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

Atomic Structure and Periodicity

Quantum Numbers

Quantum numbers describe the unique quantum state of an electron in an atom. They are essential for understanding electron configurations and the arrangement of electrons in atoms.

• Principal Quantum Number (n): Indicates the main energy level or shell. Values: n = 1, 2, 3, ...

• Angular Momentum Quantum Number (l): Defines the subshell (s, p, d, f). Values: l = 0 to n-1.

• Magnetic Quantum Number (ml): Specifies the orientation of the orbital. Values: ml = -l to +l.

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

Example: For a 3p electron: n = 3, l = 1, ml = -1, 0, or +1, ms = +1/2 or -1/2.

Emission of Light from an Atom

When electrons transition between energy levels, atoms emit or absorb light at specific wavelengths.

• Emission: Electron drops to a lower energy level, releasing a photon.

• Absorption: Electron absorbs energy and moves to a higher energy level.

Formula:

$E = h\nu = \frac{hc}{\lambda}$

where $E$ is energy, $h$ is Planck's constant, $\nu$ is frequency, $c$ is the speed of light, and $\lambda$ is wavelength.

Electron Configurations

Electron configuration describes the arrangement of electrons in an atom's orbitals.

• Electrons fill orbitals in order of increasing energy (Aufbau principle).

• 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.

Example: Oxygen (O): 1s2 2s2 2p4

Orbital Diagrams

Orbital diagrams visually represent electron configurations using boxes (orbitals) and arrows (electrons).

• Each box represents an orbital; arrows indicate electron spin.

• Unpaired electrons are shown as single arrows in a box.

Effective Nuclear Charge

Effective nuclear charge ($Z_{eff}$) is the net positive charge experienced by an electron in a multi-electron atom.

• Calculated as $Z_{eff} = Z - S$, where $Z$ is the atomic number and $S$ is the shielding constant.

• Increases across a period and decreases down a group.

Periodic Trends

Periodic trends describe how certain properties of elements change across periods and groups in the periodic table.

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

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

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

Chemical Bonding

Ionic vs Covalent Bonds

Chemical bonds form between atoms to achieve stable electron configurations. The two main types are ionic and covalent bonds.

• Ionic Bonds: Formed by the transfer of electrons from a metal to a nonmetal, resulting in oppositely charged ions.

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

Example: NaCl (ionic), H2O (covalent)

Representing Ionic and Molecular Compounds

Compounds can be represented using chemical formulas and Lewis structures.

• Chemical Formula: Indicates the types and numbers of atoms (e.g., H2O, NaCl).

• Lewis Structure: Shows the arrangement of valence electrons around atoms.

Assessment Format and Logistics

Assessment Format

• Multiple Choice

• Matching

• Short Answer

Authorized Resources

• T1-36X Pro Calculator

• 13th ed. Reference Data Card (RDC) – NO MARKINGS

• Pencil

Logistics

• Quiz will occur during the final 30 minutes of Lesson 9.

• Maximum time limit: 30 minutes.

• Lesson 9 material is included on the assessment.

• Instructors will not have extra RDCs or calculators.