Inorganic Chemistry Module Notes

Module 1: Overview to General Chemistry

Matter can be classified by composition, properties, and the kinds of changes it can undergo. According to composition, matter is divided into pure substances (elements and compounds) and mixtures (homogeneous solutions and heterogeneous mixtures). A molecule of oxygen, an element, is a pure substance, while water is a compound, and air is a mixture containing several gases. This classification helps explain why substances have particular properties and how they react. In terms of properties, matter exhibits physical properties such as color, density, melting point, and solubility, which can be observed without changing the substance, and chemical properties such as reactivity or flammability, which describe how a substance may change into new substances under certain conditions. Changes to matter are categorized as physical changes (phase transitions like melting or dissolving without altering the substance’s identity), chemical changes (reactions that form new substances), and nuclear changes (rearrangements in the nucleus, such as radioactive decay, fission, or fusion) that often involve energy changes at the atomic level. To handle very large or very small numbers that appear in chemistry, scientists use scientific notation, where a number is written as a \times 10^{n} with 1 \le a < 10 and n an integer, for example 6.022 \times 10^{23} (Avogadro’s number) or 4.5 \times 10^{-4}. This notation simplifies calculations and keeps track of significant figures in measurements.

Module 2: Atoms and Elements

In this module, the study begins with the evolution of atomic theories and the fundamental structure of atoms. Historical atomic theories include Dalton’s atomic theory (matter consists of indivisible atoms), Thomson’s plum pudding model (electrons embedded in a positively charged sphere), Rutherford’s nuclear model (a dense nucleus with electrons surrounding it), Bohr’s model (electrons orbit specific energy levels), and the modern quantum mechanical model (electrons described by orbitals and probability distributions). The atomic structure comprises a dense nucleus containing protons (p+) and neutrons (n0), with electrons (e−) surrounding the nucleus in electron clouds. The nucleus gives the atom its mass, while the electrons determine its chemical behavior. The mass number is denoted by A = Z + N, where Z is the atomic number (number of protons) and N is the number of neutrons. Atoms can gain or lose electrons to form ions, resulting in a net charge. Ions are categorized as cations (positive charge, formed by loss of electrons) and anions (negative charge, formed by gain of electrons).

Module 3: Electronic Structure

This module covers the arrangement of electrons around the nucleus and how it relates to periodic properties. The Periodic Table groups elements into families (groups) and rows (periods), with elements in the same group sharing similar valence electron configurations and chemical properties. The electronic structure of atoms is described by quantum numbers and electron configurations. Quantum numbers describe the allowed states of electrons: the principal quantum number n (energy level), the orbital angular momentum quantum number l (subshel level, where l = 0,1,2,\ldots, n-1), the magnetic quantum number ml (orbital orientation, where ml = -l, -l+1, \ldots, +l), and the spin quantum number ms (spin orientation, ms = \pm \frac{1}{2}). The Pauli exclusion principle, Hund’s rule, and Aufbau principle guide electron configurations, for example the ground-state configuration of oxygen is 1s^2 2s^2 2p^4. Periodic trends describe how properties change across periods and within groups: atomic radius generally increases down a group and decreases across a period; ionization energy and electron affinity exhibit characteristic trends; electronegativity tends to increase across a period and decrease down a group. General rules of electron configuration involve filling lowest-energy orbitals first, obeying the order dictated by the Aufbau principle, while ensuring that no two electrons in the same atom have identical quantum numbers (Pauli exclusion) and maximizing unpaired electrons in degenerate orbitals (Hund’s rule).

Module 4: Classifications, Naming and Writing of Formulas

This module focuses on naming compounds and writing their formulas. Compounds are classified as ionic (typically metal with a nonmetal) or covalent (two nonmetals). Ionic compounds are named by giving the name of the cation followed by the name of the anion; for transition metals, oxidation states are indicated with a Roman numeral (stock system), e.g., iron(III) chloride. Polyatomic ions (such as sulfate, NO₃⁻, NH₄⁺) are common species with fixed formulas. Writing formulas involves balancing the charges to obtain a neutral overall compound (e.g., Na⁺ and Cl⁻ form NaCl; Ca²⁺ with SO₄²⁻ forms CaSO₄). Covalent compound naming uses prefixes to indicate the number of atoms of each element (e.g., CO₂ is carbon dioxide, P₄O₁₀ is tetraphosphorus decaoxide). The general approach includes recognizing ionic versus covalent bonding, applying appropriate naming conventions, and constructing correct chemical formulas from names or vice versa.

Module 5: Stoichiometry

Stoichiometry connects the macroscopic world of masses and moles to the microscopic world of atoms and molecules. Atomic and molecular masses are used to determine molar masses: the molar mass M of a substance is numerically equal to its relative atomic or molecular mass in g/mol, and it serves as a bridge between grams and moles via n = \dfrac{m}{M}, where n is the amount in moles and m is the mass in grams. Avogadro’s number, NA = 6.022 \times 10^{23} \, \text{mol}^{-1}, connects the number of particles to moles: N = n \times NA. Percent composition expresses the fraction of each element’s mass in a compound: \%\text{ composition of element i} = \left( \dfrac{mi}{m{\text{compound}}} \right) \times 100\%. The empirical formula gives the simplest whole-number ratio of atoms in a compound, derived from percent composition by converting percentages to moles and simplifying the mole ratio; the molecular formula gives the actual number of atoms in a molecule and is related to the empirical formula by the ratio \dfrac{M{\text{actual}}}{M{\text{empirical}}} = n \in \mathbb{N}, with the actual formula being (\text{empirical formula})_n.

In reactions, the concept of limiting and excess reactants determines the theoretical yield. The limiting reactant is the one consumed first, limiting the amount of product formed, while any remaining reagent is in excess. To predict product amounts, one converts the given masses or moles of reactants to moles using their molar masses, applies the balanced chemical equation, identifies the limiting reagent, and computes the theoretical yield of products using the stoichiometric coefficients. The general types of chemical reactions include synthesis (A + B → AB), decomposition (AB → A + B), single-displacement (AB + C → AC + B), double-displacement (AB + CD → AD + CB), and combustion (fuel + O₂ → CO₂ + H₂O). These concepts are foundational for quantitative analysis in chemistry and have broad real-world relevance in fields ranging from materials science to pharmacology.