Chapter 3 Review98 Total Terms
Chapter 3 Key Terms & Vocabulary
Comprehensive index of all scientific terms, definitions, and examples across Chapter 3: Electronic Structure and Periodic Properties of Elements.
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From Section 3.1: Electromagnetic Energy
NewtonHistorical Figures
17th-century scientist who advanced a corpuscular (particle) view of light using prisms and lenses.
From Section 3.1: Electromagnetic Energy
HuygensHistorical Figures
17th-century scientist who explained optical reflection and refraction using a wave model of light.
From Section 3.1: Electromagnetic Energy
YoungHistorical Figures
Physicist who proved the wave nature of light by observing interference patterns in double-slit experiments.
From Section 3.1: Electromagnetic Energy
MaxwellHistorical Figures
Developed classical electromagnetic theory, proving light consists of oscillating electric and magnetic waves.
From Section 3.1: Electromagnetic Energy
electromagnetic radiationWaves & Light
Energy transmitted via oscillating electric and magnetic fields traveling at the speed of light.
Example: X-rays, visible light, radio waves
From Section 3.1: Electromagnetic Energy
waveWaves & Light
An oscillation or periodic movement that transports energy through space without permanently displacing matter.
From Section 3.1: Electromagnetic Energy
wavelengthWave Properties
The linear distance between two consecutive peaks or troughs of a wave (denoted by λ).
From Section 3.1: Electromagnetic Energy
frequencyWave Properties
The number of wave cycles passing a specific point per second (denoted by ν).
From Section 3.1: Electromagnetic Energy
amplitudeWave Properties
The magnitude of a wave
From Section 3.1: Electromagnetic Energy
hertz (Hz)Units
The SI unit for frequency, defined as cycles per second (s⁻¹).
From Section 3.1: Electromagnetic Energy
electromagnetic spectrumWaves & Light
The complete range of all types of electromagnetic radiation arranged by wavelength and frequency.
From Section 3.1: Electromagnetic Energy
interference patternsWave Properties
Fringe patterns resulting from the constructive and destructive overlap of waves.
From Section 3.1: Electromagnetic Energy
Standing wavesWaves & Light
Waves that remain constrained within a specific region of space, foundational to atomic structure.
From Section 3.1: Electromagnetic Energy
stationary wavesWaves & Light
Alternative term for standing waves that do not travel through space.
From Section 3.1: Electromagnetic Energy
quantizationQuantum Mechanics
The restriction of a property, such as energy, to discrete, specific values rather than a continuous range.
From Section 3.1: Electromagnetic Energy
nodesWave Properties
Points or lines in a standing wave where displacement and motion are zero.
From Section 3.1: Electromagnetic Energy
continuous spectrumSpectroscopy
An unbroken series of all wavelengths of light emitted by heated solids, liquids, or dense gases.
From Section 3.1: Electromagnetic Energy
blackbodyThermodynamics
An idealized emitter that absorbs and re-emits all incident thermal radiation.
From Section 3.1: Electromagnetic Energy
photonsQuantum Mechanics
Discrete packets or particles of electromagnetic radiation whose energy depends directly on frequency.
From Section 3.1: Electromagnetic Energy
endothermicThermodynamics
Chemical or physical processes that absorb energy (such as light absorption) from their surroundings.
From Section 3.1: Electromagnetic Energy
exothermicThermodynamics
Chemical or physical processes that release energy (such as light emission) to their surroundings.
From Section 3.1: Electromagnetic Energy
wave-particle dualityQuantum Mechanics
The fundamental principle that light and matter exhibit both wavelike and particle-like properties.
From Section 3.1: Electromagnetic Energy
line spectraSpectroscopy
Discrete, narrow lines of light emitted by excited low-pressure gases, unique to each chemical element.
From Section 3.1: Electromagnetic Energy
BalmerHistorical Figures
Physicist who derived an empirical formula for the visible spectral lines of hydrogen.
From Section 3.1: Electromagnetic Energy
RydbergHistorical Figures
Developed a generalized mathematical formula to calculate all spectral lines of atomic hydrogen.
From Section 3.1: Electromagnetic Energy
BohrHistorical Figures
Physicist who incorporated Planck
From Section 3.2: The Bohr Model
RutherfordHistorical Figures
Physicist who established the nuclear model of the atom, depicting a dense positive nucleus surrounded by moving electrons.
From Section 3.2: The Bohr Model
BohrHistorical Figures
Physicist who introduced quantization to the atomic model, successfully explaining hydrogen
From Section 3.2: The Bohr Model
Bohr’s modelAtomic Models
An early atomic model where electrons travel in specific circular orbits with quantized energies without radiating energy.
Example: En = -k(Z² / n²)
From Section 3.2: The Bohr Model
ground electronic stateEnergy States
The lowest energy state of an atom where the electron occupies the n = 1 orbit.
Example: Hydrogen atom with its electron at n = 1
From Section 3.2: The Bohr Model
quantum numbersQuantum Mechanics
Integer values that specify the allowed energy levels and physical characteristics of electrons in atoms.
From Section 3.3: Development of Quantum Theory
de BroglieWave-Particle Duality
French physicist who predicted that material particles exhibit wavelike characteristics governed by wavelength \(\lambda = \frac{h}{mv}\).
From Section 3.3: Development of Quantum Theory
DavissonWave-Particle Duality
Scientist who experimentally proved the wavelike behavior of electrons by observing diffraction patterns through nickel crystals.
From Section 3.3: Development of Quantum Theory
GermerWave-Particle Duality
Collaborator with Davisson who co-discovered electron interference patterns using crystal lattices.
From Section 3.3: Development of Quantum Theory
Heisenberg uncertainty principleQuantum Principles
Fundamental physical limit stating that simultaneous exact measurement of a particle
From Section 3.3: Development of Quantum Theory
wavefunctionsQuantum Mechanics
Mathematical functions (\(\psi\)) representing three-dihydroxy stationary waves that describe quantum states of electrons.
From Section 3.3: Development of Quantum Theory
BornQuantum Mechanics
Proposed that the square of a wavefunction
From Section 3.3: Development of Quantum Theory
quantum mechanicsQuantum Mechanics
The theoretical framework describing the energy and behavior of microscopic systems using wave equations and quantized states.
From Section 3.3: Development of Quantum Theory
principal quantum numberQuantum Numbers
Quantum number (\(n\)) defining the primary energy level and general size of an atomic shell.
From Section 3.3: Development of Quantum Theory
shellsQuantum Numbers
Concentric energy levels radiating outward from the atomic nucleus, designated by \(n = 1, 2, 3, \dots\).
From Section 3.3: Development of Quantum Theory
atomic orbitalQuantum Mechanics
A specific three-dihydroxy region in space where an electron has a high probability of being found.
From Section 3.3: Development of Quantum Theory
secondary (angular momentum) quantum numberQuantum Numbers
Quantum number (\(l\)) that dictates the 3D shape of an atomic subshell (\(l = 0, 1, \dots, n-1\)).
From Section 3.3: Development of Quantum Theory
subshellQuantum Numbers
A collection of orbitals within a given shell that share the same angular momentum quantum number (\(l\)).
From Section 3.3: Development of Quantum Theory
s orbitalsOrbital Types
Spherical atomic orbitals corresponding to an angular momentum quantum number of \(l = 0\).
From Section 3.3: Development of Quantum Theory
p orbitalsOrbital Types
Dumbbell-shaped atomic orbitals corresponding to an angular momentum quantum number of \(l = 1\).
From Section 3.3: Development of Quantum Theory
d orbitalsOrbital Types
Complex multi-lobed atomic orbitals corresponding to an angular momentum quantum number of \(l = 2\).
From Section 3.3: Development of Quantum Theory
f orbitalsOrbital Types
Highly complex atomic orbitals corresponding to an angular momentum quantum number of \(l = 3\).
From Section 3.3: Development of Quantum Theory
magnetic quantum numberQuantum Numbers
Quantum number (\(m_l\)) specifying the spatial orientation of an orbital within a subshell (\(-l\) to \(+l\)).
From Section 3.3: Development of Quantum Theory
degenerate orbitalsQuantum Mechanics
Orbitals that possess exactly the same energy level (e.g., orbitals within the same subshell).
From Section 3.3: Development of Quantum Theory
spin quantum numberQuantum Numbers
Quantum number describing the intrinsic quantum spinning or angular momentum state of an electron.
From Section 3.3: Development of Quantum Theory
msQuantum Numbers
Symbol for the spin quantum number, which can take values of \(+\frac{1}{2}\) or \(-\frac{1}{2}\).
From Section 3.3: Development of Quantum Theory
Pauli exclusion principleQuantum Principles
Principle stating that no two electrons in the same atom can share an identical set of all four quantum numbers.
From Section 3.4: Electronic Structure of Atoms (Electron Configurations)
electron configurationQuantum Structure
The specific distribution of electrons among the atomic orbitals of an atom.
Example: 1s² 2s² 2p⁴ for oxygen
From Section 3.4: Electronic Structure of Atoms (Electron Configurations)
Aufbau principleQuantum Mechanics
A method of building ground-state electron configurations by sequentially adding electrons to the lowest available energy subshells.
Example: Filling 4s before 3d
From Section 3.4: Electronic Structure of Atoms (Electron Configurations)
Orbital diagramsQuantum Structure
Pictorial representations of electron configurations using boxes for orbitals and arrows for electron spins.
Example: Up and down arrows in a 1s box
From Section 3.4: Electronic Structure of Atoms (Electron Configurations)
Hund’s ruleQuantum Rules
States that the lowest-energy arrangement of electrons in degenerate orbitals maximizes the number of unpaired electrons with parallel spins.
Example: Placing single electrons in three separate 2p boxes before pairing
From Section 3.4: Electronic Structure of Atoms (Electron Configurations)
valence electronsAtomic Properties
Electrons occupying the outermost principal shell of an atom that determine chemical reactivity.
Example: The 3s¹ electron in sodium ([Ne]3s¹)
From Section 3.4: Electronic Structure of Atoms (Electron Configurations)
core electronsAtomic Properties
Inner-shell electrons that correspond to a noble gas configuration.
Example: The [Ne] core in sodium
From Section 3.4: Electronic Structure of Atoms (Electron Configurations)
valence shellAtomic Properties
The outermost electron shell of an atom containing the highest principal quantum number (n).
Example: The n = 3 shell in phosphorus
From Section 3.4: Electronic Structure of Atoms (Electron Configurations)
Main group elementsPeriodic Table
Elements in which the last added electron enters an s or p orbital in the outermost shell.
Example: Groups 1, 2, and 13–18
From Section 3.4: Electronic Structure of Atoms (Electron Configurations)
representative elementsPeriodic Table
Alternative terminology for main group elements, encompassing all s-block and p-block elements.
Example: Carbon, fluorine, potassium
From Section 3.4: Electronic Structure of Atoms (Electron Configurations)
Transition elements or transition metalsPeriodic Table
Metallic elements in which the last electron added enters a d orbital, or possessing partially filled d orbitals.
Example: Iron, copper, titanium
From Section 3.4: Electronic Structure of Atoms (Electron Configurations)
Inner transition elementsPeriodic Table
Metallic elements in which the last electron added occupies an f orbital, consisting of the lanthanides and actinides.
Example: Uranium, neodymium
From Section 3.5: Periodic Variations in Element Properties
covalent radiusAtomic Structure
One-half the distance between the nuclei of two identical atoms joined by a covalent bond.
Example: Used to gauge relative atomic sizes.
From Section 3.5: Periodic Variations in Element Properties
effective nuclear charge, ZeffPeriodic Trends
The net positive charge experienced by an outer electron after accounting for core electron shielding.
Example: Zeff = Z - shielding
From Section 3.5: Periodic Variations in Element Properties
isoelectronicAtomic Structure
Describing atoms or ions that possess the exact same electron configuration.
Example: N³⁻, O²⁻, F⁻, Ne, Na⁺, Mg²⁺, and Al³⁺ all share 1s²2s²2p⁶.
From Section 3.5: Periodic Variations in Element Properties
ionization energyPeriodic Trends
The energy required to remove the most loosely bound electron from a gaseous atom or ion in its ground state.
Example: X(g) → X⁺(g) + e⁻
From Section 3.5: Periodic Variations in Element Properties
electron affinityPeriodic Trends
The energy change associated with the addition of an electron to a gaseous atom to form an anion.
Example: X(g) + e⁻ → X⁻(g)
From Section 3.6: The Periodic Table
MendeleevHistory
Russian chemist widely credited with creating the first periodic table arranged by atomic mass and predicting missing elements.
From Section 3.6: The Periodic Table
MeyerHistory
German chemist who independently published a periodic table ordered by atomic mass around 1870.
From Section 3.6: The Periodic Table
periodic lawPrinciples
The principle stating that the properties of the elements are periodic functions of their atomic numbers.
From Section 3.6: The Periodic Table
periodic tableStructure
A tabular arrangement of elements ordered by atomic number, highlighting recurring chemical patterns.
From Section 3.6: The Periodic Table
periodsStructure
Horizontal rows of elements across the periodic table (also referred to as series).
From Section 3.6: The Periodic Table
seriesStructure
Alternative term for horizontal rows or periods in the periodic table.
From Section 3.6: The Periodic Table
groupsStructure
Vertical columns of the periodic table containing elements with similar chemical behaviors.
From Section 3.6: The Periodic Table
metalsElement Classes
Elements that are shiny, malleable, ductile, and good conductors of heat and electricity.
From Section 3.6: The Periodic Table
nonmetalsElement Classes
Elements that appear dull and are poor conductors of heat and electricity.
From Section 3.6: The Periodic Table
metalloidsElement Classes
Elements with intermediate conductivity and a mixture of metallic and nonmetallic properties.
From Section 3.6: The Periodic Table
main-group elementsSubdivisions
Elements in columns 1, 2, and 13–18 of the periodic table (also called representative elements).
From Section 3.6: The Periodic Table
representative elementsSubdivisions
Alternative term for main-group elements in columns 1, 2, and 13–18.
From Section 3.6: The Periodic Table
transition metalsSubdivisions
Metallic elements located in columns 3–12 of the periodic table.
From Section 3.6: The Periodic Table
inner transition metalsSubdivisions
Elements in the two bottom rows of the periodic table (lanthanides and actinides).
From Section 3.6: The Periodic Table
lanthanidesSubdivisions
The top row of inner transition metals, spanning atomic numbers 57–71.
From Section 3.6: The Periodic Table
actinidesSubdivisions
The bottom row of inner transition metals, spanning atomic numbers 89–103.
From Section 3.6: The Periodic Table
alkali metalsGroup Families
Reactive metals in Group 1 (excluding hydrogen) that form 1:1 ratio compounds with oxygen.
From Section 3.6: The Periodic Table
alkaline earth metalsGroup Families
Reactive metals in Group 2 that form compounds with a 1:2 ratio to hydrogen.
From Section 3.6: The Periodic Table
pnictogensGroup Families
Elements located in Group 15 of the periodic table.
From Section 3.6: The Periodic Table
chalcogensGroup Families
Elements located in Group 16 of the periodic table, also known as the oxygen family.
From Section 3.6: The Periodic Table
halogensGroup Families
Reactive nonmetals located in Group 17 of the periodic table.
From Section 3.6: The Periodic Table
noble gasesGroup Families
Unreactive, chemically stable elements in Group 18 of the periodic table.
From Section 3.6: The Periodic Table
inert gasesGroup Families
Alternative term for Group 18 noble gases, referencing their historical lack of reactivity.
From Section 3.7: Ionic and Molecular Compounds
monatomic ionsIons
Ions formed from only a single atom that has gained or lost electrons.
Example: Na⁺, Cl⁻, Ca²⁺
From Section 3.7: Ionic and Molecular Compounds
polyatomic ionsIons
Electrically charged molecules consisting of a group of bonded atoms that act as a single, discrete unit.
Example: NH₄⁺, SO₄²⁻, OH⁻
From Section 3.7: Ionic and Molecular Compounds
OxyanionsIons
Polyatomic ions that contain one or more oxygen atoms combined with another element.
Example: NO₃⁻ (nitrate), PO₄³⁻ (phosphate)
From Section 3.7: Ionic and Molecular Compounds
ionic bondsChemical Bonding
Electrostatic forces of attraction between oppositely charged cations and anions resulting from electron transfer.
Example: The attraction between Na⁺ and Cl⁻ in NaCl
From Section 3.7: Ionic and Molecular Compounds
covalent bondsChemical Bonding
Attractive forces between the positively charged nuclei of bonded atoms and one or more pairs of shared electrons.
Example: The bonds holding the H and O atoms together in H₂O
From Section 3.7: Ionic and Molecular Compounds
ionic compoundCompounds
A compound containing ions held together by electrostatic ionic bonds, typically formed by combining a metal and a nonmetal.
Example: NaCl, CaSO₄, Al₂O₃
From Section 3.7: Ionic and Molecular Compounds
molecular compoundsCompounds
Compounds that consist solely of discrete, neutral molecules formed when nonmetal atoms share electrons; also called covalent compounds.
Example: H₂O, CO₂, NH₃
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3.7 Ionic and Molecular Compounds
End of current section