine the number of bonds (single, double, or triple) around the central atom, including lone pairs.

Tetrahedral: Four bonds around the central atom; bond angle is 109°. Example: Methane (CH₄).
Pyramidal: Three bonds and one lone pair; bond angle is approximately 107°. Example: Ammonia (NH₃).
V-shaped (Angular): Two bonds and two lone pairs; bond angle is about 104.5°. Example: Water (H₂O).
Linear: Diatomic molecules with a bond angle of 180°. Examples: Chlorine (Cl₂).
Trigonal Planar: Three groups around a central atom with no lone pairs; bond angle is 120°. Examples: Methanal (CH₂O), Boron trihydride (BH₃).

Definition: Structural formulas visually depict the shape of molecules.
1. Replace each bonding pair of electrons in the Lewis structure with a line.
2. Draw the molecule according to expected shape based on VSEPR theory.

Neutral Molecules: Molecules like H₂O are neutral due to equal proton and electron numbers but can act charged due to electron distribution.
Electronegativity: Different elements have varying electronegativities; more electronegative atoms attract bonding electrons better.

Non-Polar Bonds: Shared equally between identical atoms; no charge on either side.
Polar Bonds: When one atom is more electronegative, it leads to a dipole, producing unequal sharing of electrons causing partial charges (δ+ and δ-).
Polarity Calculation: Polar nature depends on the electronegativity difference; the Pauling scale is used for quantification.

Electronegativity Difference: Bond types can be predicted based on electronegativity differences:
- Values from 0.5 to 2.1 indicate a polar bond.
- Very high differences suggest complete electron transfer resulting in ionic bonds.

Determining Polarity in Complex Molecules: Check for polar bonds and assess symmetry.
1. Symmetrical molecules (even if polar bonds exist) are non-polar.
2. Asymmetrical molecules are typically polar due to non-canceling dipoles, denoting partial charges.

Definition: Forces between different molecules (weaker than intramolecular covalent bonds) that break upon heating.
Types of Intermolecular Forces:
1. Dispersion Forces (Weakest): Present in all molecules; form through temporary dipoles.
2. Dipole-Dipole Bonds: Present in polar molecules due to permanent dipole attractions.
3. Hydrogen Bonds (Strongest): Form between molecules containing O—H, F—H, or N—H bonds; notable for strong dipole interactions due to hydrogen's small size, allowing close proximity to electronegative atoms.
Boiling/Melting Points: Stronger intermolecular forces correlate with higher melting and boiling points.

Formation: Occurs from the random movement of electrons leading to temporary dipoles.
Influence of Molecular Mass: Larger mass leads to stronger dispersion forces and higher melting/boiling points due to a greater density of electrons.

Hydrogen bonds considerably raise the melting and boiling points of compounds compared to those containing only weaker intermolecular forces.

Comparison of Boiling Points: Hydrogen fluoride (HF) versus neon (Ne) exemplifies the strength of hydrogen bonds over dispersion forces, despite similar molecular weights.
Justification Question: When comparing boiling points of OF₂ (-145°C) and CF₄ (-128°C), the reasons for the difference in intermolecular forces should be elaborated.

Definition: Continuous 3D structures formed through covalent bonding that create repeating atomic patterns.
Allotropes of Carbon: Examples include:
- Diamond: A covalent lattice where each carbon is bonded to four others; notable for hardness and electrical insulating.
- Graphite: Carbon atoms bonded in layers; soft due to weak forces between layers; conducts electricity due to free-moving electrons.
- Amorphous Carbon: Lacks a consistent structure; produced from incomplete combustion of carbon-containing materials.

Differences in Properties: Attributes concerning hardness and electrical conductivity can be justified by the bonding and structural aspects of diamond and graphite, showcasing the significance of covalent structures in determining physical properties.