CHAPTER 2 Atomic Structure and Periodicity - Vocabulary Flashcards

A set of vocabulary flashcards covering the key terms and concepts from Chapter 2 notes on atomic structure and periodicity, including electromagnetic radiation, quantum mechanics, orbitals, and periodic trends.

Electromagnetic Radiation

Radiation that can be described by its wavelength, frequency, and speed (c ≈ 2.9979 × 10^8 m/s); can be viewed as waves and as streams of photons.

Wavelength

Distance between successive crests (or troughs) of a wave; typically measured in meters or Angstroms; denoted by λ.

Frequency

Number of wave cycles per second; denoted by ν (or f); units s^-1 (hertz).

Speed of Light

Constant c, the speed at which all electromagnetic radiation travels in vacuum; approximately 2.9979 × 10^8 m/s.

Planck's Constant

Fundamental constant (h) relating energy and frequency of a photon via E = hν; value ≈ 6.626 × 10^-34 J·s.

Photons

Quanta or particles of light; each has energy E = hν and momentum p = h/λ.

Energy of a Photon (E = hν)

Energy carried by a photon; proportional to its frequency and inversely related to wavelength.

Photon Energy and Wavelength Relationship

Photon energy is inversely proportional to wavelength: as λ decreases, E = hc/λ increases.

Photoelectric Effect

Emission of electrons from a material (usually a metal) when light shines on it; requires photons with sufficient energy to overcome the material's work function.

Einstein (in context of light)

Explained the photoelectric effect by treating light as consisting of photons with quantized energy.

Dual Nature of Light

Light exhibits both wave-like and particle-like properties (waves and photons).

Continuous Spectrum

Spectrum showing all wavelengths without gaps; produced by incandescent solids, liquids, or gases at high temperature.

Line Spectrum

Spectrum showing discrete wavelengths; typical for atomic or ionic emissions corresponding to transitions between energy levels.

Hydrogen Spectrum

Emission spectrum of hydrogen showing discrete lines, indicating discrete energy levels.

Bohr Model of the Hydrogen Atom

Model proposing quantized angular momentum and circular orbits; explained hydrogen’s emission spectrum but is ultimately incomplete.

Quantization

Restriction of certain physical quantities (like energy) to discrete values rather than a continuum.

Wave (Quantum) Mechanical Model

Modern model where electrons are described by standing waves and probability distributions (orbitals) rather than precise orbits.

Standing Wave

Wave that remains in a constant position, used to describe electron wave behavior in atoms.

Wave Function

Mathematical function Ψ whose squared magnitude |Ψ|^2 gives the probability distribution of finding an electron in space.

Orbital

Region of space where there is a high probability of finding an electron; described by quantum numbers (n, l, m_l).

Probability Distribution

Function giving the likelihood of finding an electron at a given location; in atoms, derived from |Ψ|^2.

Radial Probability Distribution

Probability of finding an electron at a distance r from the nucleus, integrating over angles.

Heisenberg Uncertainty Principle

Δx·Δp ≥ h/4π; cannot simultaneously know exact position and momentum of a particle.

Quantum Numbers

Set of numbers (n, l, ml, ms) that describe the size, shape, and orientation of orbitals and electron spin.

Principal Quantum Number (n)

Denotes the main energy level (shell) of an electron; n = 1, 2, 3, …; determines energy and size of orbital.

Angular Momentum Quantum Number (l)

Determines orbital shape; values range from 0 to n-1 for each n (s, p, d, f…).

Magnetic Quantum Number (m_l)

Determines orbital orientation in space; m_l takes integer values between -l and +l.

Subshell

Set of orbitals within a principal shell with same n and l (e.g., 2p, 3d).

Nodal Surfaces (Nodes)

Surfaces where the probability density is zero for an electron in an orbital.

Degenerate Orbital

Orbitals with the same energy in a given atom (e.g., the three 2p orbitals in the absence of external fields).

Electron Spin

Intrinsic angular momentum of electrons, quantified by ms; possible values are +1/2 and -1/2.

Electron Spin Quantum Number (m_s)

Magnetic spin quantum number; can be +1/2 or -1/2 for electrons.

Pauli Exclusion Principle

No two electrons in an atom can have the same set of four quantum numbers (n, l, ml, ms).

Aufbau Principle

Electrons fill the lowest-energy orbitals first before occupying higher ones.

Hund's Rule

Electrons in degenerate orbitals fill singly with parallel spins before pairing occurs.

Valence Electrons

Electrons in the outermost principal energy level; primarily determine chemical properties.

Core Electrons

Electrons in inner shells not in the valence shell; shield and influence energy levels.

Polyelectronic Atoms

Atoms with more than one electron; electron–electron repulsion affects energy levels and spectra.

Transition Metals

Elements in the d-block whose chemistry is largely governed by d-electron configurations.

Lanthanide Series (Lanthanides)

Elements with electron configurations filling 4f orbitals; part of the f-block.

Actinide Series (Actinides)

Elements with electron configurations filling 5f orbitals; part of the f-block.

Main-Group Elements (Representative Elements)

Elements in the s- and p-blocks that define the typical trends in the periodic table.

Atomic Radii

Size of an atom (often half the distance between two bonded nuclei); trends vary with period and group.

First Ionization Energy

Energy required to remove the outermost (first) electron from a neutral atom in the gas phase.

Second Ionization Energy

Energy required to remove a second electron after the first has already been removed.

Metalloids (Semimetals)

Elements with properties between metals and nonmetals; lie along the dividing line on the periodic table.

Ground State

Lowest-energy arrangement of electrons in an atom.

De Broglie Wavelength

Wavelength associated with a particle, λ = h/p; shows wave-like behavior of matter.

Bonding and Periodic Trends (Overview)

Trends in ionization energy, atomic radii, and orbital filling explain periodic table patterns and chemical properties.

Wave-Particle Duality of Matter

All matter exhibits both wave-like and particle-like properties, especially evident for electrons and light.

Diffraction

Bending and spreading of waves around obstacles or through slits; evidence of wave nature.

Diffraction Pattern

Pattern produced by the interference of diffracted waves; used to study wave properties.

Standing Wave Model of the Electron

Electron described as a standing wave in an atom, leading to quantized energy levels.

Hydrogen Emission Spectrum

Discrete spectral lines indicating transitions between discrete energy levels in hydrogen.

Noble Gases

Group of inert, highly stable elements with filled outer electron shells.

Shielding and Penetration (Concepts)

Shielding: decrease of nuclear attraction felt by outer electrons due to inner electrons; penetration: the extent inner electrons shield outer electrons from the nucleus.