AP Chemistry Unit 2: Molecular and Ionic Compounds and Their Properties

Overview of Chemical Bonds

  • Study unit 2 covers molecular and ionic compounds and their properties, following unit 1's focus on atomic structure and periodic trends.

  • Section includes ionic bonds, covalent bonds, metallic bonds, and key concepts such as molecular structure, VSEPR theory, Lewis Dot diagrams, and hybridization.

Definition of Chemical Bond

  • Chemical Bond: A force that holds two atoms together, allowing them to behave as a unit (Ex: Water, H2O).

    • Without bonding, individual atoms exhibit different properties. Example: Water, composed of H2 and O, has distinct properties as a single entity compared to its separate atoms.

Importance of Chemical Bonds

  • Play a crucial role in defining physical properties of substances:

    • Ex: Different forms of carbon (diamond, graphite, charcoal) showcase varying properties influenced by structure and bonding.

Chemical Reactions and Bonds

  • Chemical Reaction: Involves breaking old bonds and forming new ones (e.g., H2 + O2 → H2O involves breaking H-H and O=O bonds).

  • Phase changes (e.g., liquid to solid) without bond formation/breaking are not classified as chemical reactions.

Types of Chemical Bonds

Ionic Bonds

  • Form between a metal and a non-metal.

  • Electron Transfer: One atom loses an electron (becoming cation) while another gains (becoming anion).

  • Strong attraction between oppositely charged ions (e.g., Na+ and Cl- in NaCl).

    • Example: Sodium Chloride formation from sodium cation and chloride anion.

Covalent Bonds

  • Form predominantly between two non-metals and involve electron sharing.

  • Two Types:

    • Nonpolar Covalent Bonds: Equal sharing of electrons (e.g., Cl2).

    • Polar Covalent Bonds: Unequal sharing (e.g., H2O).

Metallic Bonds

  • Occur between metal cations and delocalized electrons, forming the electron sea model.

  • Key properties: Conductivity, malleability, and ductility.

Electronegativity Trends

  • Electronegativity increases left to right (due to increasing nuclear charge) and decreases top to bottom (due to increasing distance from nucleus).

  • Determines bond type:

    • Ionic bonds: Large electronegativity difference (e.g., Na has 0.9, Cl has 3.0; difference: 2.1).

    • Covalent bonds: Small or no difference (e.g., N and O have a difference of only 0.5).

Coulomb's Law

  • Force between charges is calculated by F = k \frac{q1 q2}{r^2} ,

    • where q1 and q2 represent charge amounts and r is the distance between charges.

    • Stronger ionic bonds form with higher charges and smaller distances (e.g., MgO vs. NaCl).

Properties of Ionic Compounds

  • High melting/boiling points due to strong ionic bonds.

  • Poor electricity conductors in solid form due to fixed ion positions, but conduct when molten or dissolved in water (free movement of ions).

Covalent Bonding and Structure

  • Hybridization: Mixing of atomic orbitals to form new orbitals of the same energy (e.g., sp, sp2, sp3).

    • Determined by the number of electron pairs: 1 pair = sp, 2 pairs = sp2, 3 pairs = sp3.

  • Bond Length: Decreases and strength increases with more shared electron pairs (single vs. double vs. triple bonds).

Bonding Types

  • Sigma Bond: End-to-end overlap of orbitals.

  • Pi Bond: Side-by-side overlap, found in multiple bonds.

Molecular Structure and VSEPR Theory

  • VSEPR (Valence Shell Electron Pair Repulsion) helps predict molecular shapes based on repulsion between electron pairs.

  • Geometries:

    • MX2: Linear (180°)

    • MX3: Trigonal planar (120°)

    • MX4: Tetrahedral (109.5°), trigonal pyramidal (NH3), bent (H2O)

    • MX5: Trigonal bipyramidal (120° & 90°)

    • MX6: Octahedral (90°)

Resonance Structures

  • Some molecules can have equivalent Lewis structures; the actual bond is a hybrid of these forms (e.g., Ozone O3).

Exceptions to the Octet Rule

  • Some molecules have incomplete (underfilled) or expanded octets (overfilled) based on their central atom's group.

    • Ex: BeH2 has only 4 electrons.

Alloy Properties

  • Alloys (mixtures of metals) enhance properties such as strength, hardness, and resistance to corrosion (e.g., steel).

  • Types of alloys: substitutional (similar sizes) and interstitial (different sizes).

Summary

  • Bonding principles are central to understanding chemistry, influencing both the structure and behavior of compounds.