General Chemistry Study Guide: Atomic Structure, Periodicity, and Chemical Bonding

Atomic Structure and Electronic Configuration

Nature of Light and Atomic Models

The study of atomic structure begins with understanding the nature of light and how it interacts with matter. Early models of the atom evolved as scientists observed atomic spectra and developed quantum mechanical theories.

• Electromagnetic Radiation: Light exhibits both wave-like and particle-like properties. Key terms include wavelength, frequency, and the speed of light.

• Photoelectric Effect: Demonstrates the particle nature of light; electrons are ejected from a metal surface when exposed to light of sufficient energy.

• Bohr Model of the Atom: Electrons occupy quantized energy levels. The energy difference between levels corresponds to the emission or absorption of photons.

• Hydrogen Atom Energy Levels: The energy of an electron in the nth level is given by: $E_n = -2.178 \times 10^{-18} \left(\frac{1}{n^2}\right) \text{J}$

• Rydberg Equation: Used to calculate the wavelength of light emitted or absorbed during electron transitions in hydrogen: $\frac{1}{\lambda} = R_H \left(\frac{1}{n_1^2} - \frac{1}{n_2^2}\right)$ where $R_H = 1.097 \times 10^7 \text{m}^{-1}$

• Quantum Mechanical Model: Electrons are described by wavefunctions (orbitals) and have quantized energies. The Heisenberg Uncertainty Principle states that the position and momentum of an electron cannot both be precisely known.

Quantum Numbers and Atomic Orbitals

Quantum numbers describe the properties of atomic orbitals and the electrons within them.

• Principal Quantum Number (n): Indicates the energy level and size of the orbital.

• Angular Momentum Quantum Number (l): Defines the shape of the orbital (s, p, d, f).

• Magnetic Quantum Number (ml): Specifies the orientation of the orbital.

• Spin Quantum Number (ms): Describes the spin of the electron (+1/2 or -1/2).

Electron Configurations and the Periodic Table

Electron configurations describe the arrangement of electrons in an atom. The Aufbau Principle, Pauli Exclusion Principle, and Hund's Rule govern the filling of orbitals.

• Aufbau Principle: Electrons fill orbitals of lowest energy first.

• Pauli Exclusion Principle: No two electrons in an atom can have the same set of four quantum numbers.

• Hund's Rule: Electrons occupy degenerate orbitals singly before pairing.

• Core and Valence Electrons: Core electrons are in filled inner shells; valence electrons are in the outermost shell and determine chemical properties.

• Noble Gas Configuration: Shorthand notation using the previous noble gas to represent core electrons.

Periodicity and Periodic Trends

Development of the Periodic Table

The periodic table arranges elements by increasing atomic number, revealing periodic trends in properties.

• Groups (Columns): Elements with similar valence electron configurations and chemical properties.

• Periods (Rows): Elements with the same number of electron shells.

Periodic Trends

Several properties of elements show predictable trends across periods and groups.

Property	Trend Across Period (Left to Right)	Trend Down Group (Top to Bottom)
Atomic Radius	Decreases	Increases
Ionization Energy	Increases	Decreases
Electron Affinity	Generally more negative	Less negative
Electronegativity	Increases	Decreases

• Effective Nuclear Charge (Zeff): The net positive charge experienced by valence electrons. Increases across a period.

• Shielding: Inner electrons shield outer electrons from the nucleus, reducing Zeff.

• Isoelectronic Series: Ions with the same number of electrons; size decreases with increasing nuclear charge.

Chemical Bonding

Ionic and Covalent Bonds

Chemical bonds form when atoms share or transfer electrons to achieve stable electron configurations.

• Ionic Bonds: Formed by the transfer of electrons from metals to nonmetals, resulting in oppositely charged ions.

• Covalent Bonds: Formed by the sharing of electrons between nonmetals.

• Metallic Bonds: Involve a 'sea' of delocalized electrons shared among metal atoms.

Lewis Structures and Resonance

Lewis structures represent the arrangement of valence electrons in molecules. Resonance structures depict delocalization of electrons.

• Drawing Lewis Structures: Count valence electrons, arrange atoms, and assign lone pairs and bonds to satisfy the octet rule.

• Resonance: When more than one valid Lewis structure can be drawn, the actual structure is a hybrid.

• Formal Charge: Calculated as: $\text{Formal Charge} = \text{Valence electrons} - \text{Nonbonding electrons} - \frac{1}{2}(\text{Bonding electrons})$

Bond Polarity and Electronegativity

Bond polarity arises from differences in electronegativity between bonded atoms.

• Nonpolar Covalent Bond: Electrons are shared equally.

• Polar Covalent Bond: Electrons are shared unequally, creating dipoles.

• Electronegativity: The ability of an atom to attract electrons in a bond. Increases across a period and decreases down a group.

Lattice Energy and Bond Strength

Lattice energy is the energy required to separate one mole of an ionic solid into gaseous ions. Bond strength is related to bond order and bond length.

• Lattice Energy Equation (Born-Haber Cycle): $\Delta H_{\text{lattice}} = \sum \text{(ionization energies, electron affinities, etc.)}$

• Bond Order: Number of shared electron pairs; higher bond order means stronger, shorter bonds.

Additional Topics

Lab Techniques and Mathematical Operations

Familiarity with significant figures, unit conversions, and basic laboratory procedures is essential for success in general chemistry.

• Significant Figures: Reflect the precision of measurements and calculations.

• Unit Conversions: Use dimensional analysis to convert between units.

• Lab Safety: Always follow proper safety protocols in the laboratory.

Exam Preparation Tips

• Bring a non-programmable calculator to the exam.

• Review all lecture notes, textbook chapters, and assigned problems.

• Practice drawing Lewis structures, predicting molecular shapes, and applying periodic trends.

Additional info: Some sections (e.g., VSEPR theory, molecular orbital theory) are noted as being on the final exam but not on the current exam. Focus on atomic structure, periodicity, and bonding for this assessment.