Chapter 1 Glossary

Front 1

acid-dissociation constant (Ka)

Back 1

The equilibrium constant for the reaction of the acid with water to generate H₃O⁺. (p. 23) The negative logarithm of Ka is expressed as pKa: pKa = -log₁₀Ka

Front 2

Bronsted-Lowry acid:

Back 2

proton donor

Front 3

Bronsted-Lowry base:

Back 3

proton acceptor

Front 4

Lewis acid:

Back 4

electron-pair acceptor (electrophile)

Front 5

Lewis base:

Back 5

electron-pair donor (nucleophile)

Front 6

conjugate acid

Back 6

The acid that results from protonation of a base. (p. 23)

Front 7

conjugate base

Back 7

The base that results from loss of a proton from an acid. (p. 25)

Front 8

covalent bonding

Back 8

Bonding that occurs by the sharing of electrons in the region between two nuclei. (p. 7)

Front 9

single bond

Back 9

A covalent bond that involves the sharing of one pair of electrons. (p. 8)

Front 10

double bond

Back 10

A covalent bond that involves the sharing of two pairs of electrons. (p. 8)

Front 11

triple bond

Back 11

A covalent bond that involves the sharing of three pairs of electrons. (p. 8)

Front 12

curved-arrow formalism

Back 12

A method of drawing curved arrows to keep track of electron movement from nucleophile to electrophile (or within a molecule) during the course of a reaction. (p. 31)

Front 13

degenerate orbitals

Back 13

Orbitals with identical energies. (p. 4)

Front 14

dipole moment (μ)

Back 14

A measure of the polarity of a bond (or a molecule), proportional to the product of the charge separation times the bond length. (p. 10)

Front 15

electron density

Back 15

The relative probability of finding an electron in a certain region of space. (p. 3)

Front 16

electronegativity

Back 16

A measure of an element's ability to attract electrons. Elements with higher electronegativities attract electrons more strongly. (p. 10)

Front 17

electrophile

Back 17

An electron-pair acceptor (Lewis acid). (p. 29)

Front 18

electrostatic potential map (EPM)

Back 18

A computer-calculated molecular representation that uses colors to show the charge distribution in a molecule. In most cases, the EPM uses red to show electron-rich regions (most negative electrostatic potential) and blue to show electron-poor regions (most positive electrostatic potential). The intermediate colors orange, yellow, and green show regions with intermediate electrostatic potentials. (p. 10)

Front 19

empirical formula

Back 19

The ratios of atoms in a compound. (p. 20) See also molecular formula.

Front 20

formal charges

Back 20

A method for keeping track of charges, showing what charge would be on an atom in a particular Lewis structure. (p. 11)

Front 21

Hund's rule

Back 21

When there are two or more unfilled orbitals of the same energy (degenerate orbitals), the lowest-energy configuration places the electrons in different orbitals (with parallel spins) rather than paired in the same orbital. (p. 6)

Front 22

ionic bonding

Back 22

Bonding that occurs by the attraction of oppositely charged ions. Ionic bonding usually results in the formation of a large, three-dimensional crystal lattice. (p. 7)

Front 23

isotopes

Back 23

Atoms with the same number of protons but different numbers of neutrons; atoms of the same element but with different atomic masses. (p. 3)

Front 24

Lewis structure

Back 24

A structural formula that shows all valence electrons, with the bonds symbolized by dashes ( − ) or by pairs of dots, and nonbonding electrons symbolized by dots. (p. 7)

Front 25

line-angle formula (skeletal structure, stick figure)

Back 25

A shorthand structural formula with bonds represented by lines. Carbon atoms are implied wherever two lines meet or a line begins or bends. Atoms other than C and H are drawn in, but hydrogen atoms are not shown unless they are on an atom that is drawn. Each carbon atom is assumed to have enough hydrogens to give it four bonds. (p. 19)

Front 26

lone pair

Back 26

A pair of nonbonding electrons. (p. 7)

Front 27

molecular formula

Back 27

The number of atoms of each element in one molecule of a compound.

Front 28

empirical formula

Back 28

simply gives the ratios of atoms of the different elements.

Front 29

node

Back 29

A region in an orbital with zero electron density. (p. 4)

Front 30

nodal plane

Back 30

A flat (planar) region of space with zero electron density. (p. 4)

Front 31

nonbonding electrons

Back 31

Valence electrons that are not used for bonding. A pair of nonbonding electrons is often called a lone pair. (p. 7)

Front 32

nucleophile

Back 32

An electron-pair donor (Lewis base). (p. 29)

Front 33

octet rule

Back 33

Atoms generally form bonding arrangements that give them filled shells of electrons (noble-gas configurations). For the second-row elements, this configuration has eight valence electrons. (p. 6)

Front 34

orbital

Back 34

An allowed energy state for an electron bound to a nucleus; the probability function that defines the distribution of electron density in space. The Pauli exclusion principle states that up to two electrons can occupy each orbital if their spins are paired. (p. 3)

Front 35

organic chemistry

Back 35

The chemistry of carbon compounds

Front 36

pH

Back 36

A measure of the acidity of a solution, defined as the negative logarithm (base 10) of the H3O+ concentration

Front 37

Equation for pH

Back 37

pH= -log₁₀[H₃O⁺] (p. 22)

Front 38

polar covalent bond

Back 38

A covalent bond in which electrons are shared unequally. A bond with equal sharing of electrons is called a nonpolar covalent bond. (p. 9)

Front 39

resonance forms or resonance structures.

Back 39

two or more valid Lewis structures differing only in the placement of the valence electrons. Individual resonance forms do not exist, but we can estimate their relative energies.

Front 40

major contributors

Back 40

The more important (lower-energy) resonance structures

Front 41

minor contributors

Back 41

The less important (higher-energy) resonance structures

Front 42

delocalized

Back 42

When a charge is spread over two or more atoms by resonance

Front 43

complete structural formula

Back 43

shows all the atoms and bonds in the molecule.

Front 44

condensed structural formula

Back 44

shows each central atom along with the atoms bonded to it.

Front 45

line-angle formula (sometimes called a skeletal structure or stick figure)

Back 45

assumes that there is a carbon atom wherever two lines meet or a line begins or ends.

Front 46

valence

Back 46

The number of bonds an atom usually forms. (p. 9)

Front 47

valence electrons

Back 47

Those electrons that are in the outermost shell. (p. 6)