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Chemical Bond

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a strong attractive force that exists between atoms in a molecule

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Ionic bond

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a bond between oppositely charged ions. The ions are formed from atoms by transfer of one or more electrons

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Covalent Bond

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a bond formed between two or more atoms by a sharing of electrons

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Metallic Bonds

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bonding, usually in solid metals, in which the bonding electrons are relatively free to more throughout the three-dimensional structure

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Lewis Symbol

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(electron-dot symbol) the chemical symbol for an element, with a dot for each valence electron

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Octet Rule

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a rule stating that bonded atoms tend to possess or share a total of eight valence-shell electrons

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Single Bond

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a covalent bond involving one electron pair

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Double Bond

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a covalent bond involving two electron pairs

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Bond length

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the distance between the centers of two bonded atoms

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Bond polarity

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a measure of the degree to which the electrons are shared unequally between two atoms in a chemical bond

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Nonpolar Covalent Bond

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a covalent bond in which the electrons are shared equally

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Polar Covalent Bond

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a covalent bond in which the electrons are not shared equally

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Electronegativity

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a measure of the ability of an atom that is bonded to another atom to attract electrons to itself

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Dipole

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a molecule with one end having a partial negative charge and the other end having a partial positive charge; a polar molecule

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Dipole Moment

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a measure of the separation and magnitude of the positive and negative charges in polar molecules

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Polar Molecule

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a molecule that possesses a nonzero dipole moment

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Formal Charge

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the number of valence electrons in an isolated atom minus the number of electrons assigned to the atom in the Lewis structure

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Resonance Structures

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Individual Lewis structures in cases where two or more Lewis structures are equally good descriptions of a single molecule. The resonance structures in such an instance are "averaged" to give a more accurate description of the real molecule.

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Bond Enthalpy

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The enthalpy change, ΔH, required to break a particular bond when the substance is in the gas phase

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Bond angles

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the angles made by the lines joining the nuclei of the atoms in a molecule

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Valence Shell Electron-Pair Repulsion Model (VSEPR)

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A model that accounts for the geometric arrangements of shared and unshared electron pairs around a central atom in terms of the repulsions between electron pairs

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Electron Domains

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In the VSEPR model, a region about a central atom in which an electron pair is concentrated

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Bonding Pairs

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In a Lewis structure a pair of electrons that is shared by two atoms

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Nonbonding Pairs

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In a Lewis structure a pair of electrons assigned completely to one atom; also called a lone pair

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Electron-Domain Geometry

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The three-dimensional arrangement of the electron domains around an atom according to the VSEPR model

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Molecular Geometry

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the arrangement in space of the atoms of a molecule

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Bond Dipoles

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The dipole moment that is due to unequal electron sharing between two atoms in a covalent bond

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Valence Bond Theory

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A model of chemical bonding in which an electron-pair bond is formed between two atoms by the overlap of orbitals on the two atoms

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Overlap

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The extent to which atomic orbitals on different atoms share the same region of space. When the overlap between two orbitals is large, a strong bond may be formed.

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Hybrid Orbitals

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An orbital that results from the mixing of different kinds of atomic orbitals on the same atom. For example, an sp³ hybrid results from the mixing, or hybridizing, of one s orbital and three p orbitals

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Hybridization

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The mixing of different types of atomic orbitals to produce a set of equivalent hybrid orbitals

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Sigma (σ) Bonds

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a covalent bond in which electron density is concentrated along the internuclear axis

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Pi (∏) Bond

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A covalent bond in which electron density is concentrated above and below the internuclear acis

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Molecular Orbital Theory

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A theory that accounts for the allowed states for electrons in molecules

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Molecular Orbitals (MOs)

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An allowed state for an electron in a molecule. According to molecular-orbital theory, a molecular orbital is entirely analogous to an atomic orbital, which is an allowed state for an electron in an atom. Most bonding molecular orbitals can be classified as sigma or pi, depending on the disposition of electron density with respect to the internuclear axis.

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Bonding Molecular Orbital

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A molecular orbital in which the electron density is concentrated in the internuclear region. The energy of a bonding molecular orbital is lower than the energy of the separate atomic orbitals from which it forms.

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Antibonding Molecular Orbital

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a molecular orbital in which electron density is concentrated outside the region between the two nuclei of bonded atoms. Such orbitals, designated as sigma or pi, are less stable (of higher energy) than bonding molecular orbitals

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Sigma Molecular Orbitals

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a molecular orbital that centers the electron density about an imaginary line passing through two nuclei

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Energy Level Diagram

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A diagram that shows the energies of molecular orbitals relative to the atomic orbitals from which they are derived. Also called a molecular-orbital diagram.

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Molecular Orbital Diagram

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A diagram that shows the energies of molecular orbitals relative to the atomic orbitals from which they are derived. Also called an energy-level diagram.

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Bond Order

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The number of of bonding electron pairs shared between two atoms, minus the number of antibonding electron pairs: bond order = (number of bonding electrons - number of antibonding electrons)/2

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Intermolecular Forces

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The short-range attractive forces operating between the particles that make up the units of a liquid or solid substance. These same forces also cause gases to liquefy or solidify at low temperatures and high pressures.

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Dipole-Dipole Forces

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The force that exists because of the interactions of dipoles on polar molecules in close contact.

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London Dispersion Forces

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Intermolecular forces resulting from attractions between induced dipoles.

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Hydrogen Bonding

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Bonding that results from intermolecular attractions between molecules containing hydrogen bonded to an electronegative element. The most important examples involve OH, NH, and HF.

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Ion-Dipole Forces

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The force that exists between an ion and a neutral polar molecule that possesses a permanent dipole moment.

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Polarizability

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The ease with which the electron cloud of an atom or a molecule is distorted by an outside influence, thereby inducing a dipole moment.

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Viscosity

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A measure of the resistance of fluids to flow.

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Surface Tension

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The intermolecular, cohesive attraction that causes a liquid to minimize its surface area.

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Capillary Action

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The process by a which a liquid rises in a tube because of a combination of adhesion to the walls of the tube and cohesion between liquid particles.

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Phase Changes

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The conversion of a substance from one state of matter to another. The phase changes we consider are melting and freezing (solid ↔ liquid), sublimation and deposition (solid ↔ gas), and vaporization and condensation (liquid ↔ gas.)

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Heat of Fusion (Melting)

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The enthalpy change, ΔH, for melting a solid.

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Heat of Sublimation

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The enthalpy change, ΔH, for vaporization of a solid.

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Heat of Vaporization

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The enthalpy change, ΔH, for vaporization of a liquid.

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Critical Temperature

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The highest temperature at which it is possible to convert the gaseous form of a substance to a liquid. The critical temperatre increases with an increase in the magnitude of intermolecular forces.

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Critical Pressure

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The pressure at which a gas at its critical temperature is converted to a liquid state.

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Vapor Pressure

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The pressure exerted by a vapor in equilibrium with its liquid or solid phase.

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Dynamic Equilibrium

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A state of balance in which opposing processes occur at the same rate.

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Volatile

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Tending to evaporate readily.

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Normal Boiling Point

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The boiling point of 1 atm pressure.

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Alkanes

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Compounds of carbon and hydrogen containing only carbon-carbon single bonds.

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Alkenes

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Hydrocarbons containing one or more carbon-carbon double bonds.

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Alkynes

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Hydrocarbons containing one or more carbon-carbon triple bonds.

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Aromatic Hydrocarbons

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Hydrocarbon compounds that contain a planar, cyclic arrangement of carbon atoms linked by both a sigma and delocalized pi bonds.

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Structural Isomers

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Compounds possessing the same formula but differing the bonding arrangements of the atoms

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Cycloalkanes

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Saturated hydrocarbons of general formula C_nH_2n in which the carbon atoms form a closed ring

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Geometric Isomers

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Compounds with the same type and number of atoms and the same chemical bonds but different spatial arrangements of these atoms and bonds.

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