Periodic table Study Guide

📖 Core Concepts Periodic Table – Ordered grid of elements by increasing atomic number \(Z\); rows = periods, columns = groups. Atomic Number \(Z\) – Number of protons; uniquely identifies an element. Periods – Each new period starts when electrons begin filling a new principal quantum‑number shell \(n\). Groups – Elements in the same column share the same number of electrons in a particular subshell → similar chemistry. Blocks – Determined by the subshell being filled: s‑block (Groups 1‑2, He) – electrons enter \(ns\) orbitals. p‑block (Groups 13‑18) – electrons enter \(np\) orbitals. d‑block (Groups 3‑12) – electrons enter \((n-1)d\) orbitals. f‑block (lanthanides & actinides) – electrons enter \((n-2)f\) orbitals. Quantum Numbers – \((n,\ell,m{\ell},m{s})\) uniquely describe an orbital. Madelung (Aufbau) Rule – Fill orbitals in order of increasing \(n+\ell\); ties broken by lower \(n\). Periodic Trends – Systematic variations across periods & down groups for: Atomic radius (↓ across, ↑ down) First ionisation energy (↑ across, ↓ down) Electron affinity (more negative across, less negative down) Electronegativity (↑ across, ↓ down) Metallic character (↑ down, ↑ left‑to‑right) Valence vs. Core Electrons – Valence electrons = outermost‑shell electrons; drive bonding. Oxidation State – Formal charge after assigning ionic charges; main‑group: equals group number (or 8‑group for p‑block). 📌 Must Remember \(Z\) = protons = element identity. Group numbers: 1 – 18 left→right. Madelung sequence (first 20): \(1s\), \(2s\), \(2p\), \(3s\), \(3p\), \(4s\), \(3d\), \(4p\), \(5s\), \(4d\), \(5p\), \(6s\), \(4f\), \(5d\), \(6p\), \(7s\), \(5f\), \(6d\), \(7p\)… Key exceptions: Cr: \([Ar]\,3d^{5}4s^{1}\) Cu: \([Ar]\,3d^{10}4s^{1}\) Electronegativity extremes – F = 4.0 (highest), Cs = 0.79 (lowest). Ionisation energy trend – Lowest at each period’s first element (alkali). Electron affinity – Most negative for halogens (except O & F). Hydrogen – Can act as \(+1\) (alkali) or \(-1\) (halogen). Helium – Electron config \(1s^{2}\) but chemically a noble gas (group 18). 4s vs. 3d removal – First ionisation removes a 4s electron before any 3d electrons. 🔄 Key Processes Assigning Ground‑State Electron Configuration List orbitals in Madelung order. Fill each with up to 2 (s), 6 (p), 10 (d), 14 (f) electrons. Apply known exceptions (Cr, Cu, etc.). Predicting Periodic Trend Direction Across a period: ↑ nuclear charge, same \(n\) → stronger attraction → ↓ radius, ↑ IE, ↑ EN, more negative EA. Down a group: ↑ \(n\) → larger radius, ↓ IE/EN, less negative EA, ↑ metallic character. Determining Oxidation States Main‑group: use group number (or 8 − group for p‑block). Transition metals: start with +2 (remove \(ns\) electrons), then consider removal of \( (n-1)d \) electrons for higher states. Choosing Block Placement Identify the subshell being filled by the last electron in the neutral atom → block. Assessing Metallic vs. Nonmetallic Character Left‑most & lower periods → metallic; right‑most & upper periods → nonmetallic. 🔍 Key Comparisons s‑block vs. p‑block – s‑block metals (low IE, low EN) vs. p‑block mix of metals, metalloids, nonmetals (higher IE, higher EN). Hydrogen vs. Alkali Metals – H: 1 valence electron, can gain or lose; alkali metals: readily lose \(e^{-}\) only. Helium vs. Group 2 – He has \(1s^{2}\) (s‑block) but inert like noble gases; Group 2 elements are reactive metals. Cr vs. Expected \(3d^{4}4s^{2}\) – Half‑filled \(3d^{5}\) gives extra stability. Cu vs. Expected \(3d^{9}4s^{2}\) – Fully filled \(3d^{10}\) is more stable. Group 3 (Sc, Y, La) vs. (Sc, Y, Lu) – Lu’s 4f\(^{14}\) makes its chemistry align better with Sc/Y (metallurgist’s view). ⚠️ Common Misunderstandings Atomic number vs. atomic mass – Modern table ordered by \(Z\), not mass. All elements obey Madelung – Cr, Cu, and a few others are exceptions. Helium belongs in group 2 – Chemically a noble gas; placed in group 18. Ionisation energy always increases down a group – It generally decreases; only slight irregularities (e.g., O vs. N). Electron affinity = electronegativity – EA is energy change on adding an electron; EN is a relative tendency in bonds. All transition metals have +2 oxidation state – Many have multiple states (+1 to +7). 🧠 Mental Models / Intuition “Staircase of \(n+\ell\)” – Visualize the Madelung order as a diagonal staircase; each step down adds a higher‑energy subshell. “Valence‑electron‑count = group” – For s‑ and p‑block, the group number directly tells you how many valence electrons. “Half‑filled/fully‑filled stability” – d‑subshells like \(d^{5}\) or \(d^{10}\) give extra stability → explains Cr, Cu anomalies. “Metal‑nonmetal gradient” – Imagine a smooth gradient from metallic left‑bottom to nonmetallic right‑top; metalloids sit near the middle. “Diagonal relationship” – Elements one step down and one step right often behave similarly (e.g., Li ↔ Mg). 🚩 Exceptions & Edge Cases Cr, Cu (and a few others) – Deviate from Madelung due to d‑subshell stability. Hydrogen – Dual placement (group 1 or 17) because it can lose or gain an electron. Helium – Electron configuration suggests group 2, but inertness forces placement in group 18. Ionisation energy of O – Higher than N because removing an electron reduces electron‑electron repulsion in the half‑filled \(2p^{4}\) subshell. Electron affinity of O & F – Less negative than expected due to repulsion in a small, compact orbital. Relativistic effects – Significant for heavy elements (Au, Hg, superheavy), shrinking s‑orbitals and expanding d/f, altering colour, state, and reactivity. 📍 When to Use Which Predicting electron configuration → Use Madelung rule unless the element is Cr, Cu, or known heavy‑element relativistic cases. Estimating oxidation state range → Main‑group: use group number; Transition metals: start with +2 (remove \(ns\)) then consider \( (n-1)d \) removals. Assessing metallic character → Look at position: left‑bottom = metal, right‑top = nonmetal; check metalloids on the “staircase” boundary. Choosing periodic trend direction → Across → left‑to‑right; down → top‑to‑bottom. Deciding block → Identify the highest‑energy subshell occupied in the neutral atom. 👀 Patterns to Recognize Left‑to‑right decrease in atomic radius & increase in EN/IE/EA. Down‑group increase in radius & metallic character, decrease in IE/EN/EA. Sudden jump in IE after a noble gas (e.g., from Ne to Na). Diagonal relationships (Li ↔ Mg, Be ↔ Al). Consistent oxidation‑state limits within a group (e.g., +1 for alkali metals, –1 for halogens). Relativistic contraction evident in heavy metals (gold’s colour, mercury’s liquidity). 🗂️ Exam Traps Confusing group number with period number – Group = column, Period = row. Assigning helium to group 2 – Remember its chemical behavior places it in group 18. Assuming Madelung always works – Watch for Cr, Cu, and other known exceptions. Choosing the highest IE element as “most metallic” – Metallic character is opposite of IE trend. Neglecting that 4s electrons are removed before 3d – Important for ionisation and oxidation‑state predictions. Mix‑up of electron affinity sign – More negative = stronger tendency to gain an electron; noble gases have essentially zero EA. Over‑generalising electronegativity to noble gases – They are not assigned Pauling values because they don’t form stable anions. Treating all transition metals as having the same IE – IE is relatively constant across a d‑block but shows small variations due to subshell stability. --- Use this guide for rapid recall before your exam – each bullet is a high‑yield fact or shortcut you can trust.