JAMB 2025 Chemistry Preparation Handout For UTME Exam

1. Atomic Structure

Atomic Number and Mass Number: The atomic number represents the number of protons in the nucleus of an atom, while the mass number is the total number of protons and neutrons. These two numbers define the identity and characteristics of an atom.
Isotopes: Atoms of the same element with the same atomic number but different mass numbers due to varying numbers of neutrons. Example: Carbon-12 and Carbon-14.
Electronic Configuration: The arrangement of electrons in orbitals around the nucleus. Understanding the distribution of electrons in s, p, d, and f orbitals is critical. Use of Aufbau principle, Pauli exclusion principle, and Hund’s rule are essential in filling the orbitals.
Periodic Table and Periodicity: The periodic table is organized based on atomic number. Elements are arranged in periods and groups, with trends in atomic size, ionization energy, electron affinity, and electronegativity moving across the table. For example, atomic size decreases across a period and increases down a group.

JAMB 2025 Chemistry Preparation Handout For UTME Exam

2. Chemical Bonding

Ionic Bonding: A type of bonding where electrons are transferred from a metal to a non-metal, leading to the formation of positively charged cations and negatively charged anions. Example: NaCl.
Covalent Bonding: In covalent bonds, atoms (typically non-metals) share pairs of electrons to achieve a full outer shell. Example: H2O, where oxygen shares electrons with hydrogen atoms.
Metallic Bonding: Occurs between metal atoms where electrons are delocalized and shared among a lattice of atoms, giving metals their conductive and malleable properties.
Intermolecular Forces: Forces of attraction between molecules, including hydrogen bonding (strongest), dipole-dipole interactions, and van der Waals forces (weakest).
Shapes of Molecules: Determined by VSEPR (Valence Shell Electron Pair Repulsion) theory. Common shapes include linear, tetrahedral, trigonal planar, and bent. Example: Methane (CH4) is tetrahedral.
Hybridization: The mixing of atomic orbitals to form new hybrid orbitals suitable for bonding. Types include sp, sp2, and sp3 hybridization. For instance, the carbon atom in ethene (C2H4) is sp2 hybridized.

3. Stoichiometry and Chemical Calculations

Mole Concept: A mole represents 6.022 × 10^23 particles (Avogadro’s number). It is used to quantify atoms, ions, or molecules. Molar mass (g/mol) relates mass to the number of moles. For example, 1 mole of H2O = 18 g (2 × 1 + 16).
Empirical and Molecular Formula: Empirical formula represents the simplest whole-number ratio of elements in a compound. The molecular formula is the actual number of atoms in a molecule. For example, the empirical formula of glucose is CH2O, while its molecular formula is C6H12O6.
Balancing Chemical Equations: To balance a chemical equation, ensure that the number of atoms for each element is equal on both sides. For example: 2H2 + O2 → 2H2O.
Quantitative Analysis: Includes calculations such as determining molarity (moles per liter of solution) and using stoichiometric ratios to determine the amount of reactants or products in a reaction.

4. Acids, Bases, and Salts

Definitions: Acids are substances that release hydrogen ions (H+) in solution, while bases release hydroxide ions (OH-). According to the Bronsted-Lowry theory, acids donate protons, and bases accept protons. Lewis theory defines acids as electron pair acceptors and bases as electron pair donors.
Titration: A laboratory technique used to determine the concentration of an acid or base. This involves adding a titrant to a solution of known volume until neutralization occurs, which is often indicated by a color change using an indicator (e.g., phenolphthalein or methyl orange).
Neutralization: A reaction between an acid and a base that results in the formation of salt and water. Example: HCl + NaOH → NaCl + H2O.

5. Chemical Equilibria

Dynamic Equilibrium: Occurs when the forward and reverse reactions take place at the same rate, and the concentration of reactants and products remains constant over time. This is common in closed systems.
Le Chatelier’s Principle: States that if an external condition (such as concentration, temperature, or pressure) is changed, the equilibrium will shift to counteract that change. For example, increasing the concentration of reactants will shift the equilibrium toward product formation.
Solubility Product (Ksp): The product of the molar concentrations of the ions in a saturated solution, each raised to the power of its coefficient in the balanced equation. The common ion effect reduces the solubility of a salt in the presence of a common ion.
Buffers: Solutions that resist changes in pH when small amounts of acid or base are added. They are made from a weak acid and its conjugate base or a weak base and its conjugate acid.

6. Electrochemistry

Redox Reactions: Oxidation involves the loss of electrons, while reduction involves the gain of electrons. The substance that loses electrons is oxidized, and the one that gains electrons is reduced.
Electrolysis: The process of using electricity to drive a non-spontaneous chemical reaction. Faraday’s laws of electrolysis relate the quantity of electric charge to the amount of substance deposited or dissolved during electrolysis.
Electrochemical Cells: Galvanic (voltaic) cells convert chemical energy into electrical energy through spontaneous redox reactions. Electrolytic cells use electrical energy to drive non-spontaneous reactions. The standard electrode potential determines the cell’s ability to generate electricity.

7. Organic Chemistry

Hydrocarbons: Organic compounds made up of hydrogen and carbon atoms. Alkanes (C-C single bonds), alkenes (C=C double bonds), and alkynes (C≡C triple bonds) are important classes of hydrocarbons with various reactions (e.g., combustion, addition, substitution).
Functional Groups: Specific groups of atoms within molecules that determine the characteristics of chemical reactions. Examples: -OH (alcohol), -COOH (carboxylic acid), -NH2 (amine).
Isomerism: Compounds with the same molecular formula but different structures. Types include structural isomerism (e.g., chain isomerism) and stereoisomerism (e.g., geometric isomerism).
Polymerization: The process of joining small molecules (monomers) to form large molecules (polymers). Types include addition polymerization (e.g., polyethylene) and condensation polymerization (e.g., nylon).

8. Thermodynamics

Enthalpy (ΔH): The heat content of a system. Reactions can be exothermic (release heat) or endothermic (absorb heat). Enthalpy changes can be calculated using bond energies or from experimental data (calorimetry).
Entropy (ΔS): A measure of the disorder or randomness in a system. Processes tend to proceed in the direction of increasing entropy.
Gibbs Free Energy (ΔG): Determines the spontaneity of a reaction. ΔG = ΔH – TΔS. If ΔG is negative, the reaction is spontaneous.

9. Kinetics

Rate of Reaction: The speed at which reactants are converted into products. Factors influencing rate include temperature, concentration, particle size, and the presence of catalysts. The Arrhenius equation relates the rate constant to temperature.
Collision Theory: Reactants must collide with proper orientation and sufficient energy to overcome the activation energy barrier for a reaction to occur.
Activation Energy: The minimum energy required for a reaction to proceed. Catalysts lower the activation energy, increasing the rate of reaction.

10. Separation Techniques

Filtration: A method for separating an insoluble solid from a liquid using a filter.
Distillation: Separates liquids based on differences in boiling points. Used for purifying liquids or separating mixtures of liquids.
Chromatography: A technique for separating the components of a mixture based on their different affinities to a stationary phase and a mobile phase. Common types include paper and thin-layer chromatography.
Crystallization: The formation of pure solid particles from a solution as it cools and the solubility of the solute decreases.