Chemical element Study Guide

📖 Core Concepts Chemical element – a species of atom defined uniquely by its atomic number (Z) = number of protons. Isotope – same Z, different mass number (A = Z + N) (different neutrons). Nuclide – a specific combination of Z, N, and nuclear energy state. Periodic table – arranges elements by increasing Z into periods (rows) and groups (columns) that share physical/chemical traits. Electron shells – discrete energy levels; max electrons per shell = $2n^{2}$ (n = principal quantum number). Blocks (s, p, d, f) – defined by the type of atomic orbital that receives the valence electrons. Strong nuclear force – short‑range force that overcomes proton‑proton repulsion, holding the nucleus together. Radioisotope – an unstable nuclide that decays (α, β⁻, β⁺, γ, fission). Half‑life (t½) – time for half of a radioactive sample to decay. --- 📌 Must Remember 118 elements are officially recognized (94 natural, 24 synthetic). Atomic number = protons = defines element identity. Mass number (A) = protons + neutrons, always an integer. Notation: $^{A}{Z}\mathrm{X}$ (e.g., $^{235}{92}\mathrm{U}$). Standard atomic weight = isotopic‑abundance‑weighted average of an element’s isotopes. Periodic trends: Atomic radius ↓ across a period, ↑ down a group. Ionization energy ↑ across a period, ↓ down a group. Electronegativity ↑ across a period, ↓ down a group (F most electronegative). Alpha decay: emits $^{4}\mathrm{He}$ (α particle). Beta‑minus decay: $n → p + e^{-} + \barν{e}$; Beta‑plus decay: $p → n + e^{+} + ν{e}$. Gamma decay: emission of a photon, no change in Z or A. Units: 1 Bq = 1 decay · s⁻¹; 1 Ci = $3.7×10^{10}$ Bq. 1 Gy = 1 J·kg⁻¹; 1 Sv = Gy × radiation‑weighting factor. --- 🔄 Key Processes Aufbau (electron‑configuration) sequence Fill lowest‑energy subshells first (1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p …). Radioactive decay chain Identify parent nuclide → apply decay type (α, β⁻/β⁺, γ) → calculate new Z and A: α: $Z{\text{new}} = Z - 2$, $A{\text{new}} = A - 4$ β⁻: $Z{\text{new}} = Z + 1$, $A$ unchanged β⁺: $Z{\text{new}} = Z - 1$, $A$ unchanged Half‑life calculation (first‑order decay): $$ N(t) = N{0}\,e^{-\lambda t}, \qquad t{½} = \frac{\ln 2}{\lambda} $$ $\lambda$ = decay constant (s⁻¹). --- 🔍 Key Comparisons Metal vs. Nonmetal Metals: good electrical conductors, form cations. Nonmetals: poor conductors, form anions or covalent bonds. Alpha vs. Beta decay α: heavy, +2 charge, reduces A by 4, Z by 2. β⁻: electron emission, increases Z by 1, A unchanged. s‑block vs. p‑block elements s‑block: valence electrons in s orbital (Group 1‑2, He). p‑block: valence electrons in p orbitals (Groups 13‑18). Natural vs. Synthetic elements Natural: Z ≤ 94, found in Earth’s crust or atmosphere. Synthetic: Z > 94, produced in nuclear reactions. --- ⚠️ Common Misunderstandings Atomic number vs. atomic mass: Z is protons only; atomic mass ≈ A (protons + neutrons) and is not an integer for the element’s average weight. All isotopes are radioactive: Many isotopes are stable; only those with unfavorable N/Z ratios decay. “Heavier” element → “larger” atom: Across a period atoms get smaller despite increasing mass because nuclear charge pulls electrons inward. Electron configuration always follows the order of increasing n: d‑subshells (n‑1) fill before the higher‑n p‑subshell (e.g., 4s before 3d). --- 🧠 Mental Models / Intuition N/Z ratio “balance beam”: Visualize protons pulling electrons inward; neutrons act as “counterweights” that stabilize the nucleus. More neutrons needed as Z grows → explains why heavy nuclei need higher N/Z. Periodic trends as a “slope”: Across a period, think of a downhill slope for radius and uphill for ionization energy; down a group, the slope reverses. Decay as “step‑wise ladder”: Each decay type moves you a fixed step in the (Z, A) grid (α = –2, –4; β⁻ = +1, 0; β⁺ = –1, 0). --- 🚩 Exceptions & Edge Cases Hydrogen: behaves both like an alkali metal (forms H⁺) and a halogen (forms H⁻). Helium: placed in s‑block (1s²) but chemically behaves like a noble gas. Transition metals: often have multiple oxidation states; not strictly predictable by group number. Lanthanides & Actinides: display contraction (lanthanide contraction) affecting ionic radii of subsequent elements. --- 📍 When to Use Which Predict oxidation state: use group number for main‑group elements (Group 1 → +1, Group 2 → +2, Group 17 → –1). Choose electron‑configuration rule: apply Aufbau → Hund’s rule → Pauli exclusion in that order. Identify decay type from daughter nucleus: if A decreases by 4 and Z by 2 → α decay; if A unchanged but Z ↑1 → β⁻; if Z ↓1 → β⁺. Select symbol for generic element: “X” for variable halogen‑like group; “R” for organic radical; “Ln” for any lanthanide, “An” for any actinide. --- 👀 Patterns to Recognize Isotope notation pattern: superscript left of symbol (e.g., $^{14}$C). Group trends: same group = similar valence electron configuration → similar chemical behavior (e.g., alkali metals all +1). Block identification: look at the highest‑energy electron’s subshell (s → left‑most block, p → right‑most, d → middle, f → bottom). Radioactive series: many natural radionuclides belong to the U‑238, U‑235, or Th‑232 decay chains – recognize the repeating pattern of α and β decays. --- 🗂️ Exam Traps Confusing atomic number with mass number: a question may give “element 56” (Ba) but list a mass number of 137 – remember Z = 56, A = 137. Assuming all heavy elements are liquids/gases: only bromine and mercury are liquids at STP; all others are solids (except a few gases). Misreading element symbols: Fe ≠ “F”, Hg ≠ “H”; always check the Latin/Neo‑Latin origin. “First‑ionization energy” vs. “second‑ionization energy”: the first is always lower; a trap may ask for the higher value. Half‑life vs. mean lifetime: t½ = ln 2 / λ; some items give λ and expect you to compute t½. ---