Unit 3 – The Periodic Table, Electron Configuration, and Atomic & Light Properties

Vocabulary flashcards covering key terms from the lecture on the periodic table, electron configurations, periodic trends, Coulombic interactions, and the wave-particle nature of light.

Periodic Law

Chemical and physical properties of elements repeat periodically when elements are arranged by increasing atomic number.

Group (Family)

Vertical column on the periodic table; elements share valence-electron configuration and similar properties.

Period

Horizontal row on the periodic table; indicates the number of occupied principal energy levels in the atoms.

Atomic Number

Number of protons in an atom’s nucleus; determines element identity and its order on the periodic table.

Valence Electron

Electron located in the outermost (valence) shell; governs an atom’s bonding behavior and chemical reactivity.

Valence Shell

Highest occupied principal energy level of an atom containing valence electrons.

Noble Gas

Group 18 element with a full valence shell (8 electrons, except He with 2); chemically inert.

Halogen

Group 17 non-metal with 7 valence electrons; extremely reactive and forms salts.

Alkali Metal

Group 1 metal with one valence electron; forms +1 cations, very reactive, never found free in nature.

Alkaline Earth Metal

Group 2 metal with two valence electrons; forms +2 cations, reactive, found only in compounds.

Transition Metal

Element in the d-block (Groups 3–12); often forms colored ions and uses two outer shells for bonding.

Main-Group Element

Element in the s- or p-block (Groups 1–2 and 13–18); shows predictable valence-electron patterns.

Post-Transition Metal

Metal located between d-block and metalloids; follows regular valence-electron rules and is relatively soft.

Metalloid

Element bordering the staircase (B, Si, Ge, As, Sb, Te, Po); exhibits both metallic and non-metallic properties.

Metal

Element that is lustrous, malleable, ductile, conductive, and typically solid at room temperature.

Nonmetal

Element usually gaseous or brittle solid, poor conductor, right of staircase; forms anions or covalent bonds.

Atomic Mass

Weighted average mass of an element’s isotopes measured in atomic mass units (amu).

Malleable

Ability of a metal to be hammered or rolled into thin sheets without breaking.

Ductile

Ability of a substance to be drawn into wires.

Conductor

Material that readily transmits heat or electricity.

Cation

Positively charged ion formed when an atom loses one or more electrons.

Anion

Negatively charged ion formed when an atom gains one or more electrons.

Electron Configuration

Shorthand that shows the distribution of electrons among orbitals (e.g., 1s² 2s² 2p⁶).

Sublevel

Set of orbitals of equal energy within a principal level (s, p, d, f).

Orbital Notation

Diagrammatic representation of electron configuration using boxes and arrows to show electron spins.

Noble Gas Configuration

Abbreviated electron configuration that starts with the symbol of the previous noble gas in brackets.

Aufbau Principle

Electrons occupy lowest-energy orbitals available before filling higher ones.

Pauli Exclusion Principle

No two electrons in an atom can have the same set of four quantum numbers; an orbital holds max two electrons with opposite spins.

Hund’s Rule

Electrons fill degenerate orbitals singly with parallel spins before pairing up.

Shielding Electron

Inner-shell electron that reduces the nuclear attraction felt by valence electrons.

Shielding Effect

Reduction of effective nuclear charge on valence electrons due to inner electrons; increases down a group.

Effective Nuclear Charge (ENC)

Net positive charge experienced by valence electrons after accounting for shielding; increases across a period.

Coulomb’s Law

Electrostatic energy ∝ (charge₁·charge₂)/distance²; greater charge or smaller distance means stronger attraction/repulsion.

Atomic Radius

Distance from nucleus to outer edge of electron cloud; increases down a group, decreases across a period.

Ionic Radius

Distance from nucleus to outer edge of electron cloud in an ion; decreases when metals lose electrons, increases when nonmetals gain electrons.

Ionization Energy

Energy required to remove an electron from a gaseous atom/ion; increases across a period, decreases down a group.

Electronegativity

Relative ability of an atom in a bond to attract shared electrons; highest for F, increases across a period, decreases down a group.

Electron Affinity

Energy change when an atom gains an electron; typically becomes more negative across a period.

Metallic Character

Extent to which an element exhibits metallic properties; increases down a group, decreases across a period.

Ground State

Lowest-energy arrangement of electrons in an atom.

Excited State

Higher-energy state achieved when an electron absorbs energy and moves to a higher orbital.

Photon

Massless quantum of electromagnetic radiation carrying energy E = hv.

Electromagnetic Radiation (EMR)

Form of energy exhibiting wave-particle duality; includes radio waves, microwaves, infrared, visible light, UV, X-rays, γ-rays.

Electromagnetic Spectrum

Continuous range of electromagnetic radiation arranged by wavelength or frequency.

Wavelength (λ)

Shortest distance between equivalent points on a wave; measured in meters, cm, or nm.

Frequency (ν)

Number of wave cycles passing a point per second; measured in hertz (Hz).

Amplitude

Height of a wave from origin to crest or trough; relates to intensity.

Speed of Light (c)

Constant velocity of EM waves in vacuum, 3.00 × 10⁸ m/s; relates wavelength and frequency (c = λν).

Planck’s Constant (h)

Proportionality constant in E = hv, 6.626 × 10⁻³⁴ J·s.

Quantum (plural Quanta)

Smallest discrete amount of energy that can be absorbed or emitted.

Energy–Frequency Relationship

Photon energy directly proportional to frequency: E = hv.

Wave–Frequency Relationship

Wavelength inversely proportional to frequency: c = λν.

Photoelectric Effect

Emission of electrons from a metal surface when light of sufficient frequency shines on it; supports particle nature of light.

Double-Slit Experiment

Experiment showing interference patterns that demonstrate wave behavior of light and electrons.

Interference

Combination of overlapping waves resulting in regions of reinforcement or cancellation.

Constructive Interference

Superposition of waves in phase, producing increased amplitude (bright bands).

Destructive Interference

Superposition of waves out of phase, producing decreased amplitude (dark bands).

Wave–Particle Duality

Concept that particles such as electrons and photons exhibit both wave and particle properties.

Emission Spectrum

Set of discrete wavelengths emitted by excited atoms returning to lower energy levels; acts as an element’s fingerprint.

Absorption Spectrum

Wavelengths absorbed by atoms as electrons transition to higher levels; complements emission spectrum.

Spectroscope

Instrument used to view and measure an element’s line spectra.

Line Spectra (Atomic Spectra)

Series of discrete emission or absorption lines corresponding to specific electron transitions.

Electron Shell (Principal Energy Level)

Region around nucleus defined by principal quantum number n; holds sublevels and orbitals.

Subshell Filling Order

Order in which orbitals fill (1s, 2s, 2p, 3s, 3p, 4s, 3d, etc.), following Aufbau principle.

4s vs 3d Rule

For transition metals, 4s electrons are filled before 3d but are removed first when forming cations.

S-Block

Left-side block of periodic table where valence electrons occupy s orbitals.

P-Block

Right-side block where valence electrons occupy p orbitals.

D-Block

Middle block containing transition metals with valence electrons in d orbitals.

F-Block

Inner transition elements (lanthanides & actinides) with valence electrons in f orbitals.

Alkali Metal Oxide Formula

General formula X₂O produced when Group 1 metals react with oxygen.

Alkaline Earth Metal Oxide Formula

General formula XO produced when Group 2 metals react with oxygen.

First Ionization Energy

Energy needed to remove the highest-energy (outermost) electron from a neutral atom.

Ionization Energy Exception

Filled or half-filled subshells (e.g., noble gases, Group 2) create local maxima in ionization energy trend.

Effective Core Charge

Alternative term for effective nuclear charge; calculated as (protons – inner electrons).

Lattice Energy

Energy released when gaseous ions form an ionic solid; magnitude follows Coulomb’s law.

Coulombic Attraction

Force of attraction between opposite charges; magnitude increases with charge and decreases with distance.

Coulombic Repulsion

Force pushing like charges apart; same dependence on charge and distance as attraction.

Microwave

EM radiation with frequencies ~10⁹ Hz and wavelengths on the centimeter scale.

Planck–Einstein Equation

Another name for E = hv, linking photon energy with frequency.

Spectral Line

Single wavelength emitted or absorbed corresponding to a specific electronic transition.

Photoelectron

Electron ejected from a metal surface during the photoelectric effect.

Duality of the Electron

Concept that electrons, like light, exhibit both particle-like and wave-like behavior.

Inverse Square Law (Electrostatic)

Coulombic force or energy diminishes with the square of the distance between charges.

Shielding Trend Down Group

Increase in inner shells leads to greater shielding and lower effective nuclear charge on valence electrons.

Atomic Radius Trend Across Period

Radius decreases left-to-right owing to increasing nuclear charge with constant shielding.

Atomic Radius Trend Down Group

Radius increases due to addition of new energy levels farther from nucleus.

Ionization Energy Trend Across Period

Ionization energy increases left-to-right because of stronger effective nuclear charge.

Ionization Energy Trend Down Group

Ionization energy decreases down a group due to increased radius and shielding.

Electronegativity Trend

Electronegativity rises across a period and falls down a group; fluorine is the most electronegative element.

Ionic Radius Trend (Metals)

Metal cations are smaller than their parent atoms due to electron loss and reduced electron-electron repulsion.

Ionic Radius Trend (Nonmetals)

Non-metal anions are larger than their parent atoms owing to electron gain and increased repulsion.

Electron Configuration (Ground State)

Lowest-energy distribution of electrons among orbitals as dictated by the Aufbau, Pauli, and Hund rules.