Unit 2: Covalent Bonding, Molecular Geometry, and Solutions

Covalent Bonding

Covalent bonds form when two atoms share electrons. The nature of the sharing determines the bond's polarity and properties.

• Nonpolar Covalent Bond: Electrons are shared equally between atoms (e.g., H2, O2).

• Polar Covalent Bond: Electrons are shared unequally, resulting in partial charges (e.g., H2O).

• Metallic Bond: Involves a 'sea' of delocalized electrons among metal atoms.

• Electronegativity: The ability of an atom to attract shared electrons. Fluorine is the most electronegative element.

Example: In H2O, the O-H bonds are polar covalent because oxygen is more electronegative than hydrogen.

Electron Configuration and the Octet Rule

Atoms tend to gain, lose, or share electrons to achieve a stable configuration, often resembling the nearest noble gas (the octet rule).

• Octet Rule: Atoms are stable with eight valence electrons.

• Exceptions: Hydrogen (stable with 2), Boron (often stable with 6), expanded octets for elements in period 3 or higher.

Molecular Geometry and VSEPR Theory

The shape of molecules is predicted by the Valence Shell Electron Pair Repulsion (VSEPR) theory, which states that electron pairs around a central atom arrange themselves to minimize repulsion.

• Linear: 2 electron groups, 180° bond angle (e.g., CO2).

• Trigonal Planar: 3 electron groups, 120° bond angle (e.g., BF3).

• Tetrahedral: 4 electron groups, 109.5° bond angle (e.g., CH4).

• Trigonal Pyramidal: 3 bonds + 1 lone pair (e.g., NH3).

• Bent: 2 bonds + 2 lone pairs (e.g., H2O).

Example: The shape of ammonia (NH3) is trigonal pyramidal due to one lone pair on nitrogen.

Bonding and Molecular Polarity

The polarity of a molecule depends on both the polarity of its bonds and its geometry.

• Nonpolar Molecule: Symmetrical shape and identical surrounding atoms (e.g., CO2).

• Polar Molecule: Asymmetrical shape or different surrounding atoms (e.g., H2O).

Example: CH4 is nonpolar, while H2O is polar.

Lewis Structures and Resonance

Lewis structures represent the arrangement of electrons in a molecule. Some molecules have resonance structures, where more than one valid Lewis structure can be drawn.

• Resonance: Delocalization of electrons across multiple atoms (e.g., O3, NO3-).

Bond Types and Overlap

Bonds form by the overlap of atomic orbitals:

• σ (sigma) bond: End-to-end overlap; all single bonds are sigma bonds.

• π (pi) bond: Side-to-side overlap; present in double and triple bonds.

Example: A triple bond (e.g., in N2) consists of one sigma and two pi bonds.

Solutions and Concentration Calculations

Solutions are homogeneous mixtures of solute and solvent. Concentration expresses the amount of solute in a given amount of solution.

• Solute: Substance being dissolved.

• Solvent: Substance doing the dissolving (usually present in greater amount).

• Molarity (M):

• Mass Percent:

• Parts per million (ppm):

Example: To find the molarity of a solution containing 10.0 g NaOH in 250.0 mL of solution:

• Calculate moles of NaOH:

• Convert volume to liters:

• Molarity:

Empirical and Molecular Formulas

The empirical formula gives the simplest whole-number ratio of atoms in a compound, while the molecular formula gives the actual number of atoms.

Periodic Trends and Electronegativity

Electronegativity increases across a period and decreases down a group. The most electronegative element is fluorine.

Sample Table: Properties of Covalent Compounds

Energy Changes in Chemical Reactions

When bonds form, energy is released; when bonds break, energy is absorbed.

• Exothermic Reaction: Releases energy (products have lower potential energy than reactants).

• Endothermic Reaction: Absorbs energy (products have higher potential energy than reactants).

Key Formulas and Equations

• Molarity:

• Mass Percent:

• ppm:

• Moles:

Sample Calculation: Molar Mass

To find the molar mass of MgCl2:

• Mg: 24.3 g/mol

• Cl: 35.5 g/mol × 2 = 71.0 g/mol

• Total: 24.3 + 71.0 = 95.3 g/mol

Summary Table: Types of Chemical Bonds

Additional info: Some explanations and examples have been expanded for clarity and completeness, based on standard General Chemistry curriculum.