An Expert-Level Report on Inorganic Chemistry for the IMAT

This part establishes the fundamental framework of inorganic chemistry: the periodic table. It explores its structure not as a mere chart, but as a direct visual representation of electron configuration, which is the ultimate determinant of an element's chemical properties. Understanding this link is the first step to predictive chemical thinking.

Section 1.1: The Periodic Table: A Chemist's Map

The periodic table of elements is the single most important organizational tool in chemistry. It arranges elements by increasing atomic number () and groups atoms with similar properties in the same vertical column (group). The seven horizontal rows are known as periods.

Elements in the same group possess the same number of valence electrons, leading to similar chemical properties. Key groups include:

• Group 1: The Alkali Metals
• Group 2: The Alkaline Earth Metals
• Group 17: The Halogens
• Group 18: The Noble Gases

The table's structure is a map of how electrons fill atomic orbitals, divided into the s-block, p-block, d-block (transition metals), and f-block. Noble gas notation (e.g., for Sodium) is used to efficiently represent electron configurations by focusing on the valence electrons.

Section 1.2: Periodic Trends in Elemental Properties

Predictable patterns in properties arise from the interplay between effective nuclear charge (), electron shells (n), and electron shielding.

• Atomic Radius: Decreases across a period (due to increasing ). Increases down a group (due to adding new electron shells).
• Ionization Energy (IE): The energy required to remove an electron. Increases across a period. Decreases down a group.
• Electronegativity: An atom's ability to attract electrons in a bond. Increases across a period. Decreases down a group. Fluorine is the most electronegative element.

📸 Source/Description: This image provides a visual summary of the major periodic trends. Atomic radius increases down and to the left. Ionization energy and electronegativity increase up and to the right.

Section 2.2: Oxides: A Study in Acidity and Basicity

Oxides are binary compounds containing oxygen. Their naming and properties depend on the element bonded to oxygen.

• Nomenclature: For metals with fixed charges, use the metal name (e.g., CaO, calcium oxide). For metals with variable charges, use a Roman numeral (Stock system), e.g., FeO is iron(II) oxide. For nonmetal oxides, use Greek prefixes (e.g., , sulfur trioxide).
• Properties: Basic oxides (from metals, e.g., ) react with water to form bases. Acidic oxides (from nonmetals, e.g., ) react with water to form acids. Amphoteric oxides (e.g., ) can act as either an acid or a base.

Section 2.3: Hydroxides, 2.4: Acids, & 2.5: Salts

Hydroxides: Contain the ion. Named as Cation + "hydroxide" (e.g., NaOH, sodium hydroxide; Fe(OH)₂, iron(II) hydroxide).

Acids: Release ions in solution. Binary acids use the prefix "hydro-" and suffix "-ic acid" (e.g., HCl(aq), hydrochloric acid). Oxyacids are named based on their polyatomic anion: "-ate" anions become "-ic" acids (e.g., , sulfuric acid), and "-ite" anions become "-ous" acids (e.g., , nitrous acid).

Salts: Ionic compounds formed from an acid-base neutralization. Named as Cation + Anion (e.g., Cu(NO₃)₂, copper(II) nitrate).

Section 3.1: The Representative Non-Metals

• Hydrogen (H): Unique with a configuration. Highly flammable, reacting with oxygen to form water ().
• The Noble Gases (Group 18): Exceptionally stable due to full valence shells (). They are monatomic, colorless, and largely unreactive gases used in lighting (Neon) and as inert atmospheres (Argon).
• The Halogens (Group 17): Highly reactive nonmetals () that exist as diatomic molecules (). They are powerful oxidizing agents, and their reactivity decreases down the group.

Section 3.2: The Active Metals

• The Alkali Metals (Group 1): Soft, highly reactive metals with a single valence electron (). They readily lose this electron, and their reactivity increases down the group. They react violently with water: .
• The Alkaline Earth Metals (Group 2): Reactive metals with two valence electrons (). They are less reactive than alkali metals but more reactive than most other metals. Their reactivity also increases down the group.

Section 3.3: The Transition Metals (d-block)

The chemistry of transition metals is defined by their partially filled d subshells. This leads to their characteristic properties:

• Variable Oxidation States: They can form multiple stable ions (e.g., and ) because the energies of the ns and (n-1)d orbitals are very close.
• Formation of Coloured Ions: Their compounds are often colored because electrons can absorb visible light to move between split d-orbital energy levels (d-d transition).
• Catalytic Activity: Their ability to change oxidation states and adsorb reactants makes them excellent catalysts (e.g., Fe in the Haber Process).

Section 4.1: Acid-Base Neutralization Reactions

An acid and a base react to form a salt and water. The net ionic equation for any strong acid-strong base reaction is:

The pH of the final solution depends on the strengths of the acid and base used.

Section 4.2: Precipitation Reactions

A precipitate (an insoluble solid) forms when two soluble ionic compounds are mixed. Predicting these reactions requires knowledge of the solubility rules.

Section 4.3: Oxidation-Reduction (Redox) Reactions

Redox reactions involve the transfer of electrons, resulting in a change in oxidation state. The mnemonic OIL RIG (Oxidation Is Loss, Reduction Is Gain) is useful. To determine if a redox reaction has occurred, one must assign oxidation states to each atom.

Example reaction:

• Zinc is oxidized (0 to +2), so it is the reducing agent.
• Copper is reduced (+2 to 0), so it is the oxidizing agent.