Chemistry of the Elements in the Periodic Table - VOCAB Flashcards

Periodic Table: Structure and Layout

  • 7 Periods, 18 Groups (7 periods, 18 groups)

  • Blocks: s, p, d, f

  • Groups: Representative (Groups 1A–8A), Transition (Groups 1B–8B), Inner Transition (f-block)

  • Lanthanide series (La–Lu) and Actinide series (Ac–Lr)

  • New elements (as of the notes): Nh (113), Mc (115), Ts (117), Og (118)

  • Glenn Seaborg era: discovery of transuranic elements (beyond U, atomic number 92)

Key Contributors to the Periodic Table

  • Antoine Lavoisier: In the late 18th century, he compiled a list of elements based on substances that could not be broken down further, defining "element" in a modern chemical sense. He also helped establish the Law of Conservation of Mass.

  • Johann Wolfgang Döbereiner: In the 1820s, he developed the "Law of Triads," observing that groups of three elements with similar chemical properties (like Cl, Br, I) had atomic weights where the middle element was approximately the average of the other two.

  • John Newlands: In 1864, he proposed the "Law of Octaves," arranging elements by increasing atomic weight and noting that properties tended to repeat every eight elements, similar to musical octaves.

  • Dmitri Mendeleev: Published the first widely recognized periodic table in 1869, arranging elements by increasing atomic mass and leaving gaps for undiscovered elements, predicting their properties.

  • Lothar Meyer: Independently developed a similar periodic classification of elements around the same time as Mendeleev, focusing on physical properties like atomic volume.

  • Henry Moseley: In 1913, refined the periodic law by arranging elements based on increasing atomic number, which resolved discrepancies in Mendeleev's table and led to the modern periodic table.

Periodic Laws and Trends

  • Periodic Law: Some properties are periodic functions of atomic number

  • Periodic variation/trends include Ionization energy, Electron affinity, Atomic radius, Electronegativity, Metallic and Nonmetallic character

  • General trends:

  • Ionization energy: increases across a period; decreases down a group

  • Electron affinity: generally increases across a period; varies down a group

  • Atomic radius: decreases across a period; increases down a group

  • Metallic character: decreases across a period; increases down a group

  • Nonmetallic character: increases across a period; decreases down a group

Electronic Configuration and Aufbau Principle

  • Aufbau principle: electrons fill from lower to higher energy subshells

  • Subshell capacities: s:2, p:6, d:10, f:14

  • Orbital capacities: s has 1 orbital, p has 3 orbitals, d has 5, f has 7 orbitals

  • Example: Calcium (Atomic No. 20) electron configuration: 1s^2 \, 2s^2 \, 2p^6 \, 3s^2 \, 3p^6 \, 4s^2 = [Ar] \, 4s^2

Quantum Numbers

  • Principal (n): n=1 \text{ to } 7; size/distance of orbital

  • Azimuthal (ℓ): \text{0 to 3}; subshell shapes

  • \text{ℓ} = 0 o s, 1 o p, 2 o d, 3 o f

  • Magnetic (mℓ): orientation; values depend on ℓ: example for p: -1, 0, 1; for d: -2,-1,0,1,2, etc.

  • Spin (ms): +1/2 or -1/2

  • Electron arrangement follows Pauli exclusion, Hund’s rule, and Heisenberg uncertainty principles

Group/Block Overview (Representative vs Transition vs Inner Transition)

  • Representative Elements (Groups 1A-8A): valence electrons equal group number (roughly) and show regular variation

  • Transition Metals (Groups 1B-8B): variable valence; multiple oxidation states; typical metallic behavior

  • Inner Transition Metals (f-block): Lanthanides and Actinides; similar properties within each series; Lanthanides mostly non-radioactive except Pm; Actinides all radioactive

Representative Element Groups (Key Families)

  • Group 1A – Alkali Metals: valence 1; common valence +1

  • H is a special case (nonmetal, forms H⁺, H₃O⁺, hydride H⁻, H₂)

  • Group 2A – Alkaline Earth Metals: valence 2; common valence +2

  • Group 3A – Boron Family: valence 3; common valence +3

  • Group 4A – Carbon Family: valence 4; common valence +4, -4

  • Group 5A – Nitrogen Family: valence 5; common valence often -3

  • Group 6A – Oxygen Family: valence 6; common valence -2

  • Group 7A – Halogens: valence 7; common valence -1

  • Group 8A – Noble Gases: valence 8; no stable valence

Transition and Subgroups (Selected)

  • Group 1B – Coinage metals: Cu, Ag, Au with multiple oxidation states (+1, +2, …)

  • Group 2B – Volatile metals: Zn, Cd, Hg with common oxidation states (+2; Hg also +1)

  • Group 3B – Scandium subgroup (Sc, Y, La–Lu, Ac–Lr)

  • Group 4B – Titanium subgroup (Ti, Zr, Hf)

  • Group 5B – Vanadium subgroup (V, Nb, Ta)

  • Group 6B – Chromium subgroup (Cr, Mo, W)

  • Group 7B – Manganese subgroup (Mn, Tc, Re, Bh)

  • Group 8B – Iron Triad (Fe, Co, Ni) and heavier metals (Rh, Pd, Os, Ir, etc.)

Important Atomic Mass Reference

  • Notable pairs: H (≈1 g/mol), C (≈12 g/mol), N (≈14 g/mol), O (≈16 g/mol), Na (≈23 g/mol), Mg (≈24 g/mol), Al (≈27 g/mol), S (≈32 g/mol), Cl (≈35.5 g/mol), K (≈39 g/mol), Ca (≈40 g/mol), Fe (≈56 g/mol), Cu (≈63.5 g/mol), Zn (≈65 g/mol), Ag (≈108 g/mol), I (≈127 g/mol)

Types of Chemical Reactions (Key Categories)

  • Synthesis/Combination/Direct Union (Addition): A + B → AB; C + O₂ → CO₂; 2H₂ + O₂ → 2H₂O; Haber process: 3H2 + N2 \rightarrow 2NH_3

  • Decomposition/Analysis (Elimination): AB → A + B; Complete combustion: CH4 + 2O2 \rightarrow CO2 + 2H2O; Electrolysis: 2H2O \rightarrow 2H2 + O_2

  • Single Displacement (Substitution): AB + X → AX + B; activity series governs reactivity (metals: Li > K > Ba > Ca > Na > Mg > Al > Zn > Fe > Cu > Au, etc.; nonmetals: F > Cl > Br > I)

  • Double Displacement (Metathesis/Neutralization/Precipitation): AB + CD → AD + CB; e.g., NaOH + HCl → NaCl + H₂O; AgNO₃ + NaCl → AgCl(s) + NaNO₃; rearrangements possible (cis/trans, etc.)

Quick Reference Concepts

  • Aufbau order for building electron configurations: 1s o 2s o 2p o 3s o 3p o 4s o 3d o 4p o 5s o 4d o 5p o 6s o 4f o 5d o 6p o 7s o 5f o 6d o 7p

  • Electron configuration notation: [\text{Noble gas}] \text{ followed by valence electrons}

  • Quantum numbers provide the allowed states and electron arrangements; Pauli exclusion, Heisenberg uncertainty, Hund’s rule underpin most stable configurations

Quick Recall Tips

  • Periodic table structure matters: 7 periods, 18 groups; s/p/d/f blocks

  • Valence electrons drive chemical properties: Groups 1A-8A show regular trends; noble gases are chemically inert due to filled valence shell

  • Reactions reveal functional group behavior: synthesis, decomposition, single/double displacement, and rearrangements cover most classic reactions

The electron arrangement in atoms is governed by fundamental principles, including the Pauli exclusion principle, Hund's rule, and the Heisenberg uncertainty principle.

  • Pauli Exclusion Principle: This principle states that no two electrons in an atom can have the same set of four quantum numbers (n, ℓ, mℓ, ms). This means that each orbital can hold a maximum of two electrons, and these two electrons must have opposite spins (+1/2 and -1/2).

  • Hund's Rule: For degenerate orbitals (orbitals with the same energy, such as the three p orbitals or five d orbitals), Hund's rule states that electrons will first occupy separate orbitals with parallel spins before pairing up in any one orbital. This maximizes the total spin and leads to a more stable configuration.

  • Heisenberg Uncertainty Principle: This principle states that it is impossible to precisely determine both the position and momentum of an electron (or any other particle) simultaneously. The more accurately one quantity is known, the less accurately the other can be known. Mathematically, this is expressed as \Delta x \cdot \Delta p \ge \frac{h}{4\pi}