Chapter 1 Vocabulary Flashcards
Anti-bonding Molecular orbitals
A high-energy molecular orbital resulting from the destructive interference between atomic orbitals
Atomic orbital
A three-dimensional plot of ψ2 of a wavefunction. It is a region of space that can accommodate electron density.
Aufbau principle
A rule that determines the order in which orbitals are filled by electrons. Specifically, the lowest energy orbital is filled first.
Bent geometry
A geometric arrangement for three atoms connected to each other (such as H−O−H) in a nonlinear fashion.
Bonds
A strong force of attraction holding atoms together in a molecule or crystal, resulting from the sharing or transfer of electrons
Bonding Molecular orbitals
A low-energy molecular orbital resulting from the constructive interference between atomic orbitals.
Constitutional Isomers
Compounds that have the same molecular formula but differ in the way the atoms are connected.
Constructive interference
When two waves interact with each other in a way that produces a wave with a larger amplitude.
Covalent bond
A bond that results when two atoms share a pair of electrons.
Debye
A unit of measure for dipole moments, where 1 debye = 10 − 18 esu ⋅ cm.
Degenerate
Having the same energy.
Degenerate orbital
Orbitals that have the same energy.
Divalent
An element that forms two bonds, such as oxygen.
Dipole moment (µ)
The amount of partial charge (δ) on either end of a dipole multiplied by the distance of separation (d): μ = δ × d
Dipole-dipole interactions
The resulting net attraction between two dipoles.
Electron density
A term associated with the probability of finding an electron in a particular region of space.
Electrostatic potential maps
A three-dimensional, rainbow like image used to visualize partial charges (δ) in a compound.
Electronegativity
the degree to which an element tends to gain electrons and form negative ions in chemical reactions
Formal Charge
A charge associated with any atom that does not exhibit the appropriate number of valence electrons.
H. O. M. O.
The Highest Occupied Molecular Orbital.
Hund’s Rule
When considering electrons in atomic orbitals, a rule that states that one electron is placed in each degenerate orbital first, before electrons are paired up.
Hybridization
A mathematical procedure in which standard atomic orbitals are combined to form new, hybrid orbitals
Hydrogen bonding
A special type of dipole-dipole interaction that occurs between an electronegative atom and a hydrogen atom that is connected to an electronegative atom.
Hydrophilic
A polar group that has favorable interactions with water.
Hydrophobic
A nonpolar group that does not have favorable interactions with water.
Induction
The withdrawal of electron density that occurs when a bond is shared by two atoms of differing electronegativity.
Ionic bond
A bond that results from the force of attraction between two oppositely charged ions.
Lewis Structure
A drawing style in which the electrons take center stage.
London Dispersion forces
Attractive forces between transient dipole moments, observed in alkanes.
Lone Pair
A pair of unshared, or nonbonding, electrons.
L. U. M. O.
The lowest unoccupied molecular orbital.
Micelle
A group of molecules arranged in a sphere such that the surface of the sphere is comprised of polar groups, rendering the micelle water soluble.
Molecular dipole moment
The vector sum of the individual dipole moments in a compound.
Molecular orbitals
Orbitals associated with an entire molecule rather than an individual atom.
Molecular orbital theory
A description of bonding in terms of molecular orbitals, which are orbitals associated with an entire molecule rather than an individual atom.
Monovalent
An element that can form one bond (such as hydrogen).
Nodes
In atomic and molecular orbitals, a location where the value of ψ is zero.
Octet Rule
The observation that second-row elements (C, N, O, and F) will generally form the necessary number of bonds to achieve a full valence shell (eight electrons).
Pauli Exclusion Principle
The rule that states that an atomic orbital or molecular orbital can accommodate a maximum of two electrons with opposite spin.
Partial charge
(δ-) negative, (δ+) positive.
Charges created due to the asymmetric distribution of electrons in chemical bonds.
Pi (π) bond
A bond formed from adjacent, overlapping p orbitals.
Polar Covalent Bond
A bond in which the difference in electronegative values of the two atoms is between 0.5 and 1.7.
Quantum mechanics
A mathematical description of an electron that incorporates its wavelike properties.
Sigma (σ) bond
A bond that is characterized by circular symmetry with respect to the bond axis.
sp-hybridized:
Atomic orbitals that are achieved by mathematically averaging one s orbital with only one p orbital to form 2 hybridized atomic orbitals.
sp2 Hybridized orbital
Atomic orbitals that are achieved by mathematically averaging one s orbital with two p orbitals to form 3 hybridized atomic orbitals.
sp3 Hybridization:
Atomic orbitals that are achieved by mathematically averaging one s orbital with three p orbitals to form 4 hybridized atomic orbitals.
Steric number
The total of electron pairs (single bonds + lone pairs) for an atom in a compound.
Tetrahedral
The geometry of an atom with four bonds separated from each other by 109.5°.
Tetravalent
An element, such as carbon, that forms four bonds.
Triagonal pyramidal
A geometry adopted by an atom that has one lone pair and a steric number of 4.
Trigonal planar
A geometry adopted by an atom with a steric number of 3. All three groups lie in one plane and are separated by 120°.
Trivalent
An element, such as nitrogen, that forms three bonds.
Valence bond theory
A theory that treats a bond as the sharing of electrons that are associated with individual atoms, rather than being associated with the entire molecule.
VESPR Theory
Valence shell electron pair repulsion theory, which can be used to predict the geometry around an atom.
Wave equations (ψ)
Describes the total energy of an electron when in the vicinity of a proton.
Wave functions
Solutions to wave equations (ψ) where ψ2 represents the probability of finding an electron in a particular location.
Zeff
effective nuclear charge
