Chem 220: Exam 3 Review – Light, Electron Configurations, Periodic Trends, and Chemical Bonding

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: $c = \lambda \nu$

• Energy of a Photon: $E = h \nu$

• Planck's Constant: $h = 6.626 \times 10^{-34} \text{J} \cdot \text{s}$

• 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: $500 \text{nm} = 5.00 \times 10^{-7} \text{m}$

• Calculate frequency: $\nu = \frac{c}{\lambda}$

• Calculate energy: $E = h \nu$

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: $1s^2 2s^2 2p^6 3s^2 3p^6$ etc.

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

Example: Electron configuration of $\text{Na}^+$: $1s^2 2s^2 2p^6$

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 ($\text{H}_2\text{O}$):

• 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:

Electronegativity Difference	Bond Type
0	Nonpolar covalent
0.1 - 1.7	Polar covalent
> 1.7	Ionic

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.

Electron Groups	Geometry	Bond Angle
2	Linear	180°
3	Trigonal planar	120°
4	Tetrahedral	109.5°
5	Trigonal bipyramidal	90°, 120°
6	Octahedral	90°

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 ($\text{CH}_4$), 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.