Back/Chemistry: Atoms First 2e
Chapter 4 Review50 Total Terms

Chapter 4 Key Terms & Vocabulary

Comprehensive index of all scientific terms, definitions, and examples across Chapter 4: Chemical Bonding and Molecular Geometry.

Filter:
Showing 50 of 50 terms
From Section 4.1: Ionic Bonding
cationIons
A positively charged ion formed when a neutral atom loses one or more valence electrons.
Example: Ca2+\text{Ca}^{2+}
From Section 4.1: Ionic Bonding
anionIons
A negatively charged ion formed when a neutral atom gains one or more valence electrons.
Example: Cl−\text{Cl}^-
From Section 4.1: Ionic Bonding
ionic compoundsCompounds
Chemical compounds composed of cations and anions held together by electrostatic forces.
Example: NaCl\text{NaCl} (table salt)
From Section 4.1: Ionic Bonding
ionic bondsBonding
Electrostatic forces of attraction holding oppositely charged cations and anions together in an ionic substance.
From Section 4.1: Ionic Bonding
binary ionic compoundCompounds
An ionic compound consisting of exactly two elements: a metallic cation and a nonmetallic anion.
Example: Al2O3\text{Al}_2\text{O}_3
From Section 4.1: Ionic Bonding
inert pair effectPeriodic Trends
The tendency of heavy post-transition elements in groups 13–15 to retain their valence s-electron pair, resulting in lower-than-expected positive charges.
Example: Tl+,Sn2+,Pb2+\text{Tl}^+, \text{Sn}^{2+}, \text{Pb}^{2+}
From Section 4.1: Ionic Bonding
isoelectronicElectronic Structure
Having the exact same electron configuration as another chemical species, typically a noble gas.
Example: Ca2+\text{Ca}^{2+} is isoelectronic with Ar\text{Ar}
From Section 4.2: Covalent Bonding
covalent bondsBonding
Chemical bonds formed by the mutual attraction of atoms for a shared pair of valence electrons.
Example: The H–H bond in H₂
From Section 4.2: Covalent Bonding
bond lengthMolecular Structure
The internuclear distance between two bonded atoms where the lowest potential energy of the system is achieved.
Example: Optimal internuclear distance in H₂ molecule
From Section 4.2: Covalent Bonding
pure covalent bondBonding
A covalent bond in which bonding electrons are shared equally between identical atoms with zero electronegativity difference.
Example: Cl–Cl bond in Cl₂
From Section 4.2: Covalent Bonding
polar covalent bondBonding
A covalent bond where electrons are shared unequally, producing an uneven electron density distribution and partial charges.
Example: H–Cl with Hδ⁺ and Clδ⁻
From Section 4.2: Covalent Bonding
electronegativityPeriodic Properties
A dimensionless measure of an atom's tendency to attract shared electrons or electron density toward itself in a bond.
Example: Fluorine is the most electronegative element (EN = 4.0)
From Section 4.2: Covalent Bonding
PaulingHistorical Figures
Linus Pauling, the pioneering chemist who established the first relative electronegativity scale and foundational bonding theories.
Example: Pauling electronegativity values
From Section 4.3: Chemical Nomenclature
NomenclatureGeneral Principles
A systematic collection of rules and conventions used for naming chemical compounds.
Example: IUPAC guidelines for inorganic chemistry.
From Section 4.3: Chemical Nomenclature
binary compoundsCompound Classes
Chemical compounds composed of exactly two different elements.
Example: NaCl\text{NaCl}, CO2\text{CO}_2, and H2O\text{H}_2\text{O}.
From Section 4.3: Chemical Nomenclature
acidsAcids & Bases
Substances that release hydrogen ions (H+\text{H}^+) when dissolved in aqueous solutions.
Example: HCl\text{HCl} and H2SO4\text{H}_2\text{SO}_4.
From Section 4.3: Chemical Nomenclature
hydratesIonic Compounds
Ionic compounds that contain water molecules bound as an integral part of their crystal lattices in specific ratios.
Example: CuSO4⋅5H2O\text{CuSO}_4 \cdot 5\text{H}_2\text{O} (copper(II) sulfate pentahydrate).
From Section 4.3: Chemical Nomenclature
binary acidAcids & Bases
An acid composed of hydrogen combined with a single nonmetallic element, dissolved in water.
Example: HCl(aq)\text{HCl}(aq) (hydrochloric acid).
From Section 4.3: Chemical Nomenclature
oxyacidsAcids & Bases
Acids containing hydrogen, oxygen, and at least one other element, typically derived from an oxoanion.
Example: HNO3\text{HNO}_3 (nitric acid) and H3PO4\text{H}_3\text{PO}_4 (phosphoric acid).
From Section 4.4: Lewis Symbols and Structures
Lewis symbolAtomic Structure
An elemental symbol surrounded by dots that represents the valence electron configuration of an atom or monatomic ion.
Example: A sodium atom represented with one dot.
From Section 4.4: Lewis Symbols and Structures
Lewis structuresMolecular Bonding
Diagrams that show the bonding between atoms in a molecule or polyatomic ion, using lines for covalent bonds and dots for lone pairs.
Example: H−O¨−HH-\ddot{O}-H for water.
From Section 4.4: Lewis Symbols and Structures
lone pairsMolecular Bonding
Pairs of valence electrons that are not involved in covalent bond formation.
Example: Three pairs of unshared electrons on each chlorine atom in Cl2Cl_2.
From Section 4.4: Lewis Symbols and Structures
single bondCovalent Bonding
A covalent bond formed by the sharing of a single pair of electrons between two atoms.
Example: The H−ClH-Cl bond in hydrogen chloride.
From Section 4.4: Lewis Symbols and Structures
octet ruleBonding Rules
The tendency of main-group atoms to form enough covalent bonds to achieve eight valence electrons in their outer shell.
Example: Carbon forming four bonds in CH4CH_4.
From Section 4.4: Lewis Symbols and Structures
double bondMultiple Bonding
A covalent bond formed by the sharing of two pairs of electrons between two atoms.
Example: The C=OC=O bond in formaldehyde (CH2OCH_2O).
From Section 4.4: Lewis Symbols and Structures
triple bondMultiple Bonding
A covalent bond formed by the sharing of three pairs of electrons between two atoms.
Example: The C≡OC \equiv O bond in carbon monoxide (COCO).
From Section 4.4: Lewis Symbols and Structures
SmalleyHistory & Discovery
Richard Smalley, a pioneer in fullerene chemistry who co-discovered buckminsterfullerene (C60C_{60}) and was honored as a father of nanotechnology.
Example: Co-recipient of the 1996 Nobel Prize in Chemistry.
From Section 4.4: Lewis Symbols and Structures
free radicalsExceptions to Octet Rule
Molecules that contain an odd number of valence electrons and possess an unpaired electron.
Example: Nitric oxide (NONO).
From Section 4.4: Lewis Symbols and Structures
hypervalent moleculesExceptions to Octet Rule
Molecules with a central atom from period 3 or higher that holds more than eight valence electrons using empty d-orbitals.
Example: PCl5PCl_5 and SF6SF_6.
From Section 4.5: Formal Charges and Resonance
Formal ChargeChemical Bonding
The hypothetical charge an atom would have if bonding electrons were distributed evenly; calculated as valence electrons minus lone pair electrons minus half of bonding electrons.
Example: Used to evaluate the relative stability of competing Lewis structures.
From Section 4.5: Formal Charges and Resonance
Molecular StructureChemical Bonding
The specific three-dimensional arrangement or geometry of atoms within a molecule or polyatomic ion.
Example: Determined using formal charge guidelines when multiple Lewis structures are possible.
From Section 4.5: Formal Charges and Resonance
ResonanceChemical Bonding
A condition where two or more valid Lewis structures with identical atom placements can be written for a single species.
Example: Observed in molecules like ozone (O₃) and the carbonate ion (CO₃²⁻).
From Section 4.5: Formal Charges and Resonance
Resonance FormsChemical Bonding
Individual valid Lewis structures that contribute to the true electronic description of a resonance-stabilized molecule.
Example: The two distinct drawings used to represent the nitrite ion (NO₂⁻).
From Section 4.5: Formal Charges and Resonance
Resonance HybridChemical Bonding
The actual electronic structure of a molecule or ion, representing the weighted average of all valid resonance forms.
Example: The true state of NO₂⁻ where both N–O bonds are identical in length and strength.
From Section 4.6: Molecular Structure and Polarity
bond angleGeometry
The angle between any two bonds that share a common central atom.
Example: The H-O-H bond angle in water is 104.5°.
From Section 4.6: Molecular Structure and Polarity
bond distanceGeometry
The straight-line distance between the nuclei of two bonded atoms.
Example: Measured in picometers (pm) or Ångstroms (Å).
From Section 4.6: Molecular Structure and Polarity
Valence shell electron-pair repulsion theory (VSEPR theory)Theory
A model used to predict 3D molecular structures by assuming valence electron pairs maximize distance to minimize electrostatic repulsion.
Example: Predicts a tetrahedral shape for methane (CH4\text{CH}_4).
From Section 4.6: Molecular Structure and Polarity
linearGeometries
A molecular or electron-pair geometry where two regions of electron density lie 180° apart on opposite sides of a central atom.
Example: CO2\text{CO}_2 or BeF2\text{BeF}_2.
From Section 4.6: Molecular Structure and Polarity
trigonal planarGeometries
A geometry where three regions of electron density lie in a single plane 120° apart around a central atom.
Example: BCl3\text{BCl}_3.
From Section 4.6: Molecular Structure and Polarity
tetrahedralGeometries
A geometry where four regions of electron density point toward the corners of a tetrahedron with 109.5° angles.
Example: CH4\text{CH}_4.
From Section 4.6: Molecular Structure and Polarity
trigonal bipyramidalGeometries
A geometry where five regions of electron density form three equatorial positions (120° apart) and two axial positions (90° apart).
Example: PCl5\text{PCl}_5.
From Section 4.6: Molecular Structure and Polarity
octahedralGeometries
A geometry where six regions of electron density point toward the corners of an octahedron with 90° angles.
Example: SF6\text{SF}_6.
From Section 4.6: Molecular Structure and Polarity
electron-pair geometryGeometry
The spatial arrangement of all regions of high electron density, including both chemical bonds and lone pairs, around a central atom.
Example: Tetrahedral for water's four electron regions.
From Section 4.6: Molecular Structure and Polarity
molecular structureGeometry
The 3D spatial arrangement and shape defined strictly by the positions of the atoms in a molecule, excluding lone pairs.
Example: Bent for water, trigonal pyramidal for ammonia.
From Section 4.6: Molecular Structure and Polarity
axial positionPositions
In a trigonal bipyramidal geometry, one of the two positions located directly vertical along the main rotational axis (90° to the equatorial plane).
Example: Top and bottom positions of PCl5\text{PCl}_5.
From Section 4.6: Molecular Structure and Polarity
equatorial positionPositions
In a trigonal bipyramidal geometry, one of the three positions forming a triangular plane around the equator (120° apart).
Example: Occupied by lone pairs in SF4\text{SF}_4 or ClF3\text{ClF}_3.
From Section 4.6: Molecular Structure and Polarity
bond dipole momentPolarity
The separation of electrical charge within an individual polar covalent bond.
Example: H-F bond dipole.
From Section 4.6: Molecular Structure and Polarity
vectorMath/Physics
A mathematical quantity possessing both magnitude and direction, used to represent bond dipoles and molecular dipoles.
Example: Bond dipole arrows.
From Section 4.6: Molecular Structure and Polarity
polar moleculePolarity
A molecule possessing a net nonzero dipole moment due to polar bonds and an asymmetric 3D structure.
Example: H2O\text{H}_2\text{O} or NH3\text{NH}_3.
From Section 4.6: Molecular Structure and Polarity
dipole momentGeneral
A quantitative measure of the net charge separation and overall polarity of an entire molecule.
Previous
4.6 Molecular Structure and Polarity
Next
5.1 Valence Bond Theory