General Chemistry: Atomic Structure, Bonding, and Molecular Geometry Study Guide

Atomic Structure and Electron Configuration

Valence Electrons

Valence electrons are the outermost electrons of an atom and are crucial in determining chemical reactivity and bonding.

• Definition: Valence electrons are electrons in the highest occupied energy level of an atom.

• Example: Oxygen has 6 valence electrons.

• Application: The number of valence electrons affects the formation of ions and chemical bonds.

Electron Configuration and Ion Formation

Atoms gain or lose electrons to achieve a stable electron configuration, often resembling the nearest noble gas.

• Example: Sulfur (S) with 16 protons and 18 electrons forms the sulfide ion, $\mathrm{S^{2-}}$.

• Octet Rule: Atoms tend to have eight electrons in their valence shell.

Ionic and Covalent Compounds

Formation of Ionic Compounds

Ionic compounds are formed by the transfer of electrons from metals to nonmetals, resulting in positive and negative ions.

• Example: Potassium (K) and oxygen (O) form potassium oxide, $\mathrm{K_2O}$.

• Formula: $\mathrm{K^+}$ and $\mathrm{O^{2-}}$ combine in a 2:1 ratio.

Naming Ionic Compounds

The name of an ionic compound is based on the cation and anion present.

• Example: $\mathrm{PbO_2}$ is called lead(IV) oxide.

• Polyatomic Ions: Compounds like $\mathrm{Na_3PO_4}$ contain polyatomic ions such as phosphate ($\mathrm{PO_4^{3-}}$).

Polar and Nonpolar Compounds

Polarity depends on the difference in electronegativity between atoms and the molecular geometry.

• Polar Molecules: Have an uneven distribution of charge (e.g., $\mathrm{SO_2}$).

• Nonpolar Molecules: Have an even distribution of charge (e.g., $\mathrm{CO_2}$).

Chemical Bonding and Molecular Geometry

Types of Chemical Bonds

Chemical bonds are the forces holding atoms together in compounds.

• Ionic Bonds: Formed by the transfer of electrons.

• Covalent Bonds: Formed by the sharing of electrons.

• Polar Covalent Bonds: Electrons are shared unequally.

Intermolecular Forces

Intermolecular forces are attractions between molecules, affecting physical properties.

• Dipole-Dipole: Attraction between polar molecules.

• Hydrogen Bonds: Strong dipole-dipole interactions involving H bonded to N, O, or F.

• Dispersion Forces: Present in all molecules, especially nonpolar ones.

Molecular Geometry

The shape of a molecule is determined by the arrangement of atoms and electron pairs around the central atom.

• VSEPR Theory: Valence Shell Electron Pair Repulsion theory predicts molecular shapes.

• Examples:

  • Tetrahedral: e.g., $\mathrm{CH_4}$

  • Trigonal Planar: e.g., $\mathrm{CO_3^{2-}}$

  • Bent: e.g., $\mathrm{SeO_2}$

Table: Electron Geometry and Bonding in Selected Molecules

Molecule	Electron Geometry	Molecular Shape	$\sigma$ Bonds	$\pi$ Bonds
$\mathrm{NH_3}$	Tetrahedral	Trigonal pyramidal	3	0
$\mathrm{CH_3O^-}$	Trigonal planar	Trigonal planar	3	1
$\mathrm{SeO_2}$	Trigonal planar	Bent	2	1

Polyatomic Ions and Nomenclature

Polyatomic Ions

Polyatomic ions are charged species composed of two or more atoms covalently bonded.

• Examples: $\mathrm{PO_4^{3-}}$ (phosphate), $\mathrm{NO_3^-}$ (nitrate)

• Application: Used in naming compounds such as sodium phosphate ($\mathrm{Na_3PO_4}$).

Naming Rules for Ionic Compounds

When naming ionic compounds, the cation is named first, followed by the anion. Prefixes are not used for ionic compounds.

• Example: $\mathrm{Na_3PO_4}$ is sodium phosphate, not trisodium monophosphate.

Special Topics

Diatomic Elements

Certain elements exist naturally as diatomic molecules.

• Examples: $\mathrm{H_2}$, $\mathrm{N_2}$, $\mathrm{O_2}$, $\mathrm{F_2}$, $\mathrm{Cl_2}$, $\mathrm{Br_2}$, $\mathrm{I_2}$

• Non-diatomic: Neon ($\mathrm{Ne}$) does not exist as a diatomic molecule.

Transition Metals and Variable Oxidation States

Transition metals can form cations with different charges.

• Example: Vanadium can form $\mathrm{V^{2+}}$, $\mathrm{V^{5+}}$, etc.

• Application: Naming requires Roman numerals to indicate charge (e.g., vanadium(II) chloride).

Practice Problems and Applications

Lewis Structures

Lewis structures represent the arrangement of electrons in a molecule.

• Steps:

  1. Count total valence electrons.

  2. Arrange atoms and connect with single bonds.

  3. Distribute remaining electrons to complete octets.

• Example: $\mathrm{NH_3}$, $\mathrm{CH_3O^-}$, $\mathrm{SeO_2}$ (see table above for geometry and bonds).

Electron Geometry and Molecular Shape

Electron geometry considers all electron pairs, while molecular shape considers only bonded atoms.

• Example: $\mathrm{NH_3}$ has tetrahedral electron geometry but trigonal pyramidal molecular shape.

Bond Counting

Single bonds are $\sigma$ bonds; double and triple bonds include $\pi$ bonds.

• Example: $\mathrm{CH_3O^-}$ has 3 $\sigma$ bonds and 1 $\pi$ bond.

Summary Table: Key Concepts

Concept	Definition	Example
Valence Electrons	Electrons in the outermost shell	Oxygen: 6 valence electrons
Ionic Compound	Formed by transfer of electrons	$\mathrm{K_2O}$
Covalent Compound	Formed by sharing electrons	$\mathrm{H_2O}$
Polyatomic Ion	Charged group of covalently bonded atoms	$\mathrm{PO_4^{3-}}$
Lewis Structure	Diagram showing electron arrangement	$\mathrm{NH_3}$

Additional info: Some context and explanations have been expanded for clarity and completeness, including definitions, examples, and tables summarizing key concepts.