NATURE OF LIGHT AND MATTER

Terminology

This section introduces key terms related to the nature of light and matter, which are foundational for understanding quantum mechanics and atomic structure.

• Wavelength (λ): The distance between successive crests of a wave, typically measured in meters or nanometers.

• Amplitude: The height of a wave from its center line to its peak.

• Frequency (ν): The number of wave cycles that pass a given point per unit time, measured in hertz (Hz).

• Photon: A quantum of electromagnetic radiation, carrying energy proportional to its frequency.

• Quantum: The smallest discrete amount of a physical property, such as energy.

• Electromagnetic spectrum: The range of all types of electromagnetic radiation.

Key Equations and Applications

• Speed of Light:

• Energy of a Photon:

• Planck's Constant:

• Photoelectric Effect: Demonstrates that light has particle-like properties; electrons are ejected from a metal surface when light of sufficient frequency shines on it.

Example: Calculate the energy of a photon with a wavelength of 500 nm.

• First, convert wavelength to meters:

• Calculate frequency:

• Calculate energy:

Nature of Electromagnetic Radiation

• Electromagnetic radiation consists of oscillating electric and magnetic fields that travel through space.

• Visible light is only a small portion of the electromagnetic spectrum, which also includes radio waves, microwaves, infrared, ultraviolet, X-rays, and gamma rays.

Atomic Emission Spectra

• Atoms emit light at specific wavelengths, producing line spectra unique to each element.

• These spectra provide evidence for quantized energy levels in atoms.

ELECTRON CONFIGURATIONS AND PERIODIC PROPERTIES

Terminology

• Orbital: A region in an atom where there is a high probability of finding electrons.

• Electron configuration: The arrangement of electrons in an atom's orbitals.

• Pauli exclusion principle: No two electrons in an atom can have the same set of quantum numbers.

• Aufbau principle: Electrons fill the lowest energy orbitals first.

• Hund's rule: Every orbital in a subshell is singly occupied before any orbital is doubly occupied.

Electron Configurations

• Write electron configurations using the format: etc.

• For ions, add or remove electrons according to the charge.

Example: Electron configuration of :

Periodic Trends

• Atomic radius: Decreases across a period, increases down a group.

• Ionization energy: Increases across a period, decreases down a group.

• Electron affinity: Generally becomes more negative across a period.

• Electronegativity: Increases across a period, decreases down a group.

Additional info: These trends are explained by effective nuclear charge and electron shielding.

IONIC AND COVALENT BONDS

Terminology

• Ionic bond: Electrostatic attraction between oppositely charged ions.

• Covalent bond: Sharing of electron pairs between atoms.

• Polar covalent bond: Unequal sharing of electrons between atoms with different electronegativities.

• Nonpolar covalent bond: Equal sharing of electrons.

Lewis Symbols and Structures

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

• Lewis structures show how atoms are bonded and the arrangement of electrons.

Example: Lewis structure for water ():

• Oxygen in the center, two single bonds to hydrogen, two lone pairs on oxygen.

Bond Polarity and Electronegativity

• Bond polarity is determined by the difference in electronegativity between atoms.

• Electronegativity values can be used to predict bond type:

Bond Length and Strength

• Bond length: Distance between nuclei of bonded atoms.

• Bond strength: Energy required to break a bond.

• Shorter bonds are generally stronger.

LEWIS STRUCTURES AND MOLECULAR SHAPE

Lewis Structures

• Draw all possible resonance structures for molecules with delocalized electrons.

• Assign formal charges to atoms to determine the most stable structure.

Molecular Geometry (VSEPR Theory)

• Valence Shell Electron Pair Repulsion (VSEPR) theory predicts the 3D shape of molecules based on electron pair repulsion.

• Common geometries include linear, trigonal planar, tetrahedral, trigonal bipyramidal, and octahedral.

Molecular Polarity

• Determine if a molecule is polar by assessing both bond polarity and molecular geometry.

• Polar molecules have an uneven distribution of charge, leading to dipole moments.

VALENCE BOND THEORY

Terminology

• Hybridization: Mixing of atomic orbitals to form new hybrid orbitals suitable for bonding.

• Sigma (σ) bond: Formed by head-on overlap of orbitals.

• Pi (π) bond: Formed by side-to-side overlap of orbitals.

Hybridization Types

• sp: Linear geometry, 180° bond angle.

• sp2: Trigonal planar geometry, 120° bond angle.

• sp3: Tetrahedral geometry, 109.5° bond angle.

Example: In methane (), carbon is sp3 hybridized, forming four sigma bonds with hydrogen.

Additional info: These topics are essential for understanding atomic structure, chemical bonding, and molecular properties, which are foundational for General Chemistry.