Chemistry

Module 2: Foundations in Chemistry

1. Atoms, Ions, and Compounds

• Atomic Structure:

• Atoms consist of protons, neutrons, and electrons. Protons (+1 charge) and neutrons (neutral) are in the nucleus, while electrons (-1 charge) orbit in shells.

• Isotopes have the same number of protons but different numbers of neutrons. Chemical properties remain the same, but physical properties differ slightly.

• Relative Atomic Mass (Ar):

• The weighted mean mass of an atom compared to  the mass of a carbon-12 atom.

• Ions:

• Atoms gain or lose electrons to form ions. Metals form positive ions (cations), non-metals form negative ions (anions).

2. Amount of Substance

• The Mole:

• A mole contains  particles (Avogadro’s constant).

• Calculations:

• , where  is the number of moles.

• Empirical and Molecular Formulas:

• Empirical: Simplest whole-number ratio of atoms.

• Molecular: Actual number of atoms in a molecule (may require molar mass to determine).

• Ideal Gas Equation:

• , with  = pressure (Pa),  = volume (m),  = moles,  = gas constant (8.31 J K mol),  = temperature (K).

• Concentration:

• , where  is concentration in mol dm and  is volume in dm.

3. Acids and Redox Reactions

• Acids and Bases:

• Acids donate H, while bases accept H. Alkalis release OH ions in water.

• Neutralization: Acid + base → salt + water.

• Titrations determine unknown concentrations.

• Redox Reactions:

• Oxidation is the loss of electrons; reduction is the gain of electrons.

• Oxidation numbers track electron transfer.

4. Electron Structure

• Energy Levels:

• Shells are divided into subshells (s, p, d, f) with specific orbitals.

• Electron configuration follows the Aufbau principle (e.g., ).

5. Bonding and Structure

• Ionic Bonding:

• Transfer of electrons between metals and non-metals. Forms giant ionic lattices with high melting points.

• Covalent Bonding:

• Sharing of electrons between non-metals, forming molecules.

• Shapes of molecules are predicted by electron pair repulsion (e.g., linear, tetrahedral).

• Metallic Bonding:

• Positive metal ions surrounded by a sea of delocalized electrons, responsible for conductivity and malleability.

• Intermolecular Forces:

• Weak forces like Van der Waals, dipole-dipole interactions, and hydrogen bonding.

Module 3: Periodic Table and Energy

1. Periodicity

• Trends Across a Period:

• Atomic radius decreases (increased nuclear charge).

• Ionization energy increases (harder to remove electrons).

• Trends Down a Group:

• Atomic radius increases (more shells).

• Reactivity increases in metals but decreases in non-metals.

2. Group Chemistry

• Group 2 (Alkaline Earth Metals):

• Reactivity increases down the group (e.g., Mg < Ba).

• React with water to form hydroxides.

• Group 7 (Halogens):

• Reactivity decreases down the group (e.g., F > Cl > Br > I).

• Displacement reactions occur when a more reactive halogen replaces a less reactive one.

3. Enthalpy Changes

• Types of Enthalpy Change:

• Formation, combustion, and neutralization.

• Hess’s Law: The enthalpy change of a reaction is the same regardless of the route taken.

• Bond Enthalpies:

• Breaking bonds requires energy (endothermic).

• Making bonds releases energy (exothermic).

4. Reaction Rates and Equilibria

• Collision Theory:

• Particles must collide with sufficient energy (activation energy) to react.

• Equilibrium:

• Dynamic equilibrium occurs when forward and reverse reactions happen at the same rate.

• Le Chatelier’s Principle predicts how equilibrium shifts with changes in conditions.

Module 5: Physical Chemistry and Transition Elements

1. Rates of Reaction

• Rate Equation: .

• and  are the orders of reaction, found experimentally.

• Arrhenius Equation: , showing how temperature affects .

2. Equilibria and Kp

• Kp (Gaseous Equilibria):

• Relates equilibrium to partial pressures of gases.

3. Acids, Bases, and Buffers

• pH: .

• Buffers: Resist pH changes by neutralizing small amounts of acid or base.

4. Enthalpy, Entropy, and Free Energy

• Lattice Enthalpy: Energy to form a lattice from gaseous ions.

• Gibbs Free Energy: . Negative  indicates feasibility.

5. Redox and Electrode Potentials

• Electrode Potential: Indicates a species’ ability to gain or lose electrons.

• Electrochemical Cells: Use redox reactions to generate electricity.

6. Transition Elements

• Characteristics:

• Variable oxidation states, colored compounds, catalytic properties.

• Complex Ions: Central metal ion bonded to ligands via coordinate bonds.