Chapter 7: Thermochemistry – Energy and Chemical Reactions

I. Energy

Understanding energy and its transformations is essential in chemistry, especially in the context of chemical reactions. This section covers the fundamental concepts of energy, heat, work, and internal energy.

• Terms and Definitions:

  • Energy: The capacity to do work or produce heat.

  • Kinetic Energy (KE): Energy due to motion.

  • Potential Energy (PE): Energy due to position or composition.

  • System and Surroundings: The system is the part of the universe under study; the surroundings are everything else.

• Energy Transfer as Heat:

  • Heat (q): Energy transferred due to temperature difference.

  • Measured in joules (J) or calories (cal).

  • Heat flows from higher to lower temperature.

• Energy Transfer as Work:

  • Work (w): Energy transfer when a force moves an object. In chemistry, often refers to pressure-volume work:

• Internal Energy (U):

  • The total energy (kinetic + potential) of a system.

  • Change in internal energy:

II. Chemistry and Energy

• Enthalpy and Enthalpies of Reaction:

  • Enthalpy (H): The heat content of a system at constant pressure.

  • Change in enthalpy:

  • For reactions at constant pressure, equals the heat absorbed or released.

• Hess’s Law:

  • The total enthalpy change for a reaction is the same, no matter how many steps the reaction is carried out in.

  • Allows calculation of for complex reactions by adding enthalpy changes of individual steps.

• Enthalpies of Formation:

  • Standard Enthalpy of Formation (): The enthalpy change when one mole of a compound forms from its elements in their standard states.

  • Used to calculate reaction enthalpies:

Chapter 8: Electronic Structure of Atoms

I. Introduction to Electronic Structure

This section explores how light interacts with matter and how these interactions reveal the arrangement of electrons in atoms.

• Interaction of Light and Matter:

  • Properties of Light: Light exhibits both wave-like and particle-like properties. Key properties include wavelength (), frequency (), and speed (), related by .

  • Absorption and Emission of Light: Atoms absorb energy and electrons move to higher energy levels (excitation); when electrons return to lower levels, light is emitted.

  • Bohr Model: Electrons orbit the nucleus in quantized energy levels. Energy differences correspond to specific wavelengths of light.

II. Quantum Mechanics

• The Hydrogen Atom:

  • Atomic Orbitals and Quantum Numbers: Orbitals are regions of space where electrons are likely to be found. Four quantum numbers describe each electron:

    • Principal (): Energy level

    • Angular momentum (): Shape of orbital

    • Magnetic (): Orientation

    • Spin (): Electron spin direction

  • Shapes of Orbitals: s (spherical), p (dumbbell), d, and f orbitals, each with characteristic shapes and orientations.

• Many-Electron Atoms – Another Quantum Number:

  • Electron-electron interactions lead to additional quantum numbers and energy splitting (e.g., spin pairing).

• Electron Configurations and the Periodic Table:

  • Electron configurations describe the arrangement of electrons in orbitals.

  • Aufbau principle, Pauli exclusion principle, and Hund’s rule guide electron filling order.

  • Periodic trends (e.g., atomic size, ionization energy) arise from electron configurations.

Chapter 10: Chemical Bonding I – Basic Concepts

I. Covalent Bonds and Lewis Bonding Theory

Chemical bonding explains how atoms combine to form molecules. Lewis theory provides a simple model for covalent bonding.

• Lewis Symbols, the Octet Rule, and Covalent Bonding:

  • Lewis symbols represent valence electrons as dots around element symbols.

  • The octet rule: Atoms tend to gain, lose, or share electrons to achieve eight valence electrons (noble gas configuration).

  • Covalent bonds involve sharing of electron pairs between atoms.

• Other Types of Covalent Bonds:

  • Multiple Covalent Bonds: Double and triple bonds involve sharing two or three pairs of electrons, respectively.

  • Coordinate Covalent Bonds: Both electrons in the shared pair come from the same atom.

• Sharing of Electrons in Covalent Bonds:

  • Electronegativity: The ability of an atom to attract shared electrons. Differences in electronegativity lead to bond polarity.

  • Bond Polarity: Polar covalent bonds have unequal sharing of electrons, resulting in partial charges ( and ).

II. Molecular Compounds and Lewis Structures in Detail

• Drawing Lewis Structures:

  • Count total valence electrons, arrange atoms, connect with single bonds, complete octets, and assign remaining electrons as lone pairs.

  • Formal Charge: Used to determine the most stable Lewis structure.

  • Exceptions to the Octet Rule: Some molecules have fewer or more than eight electrons (e.g., BF3, SF6).

  • Resonance Structures: Some molecules are best represented by two or more Lewis structures; the actual structure is a resonance hybrid.

III. Molecular Shapes and Polarity

• VSEPR Theory: Valence Shell Electron Pair Repulsion (VSEPR) theory predicts molecular shapes based on electron pair repulsions.

• Polarity of Molecules: Molecular shape and bond polarity together determine if a molecule is polar (has a net dipole moment).

Chapter 11: Chemical Bonding II – Valence Bond and Hybridization Theories

I. Covalent Bonding and Orbital Overlap

Covalent bonds form when atomic orbitals overlap, allowing electrons to be shared between atoms.

II. Hybrid Orbitals

• Electron Groups and Hybridization Schemes:

  • Atomic orbitals mix to form hybrid orbitals (e.g., sp, sp2, sp3), which explain observed molecular geometries.

  • The number of electron groups around a central atom determines the hybridization scheme.

• Multiple Covalent Bonds and Hybridization:

  • Double and triple bonds involve pi (π) and sigma (σ) bonds, with pi bonds formed from unhybridized p orbitals.

Additional info: This guide summarizes the main topics for Exam #3, focusing on thermochemistry, atomic structure, and chemical bonding, including both Lewis and valence bond theories. Students should be familiar with definitions, equations, and applications for each topic.