Exam 4 Review: Periodic Properties, Lewis Model, and Molecular Shapes (Ch. 9-11)

Chapter 9. Periodic Properties of the Elements

9.1 The Development of the Periodic Table

The periodic table organizes elements based on recurring chemical properties, allowing scientists to predict element behavior. Its development was crucial for understanding atomic structure and chemical reactivity.

• Discovery of New Elements: The 19th-century discovery of new elements led to the need for systematic classification.

• Mendeleev's Contribution: Dmitri Mendeleev arranged elements by increasing atomic mass, grouping elements with similar properties.

• Periodic Law: Properties of elements recur periodically when arranged by atomic number.

9.3 Electron Configurations: How Electrons Occupy Orbitals

Electron configuration describes the arrangement of electrons in an atom, which determines chemical properties and reactivity.

• Electron Shells and Subshells: 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 quantum numbers.

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

• Order of Filling: Orbitals fill in the order: $1s \rightarrow 2s \rightarrow 2p \rightarrow 3s \rightarrow 3p \rightarrow 4s \rightarrow 3d$.

9.4 Electron Configurations, Valence Electrons, and the Periodic Table

Valence electrons are the outermost electrons and determine chemical bonding and reactivity.

• Valence Electrons: Electrons in the highest principal energy level.

• Periodic Table Prediction: The group number often indicates the number of valence electrons.

• Electron Configuration Notation: Example for sodium: $1s^2 2s^2 2p^6 3s^1$.

9.5 The Explanatory Power of the Quantum-Mechanical Model

The quantum-mechanical model explains periodic trends based on electron arrangement and energy levels.

• Valence Electrons: Elements in the same group have the same number of valence electrons, leading to similar chemical properties.

9.6 Periodic Trends in the Sizes of Atoms and Effective Nuclear Charge

Periodic trends describe how atomic properties change across periods and groups.

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

• Effective Nuclear Charge ($Z_{eff}$): The net positive charge experienced by valence electrons; increases across a period.

• Transition Metals: Atomic radii remain approximately constant across each period.

9.7 Ions: Electron Configurations, Magnetic Properties, Ionic Radii, and Ionization Energy

Ions are atoms or molecules with a net electric charge due to loss or gain of electrons.

• Ionization Energy: Energy required to remove an electron from an atom; increases across a period, decreases down a group.

• Electron Affinity: Energy change when an atom gains an electron; generally more negative across a period.

• Ionic Radius: Cations are smaller than their parent atoms; anions are larger.

• Magnetic Properties: Atoms with unpaired electrons are paramagnetic; those with all electrons paired are diamagnetic.

9.8 Electron Affinities and Metallic Character

Electron affinity and metallic character help classify elements and predict their chemical behavior.

• Metallic Character: Increases down a group, decreases across a period.

9.9 Periodic Trends Summary

Periodic trends summarize the recurring patterns in atomic and chemical properties across the periodic table.

• Groups: Group 1A (alkali metals), Group 2A (alkaline earth metals), Group 7A (halogens), Group 8A (noble gases).

Chapter 10. Chemical Bonding I: Lewis Model

10.2 Types of Chemical Bonding

Chemical bonds form when atoms share or transfer electrons, lowering the potential energy of the system.

• Ionic Bond: Transfer of electrons between metals and nonmetals.

• Covalent Bond: Sharing of electrons between nonmetals.

• Metallic Bond: Delocalized electrons among metal atoms.

10.3 Representing Valence Electrons with Dots

Lewis dot structures visually represent valence electrons and help predict bonding.

• Lewis Symbols: Dots around element symbols represent valence electrons.

• Octet Rule: Atoms tend to gain, lose, or share electrons to achieve eight valence electrons.

10.4 Ionic Bonding: Lattice Structures and Lattice Energy

Ionic compounds form crystalline lattices, and lattice energy measures the strength of these interactions.

• Lattice Energy ($U$): Energy released when gaseous ions form an ionic solid.

• Born-Haber Cycle: Used to calculate lattice energy.

10.5 Covalent Bonding: Lewis Structures

Covalent bonds involve the sharing of electron pairs between atoms.

• Single, Double, Triple Bonds: One, two, or three shared pairs of electrons.

• Resonance Structures: Multiple valid Lewis structures for a molecule.

10.6 Electronegativity and Bond Polarity

Electronegativity measures an atom's ability to attract electrons in a bond, affecting bond polarity.

• Bond Polarity: Difference in electronegativity leads to polar covalent bonds.

• Dipole Moment ($\mu$): $\mu = q \times r$ (charge times distance).

10.7 Lewis Structures of Molecular Compounds and Polyatomic Ions

Lewis structures help predict molecular geometry and reactivity for compounds and ions.

• Polyatomic Ions: Charged species composed of multiple atoms.

10.8 Resonance and Formal Charge

Formal charge helps determine the most stable Lewis structure for a molecule.

• Formal Charge Formula: $\text{Formal Charge} = \text{Valence Electrons} - \text{Nonbonding Electrons} - \frac{1}{2} \times \text{Bonding Electrons}$

10.9 Exceptions to the Octet Rule

Some molecules have odd numbers of electrons, incomplete octets, or expanded octets.

• Odd-Electron Species: Molecules with an unpaired electron.

• Expanded Octets: Elements in period 3 or higher can have more than eight electrons.

10.10 Bond Energies and Bond Lengths

Bond energy is the energy required to break a bond; bond length is the distance between nuclei.

• Bond Energy: Stronger bonds have higher bond energies and shorter bond lengths.

10.11 Bonding in Metals: The Electron Sea Model

The electron sea model explains metallic bonding and the properties of metals.

• Delocalized Electrons: Electrons move freely among metal atoms, leading to conductivity and malleability.

Chapter 11. Chemical Bonding II: Molecular Shapes, Valence Bond Theory, and Molecular Orbital Theory

11.2 VSEPR Theory: The Five Basic Shapes

Valence Shell Electron Pair Repulsion (VSEPR) theory predicts molecular shapes based on electron pair repulsion.

• Basic Shapes: Linear, trigonal planar, tetrahedral, trigonal bipyramidal, octahedral.

• Bond Angles: Determined by the number of electron groups.

11.3 VSEPR Theory: The Effect of Lone Pairs

Lone pairs affect molecular geometry by repelling bonding pairs, altering bond angles.

• Electron Geometry vs. Molecular Geometry: Electron geometry considers all electron groups; molecular geometry considers only atoms.

11.4 VSEPR Theory: Predicting Molecular Geometries

VSEPR theory allows prediction of molecular shapes for molecules with different numbers of electron groups and lone pairs.

• Examples: Water (bent), ammonia (trigonal pyramidal), methane (tetrahedral).

11.5 Molecular Shape and Polarity

Molecular shape and bond polarity determine whether a molecule is polar or nonpolar.

• Polarity: Polar molecules have an uneven distribution of charge.

11.6 Valence Bond Theory: Orbital Overlap as a Covalent Bond

Valence bond theory explains covalent bonding as the overlap of atomic orbitals.

• Orbital Overlap: Covalent bonds form when atomic orbitals overlap, sharing electrons.

11.7 Valence Bond Theory: Hybridization of Atomic Orbitals

Hybridization describes the mixing of atomic orbitals to form new, equivalent orbitals for bonding.

• Common Types: $sp$, $sp^2$, $sp^3$ hybridization.

• Expanded Octets: $sp^3d$, $sp^3d^2$ hybridization for elements in period 3 or higher.

11.8 Molecular Orbital Theory: Electron Delocalization

Molecular orbital theory describes electrons as delocalized over the entire molecule.

• Linear Combination of Atomic Orbitals (LCAO): Atomic orbitals combine to form molecular orbitals.

• Bonding and Antibonding Orbitals: Bonding orbitals increase electron density between nuclei; antibonding orbitals decrease it.

• Molecular Orbital Diagrams: Used to predict magnetic properties and bond order.

Materials to Memorize for the Exam

• Periodic trends

• Electron filling order (and electron removal order)

• Everything pertinent to drawing a Lewis structure (formal charge, electron count, hybridization, polarity, geometries, bond angles)

• How to calculate $\Delta EN$ (difference in electronegativity)