Periodic Table & Properties – Vocabulary Flashcards

Vocabulary flashcards summarizing key terms, laws, trends and concepts from the lecture on the periodic table and periodic properties.

Periodic Table

A systematic classification of elements arranged so their properties recur periodically, giving the greatest informational control with the least effort.

Dobereiner’s Triads

1829 grouping of three elements where the middle member’s atomic weight and properties are roughly the average of the other two.

Newlands’ Law of Octaves

1864 proposal that when elements (excluding H) are arranged by increasing atomic mass, every eighth element shows similar properties.

Mendeleev’s Periodic Law

Physical and chemical properties of elements are periodic functions of their atomic masses (1869).

Modern Periodic Law

Physical and chemical properties of elements are periodic functions of their atomic numbers (Moseley, 1913–14).

Period

Horizontal row of the modern periodic table; there are seven, each beginning with a new principal quantum number (n).

Group

Vertical column of the modern periodic table; 18 numbered groups in IUPAC notation.

Short Period

First three periods (n = 1–3) containing 2, 8 and 8 elements respectively.

Long Period

Periods 4–7 containing 18, 18, 32 and (currently) 19 elements.

Magic Numbers

Recurrent electron capacities 2, 8, 18 and 32 that underlie the length of successive periods.

s-Block Elements

Groups 1 & 2; metals whose valence electrons enter ns orbitals (ns¹–²).

p-Block Elements

Groups 13–18; valence electrons enter np orbitals (ns² np¹–⁶).

d-Block Elements

Groups 3–12; transition metals where differentiating electrons enter (n–1)d orbitals.

f-Block Elements

Lanthanides (4f) and actinides (5f); electrons added to (n–2)f orbitals.

Lanthanides

14 elements Ce (58) to Lu (71) with filling of 4f orbitals.

Actinides

14 elements Th (90) to Lr (103) with filling of 5f orbitals; many are synthetic.

Transition Elements

d-block elements with at least one stable oxidation state having partially filled d-orbitals.

Inner-Transition Elements

f-block elements with incomplete f, d and p shells, displaying multiple oxidation states.

Typical Elements

Third-period main-group elements (Na to Cl) whose properties best represent their vertical families.

Bridge Elements

Second-period elements (Li to F) that resemble diagonally adjacent third-period elements.

Diagonal Relationship

Similarity in properties between an element and the one diagonally below/right (e.g., Li–Mg, Be–Al).

Noble Gases

Group 18 elements with stable ns² np⁶ configurations (He 1s²) and zero valency.

Effective Nuclear Charge (Z_eff)

Net positive charge experienced by valence electrons, calculated as Z – σ (shielding constant).

Shielding Effect

Reduction of the attractive force between nucleus and valence electrons due to inner-shell electrons.

Penetration Effect

Relative ability of an orbital (s > p > d > f) to get close to the nucleus, influencing electron binding.

Atomic Radius

Half the distance between identical nuclei—covalent for molecules, metallic for solids, van der Waals for inert gases.

Ionic Radius

Size of an ion; cations are smaller and anions larger than their parent atoms.

Isoelectronic Ions

Different ions possessing the same number of electrons; size decreases with increasing nuclear charge.

Ionization Energy (IE)

Energy required to remove an electron from a gaseous atom or ion; designated IE₁, IE₂, etc.

First Ionization Energy

Energy needed to detach the first electron from a neutral gaseous atom.

Factors Affecting IE

Atomic size, nuclear charge, shielding, penetration and stable configurations (half/full shells).

Electron Affinity (EA)

Energy released when a gaseous atom gains an electron to form an anion; higher positive value means stronger tendency to gain electrons.

Electronegativity (EN)

Relative tendency of an atom in a molecule to attract shared electron pairs toward itself.

Pauling Electronegativity Scale

Empirical EN values based on extra bond energy; H = 2.1, F = 4.0 is maximum.

Mulliken Electronegativity

Mean of an element’s ionization energy and electron affinity (in eV).

Allred-Rochow Electronegativity

EN calculated from effective nuclear charge divided by covalent radius.

Metallic Character

Tendency of an element to lose electrons forming cations; increases down a group, decreases across a period.

Reducing Power

Ability of an element or ion to donate electrons; parallels metallic character and low IE.

Paramagnetism

Weak attraction of substances containing one or more unpaired electrons to an external magnetic field.

Diamagnetism

Weak repulsion of substances with all electrons paired when placed in a magnetic field.

Ferromagnetism

Strong magnetism retained after removal of external field, seen in Fe, Co, Ni.

Hydration Energy

Enthalpy released when one mole of gaseous ions become solvated by water; higher for small, highly charged ions.

Hydride

Binary compound of hydrogen with another element; may be ionic (Groups 1–2) or covalent (p-block).

Basic Oxide

Metal oxide that yields a base when reacted with water or acids (e.g., Na₂O, MgO).

Acidic Oxide

Non-metal oxide forming acids with water (e.g., SO₃ → H₂SO₄).

Amphoteric Oxide

Oxide behaving as both acid and base (e.g., Al₂O₃, ZnO).

Magic Numbers (Nuclear)

2, 8, 18, 32—electron counts at which shells/periods close, causing periodic property repetition.

Screening Constant (σ)

Numerical measure of shielding; used in Slater’s rules to estimate Z_eff.

Slater’s Rules

Empirical guidelines for computing σ by grouping electrons and assigning shielding contributions.

Covalent Radius

Half the bond length between two identical atoms joined by a single covalent bond.

Metallic Radius

Half the distance between nuclei of adjacent metal atoms in a crystal lattice.

van der Waals Radius

Half the distance between nuclei of non-bonded atoms in neighboring molecules.

Half-Filled Stability

Extra stability associated with exactly half-filled subshells (p³, d⁵, f⁷).

Completely-Filled Stability

Added stability when subshells achieve full occupancy (p⁶, d¹⁰, f¹⁴).

Typical Oxidation State (Group 1)

+1, arising from loss of the single ns¹ electron in alkali metals.

Typical Oxidation State (Group 2)

+2, due to loss of both ns² electrons in alkaline-earth metals.

Magic Number 18

Total electrons filling s, p and d subshells through the third period, producing noble-gas stability for Ar.

Hydrated Ionic Radius

Effective size of an ion in solution; inversely related to hydration energy (Li⁺ ≫ Cs⁺).

Reducing Agent

Species that donates electrons in a redox reaction; strongest among aqueous alkali metals is Li due to high hydration energy.

Paramagnetic Test

Experimental method assessing presence of unpaired electrons by magnetic susceptibility.

Activity Series (Metals)

Empirical ordering of metals by tendency to be oxidized; highest at left/top of periodic table.

Electron Configuration

Distribution of electrons among atomic orbitals, dictating periodic properties.

Penultimate Shell

Second-outermost electron shell; d-block electrons are added here during transition series filling.

Shielding Order

Effectiveness ranking: s > p > d > f for shielding outer electrons from nuclear charge.

Hydration Trend (Alkali Ions)

Li⁺ > Na⁺ > K⁺ > Rb⁺ > Cs⁺ in hydration energy; reverse order in ionic mobility.

Paramagnetic Moment

Quantitative measure (μ = √[n(n+2)] BM) proportional to number of unpaired electrons.

Ferromagnetic Elements

Iron, cobalt, nickel and a few alloys that show permanent magnetism.

Atomic Volume

Volume occupied by one mole of atoms in solid state; varies cyclically within a period, increases down a group.

Density Trend

Within a period, density rises to a mid-period maximum then falls; generally increases down a group due to mass and packing.

Transuranium Elements

Synthetic actinides with atomic numbers > 92 (e.g., Np, Pu, Am).