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Chapter 5 - Electrons in Atoms

• Wave Relationship Equation

  • Calculate the wavelength or frequency of a wave using the wave relationship equation:

  • Formula: v = f \lambda

    • where:

    • $v$ = speed of light

    • $f$ = frequency

    • $\lambda$ = wavelength

• Energy of a Photon

  • Calculate the energy of a photon using wavelength or frequency:

  • Formula: E = hf

    • where:

    • $E$ = energy of a photon

    • $h$ = Planck's constant ($6.626 \times 10^{-34} \, \text{Js}$)

    • $f$ = frequency

  • Alternatively, using wavelength:

    • E = \frac{hc}{\lambda}

    • where:

      • $c$ = speed of light ($3.00 \times 10^{8} \, \text{m/s}$)

      • $\lambda$ = wavelength

• Conversions

  • Know how to convert between:

  • Meters and nanometers:

    • 1 meter = $1 \times 10^{9}$ nanometers

  • Joules and kilojoules:

    • 1 Joule = $0.001$ kilojoules

• Electron Behavior

  • Use the terms ground state and excited state:

  • Ground State: The lowest energy state of an atom, where electrons are in the closest orbitals to the nucleus.

  • Excited State: A higher energy state where electrons have absorbed energy and are in higher orbitals than their ground state.

• Orbitals

  • Know the different orbitals, their designations, shape, and number of energy levels for each:

  • s-orbital:

    • Shape: Spherical

    • Number of energy levels: 1

  • p-orbital:

    • Shape: Dumbbell

    • Number of energy levels: 2

  • d-orbital:

    • Shape: Various (cloverleaf)

    • Number of energy levels: 3

  • f-orbital:

    • Shape: Complex

    • Number of energy levels: 4

• Electron Configuration

  • Predict and write the electron configuration of different elements:

  • Follow the Aufbau principle, Pauli exclusion principle, and Hund's rule to fill orbitals.

    • Example:

    • Oxygen: 1s^2 2s^2 2p^4

• Energy Levels and Sublevels

  • Know the proper order of the energy levels and sublevels:

  • Order:

    • $1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p < 6s < 4f < 5d < 6p$

• Isoelectronic Ions

  • Identify isoelectronic ions and compare their properties:

  • Isoelectronic ions have the same electron configuration but different nuclear charges.

Chapter 6 - The Periodic Table and Periodic Law

• Periodic Table Organization

  • Identify groups/families and periods using IUPAC nomenclature.

• Specialized Element Categories

  • Identify elements that fit into the following specialized categories:

  • Representative Elements

  • Alkali Metals

  • Alkaline Earth Metals

  • Halogens

  • Noble Gases

  • Transition Metals

  • Lanthanides

  • Actinides

• Element Classifications

  • Identify the three classifications of elements:

  • Metals

  • Metalloids

  • Nonmetals

• Electron Configuration and Position

  • Relate the position on the periodic table to the electron configuration of an element:

  • Valence electrons correspond to the group number.

• Periodic and Group Trends

  • Predict trends on the periodic table for:

  • Atomic Radius: Generally increases down a group and decreases across a period.

  • Ionic Radius: Cations are smaller than the parent atom; anions are larger.

  • First Ionization Energy: Generally increases across a period and decreases down a group.

  • Electronegativity: Generally increases across a period and decreases down a group.

Chapter 7 - Ionic Compounds and Metals

• Valence Electrons and Ion Charges

  • Identify the number of valence electrons and common ion charges for all representative elements.

  • Know how to identify the charges for transition metals, including exceptions.

• Octet Rule

  • Use the octet rule to explain the formation of common ion charges:

  • Cations (positive) form by losing electrons to achieve a full outer shell.

  • Anions (negative) form by gaining electrons for the same reason.

• Compound Types

  • Identify ionic compounds (formed from metals and nonmetals) versus molecular compounds (formed from nonmetals).

• Naming Ionic Compounds

  • Know how to name ionic compounds and write formulas:

  • Example: Sodium chloride for NaCl.

• Common Polyatomic Ions

  • Know your common polyatomic ions, their formulas, and charges:

  • Example: Hydroxide OH^-, Sulfate SO_4^{2-}.

Chapter 8 - Covalent Bonding

• Lewis Dot Structures

  • Draw the Lewis dot structure for a molecule to represent bonding between atoms.

• Molecular Polarity

  • Determine whether a molecule is polar or nonpolar based on its dipole moments:

  • Polar: Unequal sharing of electrons leading to a dipole moment.

  • Nonpolar: Equal sharing of electrons.

• VSEPR Notation

  • Learn how to determine VSEPR notation (AXE) to predict molecular shape:

  • A = central atom, X = bonded atoms, E = lone pairs.

• Bond Angles and Hybridization

  • Know the different bond angles for each molecular geometry and associated hybrid orbitals:

  • Example:

    • Linear: 180° (sp)

    • Trigonal Planar: 120° (sp²)

    • Tetrahedral: 109.5° (sp³)

• Nomenclature of Molecular Compounds

  • Write the formula/name for molecular compounds using numerical prefixes:

  • Example: CO = carbon monoxide.

• Common Acids

  • Know the formulas for 6 common acids:

  • HCl (hydrochloric acid), H₂SO₄ (sulfuric acid), HNO₃ (nitric acid), etc.

Chapter 9 - Chemical Reactions

• Writing and Balancing Equations

  • Know how to write, balance, and classify chemical equations.

• Aqueous Ionic Compounds

  • Be familiar with how aqueous ionic compounds dissociate, including complete ionic and net ionic equations.