Chemistry Study Guide

📖 Core Concepts Matter – Anything that has mass and occupies space; can be a pure substance or a mixture. Atom – Smallest unit of an element; nucleus (protons + neutrons) + electron cloud. Neutral atom: \#e⁻ = \#p⁺. Element & Atomic Number (Z) – Pure substance of one atom type; \(Z\) = number of protons, uniquely identifies the element. Isotope – Same \(Z\) but different mass number (protons + neutrons). Mole – SI unit for amount of substance; \(1\ \text{mol}=6.02214076\times10^{23}\) entities (Avogadro constant). Chemical Bonding – Covalent (share e⁻), Ionic (transfer e⁻), Hydrogen, Van der Waals. Octet/Duet Rule – Atoms tend to attain 8 (or 2 for H/Li) valence electrons for stability. Thermodynamics – ΔG predicts spontaneity; ΔG < 0 → spontaneous, ΔG = 0 → equilibrium. Kinetics – Rate constant \(k = A\,e^{-Ea/(RT)}\); \(Ea\) = activation energy. Acid–Base Theories – Arrhenius (produce \(\mathrm{H3O^+}\)/\(\mathrm{OH^-}\)), Brønsted–Lowry (proton donor/acceptor), Lewis (electron‑pair acceptor/donor). Redox – Oxidation = loss of electrons / increase in oxidation number; Reduction = gain of electrons / decrease in oxidation number. Equilibrium & Le Chatelier – System shifts to oppose changes in concentration, temperature, or pressure. --- 📌 Must Remember Avogadro constant: \(6.02214076\times10^{23}\ \text{entities·mol}^{-1}\). Ideal gas law: \(PV = nRT\). Arrhenius equation: \(k = A\,e^{-Ea/(RT)}\). Gibbs free energy: \(\Delta G = \Delta H - T\Delta S\). ΔG sign rule: \(\Delta G<0\) → spontaneous; \(\Delta G>0\) → non‑spontaneous; \(\Delta G=0\) → equilibrium. Exothermic vs. Endothermic: Heat released (–q) vs. absorbed (+q). Exergonic vs. Endergonic: Free‑energy released (ΔG < 0) vs. required (ΔG > 0). pH definition: \(pH = -\log[\mathrm{H3O^+}]\). Lower pH = stronger acid. Acid dissociation constant: Larger \(Ka\) → stronger acid. Oxidizing agent – gets reduced; Reducing agent – gets oxidized. Balancing equations – atoms must be conserved on both sides. --- 🔄 Key Processes Balancing a Chemical Equation List each element, count atoms on reactants & products. Adjust coefficients (whole numbers) to equalize counts. Verify charge balance for ionic equations. Using the Arrhenius Equation Identify \(Ea\) (kJ mol⁻¹) and temperature \(T\) (K). Plug into \(k = A\,e^{-Ea/(RT)}\) to compare rates or estimate effect of temperature change. Calculating ΔG Obtain ΔH and ΔS (from data tables). Compute \(\Delta G = \Delta H - T\Delta S\) at the temperature of interest. Applying Le Chatelier’s Principle Add/remove a reactant/product → shift toward side that counteracts the change. Increase temperature for endothermic → shift right; for exothermic → shift left. Change pressure (gases only) → shift toward side with fewer moles. Identifying Acid–Base Strength (Brønsted–Lowry) Compare \(Ka\) values: larger = stronger acid, conjugate base correspondingly weaker. --- 🔍 Key Comparisons Covalent vs. Ionic Bond Covalent: electron sharing, usually between non‑metals. Ionic: electron transfer, metal + non‑metal → cation + anion. Exothermic vs. Exergonic Exothermic: heat released (ΔH < 0). Exergonic: free energy released (ΔG < 0). Can be endothermic yet exergonic if \(T\Delta S\) dominates. Arrhenius vs. Brønsted–Lowry Acid Arrhenius: defined by production of \(\mathrm{H3O^+}\) or \(\mathrm{OH^-}\) in water. Brønsted–Lowry: proton donor/acceptor; works in any solvent. Oxidizing vs. Reducing Agent Oxidizing: gains electrons → reduced. Reducing: loses electrons → oxidized. Gas Laws (Boyle, Charles, Gay‑Lussac) – all are special cases of the Ideal Gas Law \(PV=nRT\). --- ⚠️ Common Misunderstandings “Exothermic always spontaneous.” → Not always; spontaneity depends on ΔG, not just ΔH. “All acids are strong.” → Strength is quantitative; many acids are weak (small \(Ka\)). “Ions are always in solution.” → Ions also exist in solid crystals (e.g., NaCl lattice). “Electron transfer = oxidation number change.” – Oxidation numbers can change without actual electron flow (formalism). “Mole = mass.” – Mole is a count; convert to mass using molar mass (g mol⁻¹). --- 🧠 Mental Models / Intuition “Electron balance = charge balance.” When writing ionic equations, treat electrons like a currency that must balance on both sides. “Energy landscape” – Visualize reactants climbing an activation hill (Eₐ) to reach products; catalysts lower the hill height. “Periodic trends as slopes.” Imagine atomic radius decreasing left‑to‑right (down‑ward slope) and ionization energy increasing (upward slope). “Le Chatelier as a seesaw.” Adding stress pushes the seesaw toward the side that relieves that stress. --- 🚩 Exceptions & Edge Cases Hydrogen & Lithium – follow the duet rule (2 valence electrons) instead of octet. Transition metals – may have variable oxidation states; octet rule often does not apply. Water auto‑ionization – pure water has \([\mathrm{H3O^+}] = [\mathrm{OH^-}] = 1.0\times10^{-7}\ \text{M}\) (pH = 7) despite no added acid/base. Gas law limits – Ideal gas law fails at high pressure/low temperature; real gases deviate (need Van der Waals corrections). --- 📍 When to Use Which Calculate gas volume or pressure? → Use Ideal Gas Law \(PV=nRT\). Predict reaction direction? → Compute ΔG; if unknown, compare ΔH and ΔS with temperature. Identify acid/base in non‑aqueous media? → Apply Lewis theory (electron‑pair transfer). Determine rate change with temperature? → Plug values into the Arrhenius equation. Balance redox in aqueous solution? → Use half‑reaction method (oxidation + reduction) and ensure electrons cancel. --- 👀 Patterns to Recognize “+ Heat → Endothermic shift” – Temperature increase moves equilibrium toward the endothermic direction. “Large \(Ka\) → Small pKa” – Strong acids have low pKa values (pKa = –log Ka). “Metal + Non‑metal → Ionic compound” – Look for large electronegativity difference (> 1.7). “ΔG = 0 → No net change” – System at equilibrium, forward and reverse rates equal. “Periodic trends – Down a group: radius ↑, ionization energy ↓; across a period: radius ↓, ionization energy ↑. --- 🗂️ Exam Traps Confusing ΔH with ΔG – A question may give an exothermic ΔH but ask about spontaneity; you must consider ΔS and temperature. Mixing up “oxidizing agent” with “oxidant” – Both remove electrons, but the wording can invert the role if you forget which species is reduced. Balancing equations without charge – In ionic equations, forgetting to balance total charge leads to an incorrect answer. Assuming “strong acid = strong base” – Strength is independent; a strong acid can be paired with a weak base and vice‑versa. Using Boyle’s law when temperature changes – Boyle’s law only holds at constant \(T\); if \(T\) changes, use the full ideal gas law. ---