Atoms, Ions, and Chemical Reactions: Core Concepts and Solubility Guidelines

Chapter 2: Atoms and the Periodic Table

Atomic Theory

Atomic theory forms the foundation of modern chemistry, describing the nature and behavior of matter at the atomic level.

• All matter is composed of atoms, which are indivisible particles.

• Atoms of one element cannot change into atoms of another element during chemical reactions.

• Atoms of the same element have the same number of protons and electrons, determining their chemical behavior.

• Compounds are formed by the chemical combination of two or more elements in a specific ratio.

Atomic Symbols and Numbers

• Atomic symbol (X): The one- or two-letter abbreviation for an element.

• Atomic number (Z): Number of protons in the nucleus; also equals the number of electrons in a neutral atom.

• Atomic mass number (A): Total number of protons and neutrons in the nucleus.

• Number of neutrons (N): $N = A - Z$

Isotopes and Ions

• Isotopes: Atoms of the same element with the same number of protons but different numbers of neutrons.

• Ions: Atoms or molecules with a net electric charge due to the loss or gain of electrons.

• For ions: Number of electrons = Number of protons – charge

Atomic Weight

The atomic weight is the weighted average of the isotopic masses of an element.

• Calculated as: $\text{Atomic Weight} = \sum (\text{isotopic abundance} \times \text{isotopic mass})$

Electron Configuration and Periodic Table

• Electron configuration: The arrangement of electrons in an atom or ion.

• Element blocks: s, p, d, f blocks correspond to the type of atomic orbital being filled.

• Major group elements: Groups 1A, 2A, 7A, and 8A are especially important for chemical reactivity.

Chapter 3: Ionic Compounds

Key Terms and Concepts

• Molecule: Two or more atoms bonded together.

• Diatomic molecule: Molecule consisting of two atoms (e.g., O2).

• Polyatomic molecule: Molecule with more than two atoms.

• Cation: Positively charged ion.

• Anion: Negatively charged ion.

• Monatomic ion: Ion formed from a single atom.

• Polyatomic ion: Ion composed of multiple atoms.

• Transition metal ion: Ion formed from a transition metal, often with variable charge.

Common Ions and Acids

• Know the names and charges of common ions (e.g., Na+, Cl-, SO42-).

• Know the names of common acids (e.g., HCl, H2SO4).

Constructing and Naming Ionic Compounds

• Combine cations and anions to form neutral compounds.

• Write formulas for ionic compounds (monatomic, polyatomic, transition metal).

• Name ionic compounds according to standard nomenclature rules.

Periodic Table and Ion Formation

• Major group elements (1A, 2A, 3A, 5A, 6A, 7A) form predictable ions based on their group number.

Chapter 4: Molecular Compounds and Bonding

Valence Electrons and Bonding

• Number of valence electrons: Equals the group number for main group elements.

• Valence electrons participate in chemical bonding.

Types of Bonds

• Ionic bond: Electrostatic attraction between cations and anions.

• Covalent bond: Two or more atoms share electrons.

• Polar covalent bond: Covalent bond between atoms with different electronegativities.

Electronegativity and Bond Polarity

• Electronegativity (EN): Ability of an atom to attract electrons in a bond.

• EN increases from bottom to top and left to right on the periodic table.

• Common EN values: F (4.0), O/Cl (3.5), N (3.0), C (2.5), H (2.1), Br (2.8), others below 2.5.

• Bond type by ΔEN:

  • ΔEN > 2: Ionic

  • 0.5 < ΔEN < 1.9: Polar covalent

  • ΔEN < 0.4: Nonpolar covalent

Lewis Structures

• Count valence electrons.

• Draw skeleton structure.

• Deduct 2 electrons for each bond.

• Add lone pairs to complete octets.

• If octet is not achieved, create multiple bonds as needed.

Molecular Geometry and Polarity

• Electron group = number of outer atoms + number of lone pairs.

• Geometry and bond angles:

  • 2 groups: Linear, 180°, nonpolar

  • 3 groups: Trigonal planar, 120°, nonpolar; Bent, 120°, polar

  • 4 groups: Tetrahedral, 109°, nonpolar; Trigonal pyramidal, 109°, polar; Bent, 109°, polar

• Polarity depends on both bond polarity and molecular geometry.

Chapter 5: Classification & Balancing of Chemical Reactions

Balancing Chemical Equations

Balancing chemical equations ensures the law of conservation of mass is obeyed in chemical reactions.

• List each element on both sides of the equation.

• Balance polyatomic ions or compounds first.

• Balance atoms that appear only once on each side.

• Balance hydrogens, then oxygens.

• Double-check atom counts on both sides.

• Use only coefficients; do not change subscripts.

Classifying Reactions

• Redox reactions: Involve transfer of electrons.

• Acid/base reactions: Involve transfer of protons (H+).

• Precipitation reactions: Formation of an insoluble solid from soluble reactants.

Solubility Guidelines for Ionic Compounds in Water

Solubility rules help predict whether a compound will dissolve in water or form a precipitate.

Soluble	Exceptions
Ammonium compounds (NH4+)	None
Lithium compounds (Li+)	None
Sodium compounds (Na+)	None
Potassium compounds (K+)	None
Nitrates (NO3-)	None
Perchlorates (ClO4-)	None
Acetates (CH3COO-)	None
Chlorides (Cl-)	Ag+, Hg22+, and Pb2+ compounds
Bromides (Br-)	Ag+, Hg22+, and Pb2+ compounds
Iodides (I-)	Ag+, Hg22+, and Pb2+ compounds
Sulfates (SO42-)	Ba2+, Hg22+, and Pb2+ compounds

Example: Predicting Precipitation

• Mixing solutions of NaCl and AgNO3 forms a precipitate of AgCl, since AgCl is insoluble.

Additional info: The solubility table is essential for predicting the outcomes of double displacement reactions and for identifying the formation of precipitates in aqueous solutions.