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 $\Delta x \cdot \Delta p \ge \frac{h}{4\pi}$