Physical Chemistry

Atomic Structure & Electron Shells

The arrangement of electrons in an atom determines its chemical properties and reactivity. Electron shells are filled according to the Aufbau principle, with each shell holding a specific maximum number of electrons.

• Electron Configuration: Sodium (Na, atomic number 11) has the configuration K(2), L(8), M(1). The M shell (third shell) contains only 1 electron in sodium, not 2. Magnesium, by contrast, would have 2 electrons in the M shell.

• Valence Electrons: The number of electrons in the outermost shell (valence shell) determines an element's chemical reactivity.

• Example: Na: 1s2 2s2 2p6 3s1

Chemical Bonding

Chemical bonds form when atoms share or transfer electrons to achieve stable electron configurations. The ammonium ion (NH4+) is a classic example of covalent and coordinate bonding.

• Covalent Bonds: Atoms share electron pairs. In NH4+, nitrogen forms three standard covalent bonds with hydrogen.

• Coordinate Covalent Bond: One bond in NH4+ is formed when nitrogen donates both electrons to bond with a proton (H+).

• Shared Electron Pairs: NH4+ has four shared electron pairs.

• Example: Lewis structure of NH4+ shows four N–H bonds.

Molecular Geometry

The three-dimensional arrangement of atoms in a molecule is predicted by the Valence Shell Electron Pair Repulsion (VSEPR) theory.

• NH3 (Ammonia): Has three bonding pairs and one lone pair on nitrogen, resulting in a trigonal pyramidal geometry.

• VSEPR Notation: AX3E (A = central atom, X = bonded atoms, E = lone pairs).

• Example: Ammonia is non-planar due to the lone pair repulsion.

Crystal Types

Solids are classified by the nature of their bonding and structure.

• Molecular Crystals: Held together by van der Waals forces (e.g., CO2 as dry ice).

• Covalent Network Crystals: Atoms connected by a continuous network of covalent bonds (e.g., SiO2 in quartz).

• Ionic Crystals: Composed of ions held by electrostatic attraction (e.g., CaO).

Stoichiometry & Concentration

Stoichiometry involves quantitative relationships in chemical reactions. Calculating the maximum amount of reactant or product requires balanced equations and molar relationships.

• Example Calculation: Maximum Mg reacting with HCl: 0.048 g (≈0.002 mol). If HCl volume is fixed, the required concentration is 1.0 mol/L.

• Key Equation:

Ideal Gas Law & Vapor Pressure

The ideal gas law relates pressure, volume, temperature, and amount of gas.

• Equation:

• Application: Calculating the volume at complete vaporization under specified conditions yields 0.50 L.

Salt Hydrolysis & pH

Salts formed from weak acids and strong bases hydrolyze in water, affecting pH.

• Example: Na2CO3 is derived from weak diprotic acid (H2CO3) and strong base (NaOH). Its solution is basic due to carbonate ion hydrolysis.

• Hydrolysis Equation:

Redox Titration

Redox titrations involve electron transfer reactions, often using permanganate as an oxidizing agent.

• Example: 30 mL of 0.200 M KMnO4 reacts with 50 mL of 0.300 M H2O2.

• Balanced Equation:

• Mole Ratio: 2:5 (KMnO4:H2O2).

Electrolysis & Faraday's Law

Electrolysis uses electric current to drive non-spontaneous chemical reactions. Faraday's law relates the amount of substance produced to the charge passed.

• Faraday's Law: (where Q = charge in coulombs, n = moles of electrons, F = Faraday constant ≈ 96500 C/mol).

• Example: 672 mL of gas at STP (0.03 mol) requires 3.86 × 103 C.

Chemical Equilibrium

Chemical equilibrium is the state where the rates of the forward and reverse reactions are equal. Le Chatelier's principle predicts the effect of changes in conditions.

• Endothermic Reaction: Lowering temperature shifts equilibrium toward the exothermic direction (formation of HI in the given example).

• General Equation:

Inorganic Chemistry

Reactivity & Hydrogen Gas

The reactivity series ranks metals by their ability to displace hydrogen from acids.

• Copper (Cu): Located below hydrogen in the series; does not react with dilute HCl to produce H2 gas.

• Example: Zn + 2HCl → ZnCl2 + H2 (reaction occurs); Cu + HCl (no reaction).

Metal Ionization & Reactivity

Some metals are less reactive than hydrogen and do not react with acids or reduce other ions.

• Silver (Ag): Does not react with HCl.

• Platinum (Pt): Very unreactive; does not reduce Fe2+ in FeSO4.

Acid-Base Theory (Water)

Water can act as both an acid and a base (amphoteric), depending on the reaction partner.

• As an Acid: Donates a proton to ammonia or active metals (e.g., Na).

• Example Equations:

Sulfur Compounds & Bleaching

Sulfur dioxide (SO2) is used as a reducing bleaching agent, especially for delicate materials.

• Bleaching Mechanism: SO2 reduces colored compounds, unlike chlorine, which acts as an oxidizer.

• Example: SO2 is used for bleaching wool and silk.

Inorganic Cation Analysis

Qualitative analysis of cations involves precipitation and complex formation.

• Zinc (Zn2+): Forms a white precipitate with limited NaOH (Zn(OH)2), which dissolves in excess NaOH (amphoteric behavior) or excess NH3 (forming [Zn(NH3)4]2+).

Organic Chemistry & Polymers

Organic Isomers & Halogenation

Isomerism refers to compounds with the same molecular formula but different structures. Halogenation can create chiral centers.

• C3H5Br Isomers: Four possible unsaturated isomers exist. Addition of Br2 across the double bond in two of these creates an asymmetric (chiral) carbon.

Alkenes & Chemical Testing

Alkenes react with bromine water in an addition reaction, decolorizing the solution.

• Ethylene (C2H4): Decolorizes bromine water, confirming the presence of a double bond.

• Equation:

Organic Reaction Products

Functional group transformations are central to organic chemistry.

• Nitration of Benzene: Produces nitrobenzene using concentrated HNO3 and H2SO4.

• Sulfonation: Produces benzenesulfonic acid using fuming sulfuric acid.

• Equation:

Cycloalkenes & Addition Stoichiometry

Cycloalkenes can undergo addition reactions with halogens, increasing molecular weight.

• Cyclohexadiene (C6H8): Addition of 2 moles of Br2 yields C6H8Br4 (tetrabromide).

• Molecular Weight Calculation:

Synthetic Polymers

Polymers are large molecules formed from repeating units (monomers). Many synthetic polymers require formaldehyde as a precursor.

• Phenol Resin: Produced by condensation of phenol and formaldehyde.

• Vinylon: Produced by treating polyvinyl alcohol with formaldehyde.

Test Preparation Recommendations

• Master the Reactivity Series: Essential for predicting metal reactions in inorganic and electrochemical contexts.

• Practice Redox Balancing: Key for stoichiometry and electrochemistry problems.

• Visualize Molecular Geometry: Memorize VSEPR shapes for common molecules (NH3, H2O, CH4).

• Memorize Polymer Precursors: Recognize reagents for common plastics (Phenol resin, Vinylon, PET).

Uncertainties & Assumptions

• Calculations assume standard atomic weights and ideal gas behavior.

• Minor variations may occur due to rounding or different conventions (e.g., Mg = 24.0 vs 24.3 g/mol).