Organic 1: Structure Determines Properties Flashcards
Atomic number Z
each element is characterized by this, which is equal to the number of protons in its nucleus
Wave functions
mathematical descriptions of the electron wave, symbolized by Greek letter psi; AKA orbitals
Psi^2
probability of finding an electron at a particular spot relative to an atom’s nucleus = psi^2 at that point
Principal quantum number n
letter s is preceded by this, which specifies the shell and is related to the energy of the orbital
Shell
group of orbitals that have the same principal quantum number n
Boundary surfaces
more common to represent orbitals by these; encloses the region where the probability of finding an electron is high (90-95%)
Spin
synonymous with spin quantum number
Spin quantum number
can have a value of +1/2 or -1/2
Pauli Exclusion Principle
two electrons may occupy the same orbital only when they have opposite (paired) spins; for this reason, no orbital can contain more than 2 electrons
Period
row of the periodic table; corresponds to the principal quantum number of the highest numbered occupied orbital
Nodal surfaces
regions of a single orbital may be separated by these where the wave function changes sign and the probability of finding an electron is zero
Hund’s Rule
general principal for orbitals of equal energy; when 2 orbitals are of equal energy, they are populated by electrons so that each one is half-filled before either one is doubly occupied
Valence electrons
outermost electrons, the ones most likely to be involved in chemical bonding and reactions
Valence shell
the group of orbitals, filled and unfilled, responsible for the characteristic chemical properties of an atom
Main-group elements
for these, the number of valence electrons is equal to its group number in the periodic table
Octet
having 8 electrons in the valence shell
Noble (rare) gases
helium, neon, and argon are in this class; characterized by extremely stable, “closed-shell” electron configuration; very unreactive
Compounds
atoms combine with one another to give these, having properties different from the atoms they contain
Chemical bond
attractive force between atoms in a compound
Ionic bond
force of attraction between oppositely charged species (ions)
Cations
positively charged ions
Anions
negatively charged ions
Isoelectronic
species that have the same number of electrons
Ionization energy
large amount of energy that must be transferred to any atom to dislodge an electron
Endothermic
processes that absorb energy
Exothermic
energy-releasing reactions; energy change for this process has a negative sign
Electron affinity
energy change for addition of an electron to an atom
Electrostatic
attractive forces between oppositely charged particles; AKA coulombic attractions; what is meant by an ionic bond between 2 atoms
Covalent (shared electron pair) model
first suggested by GN Lewis (1916); proposed that a sharing of 2 electrons by 2 hydrogen atoms permits each one to have a stable, close-shell electron configuration analogous to helium
Lewis structures
structural formulas in which electrons are represented as dots; customary to represent a shared electron-pair bond by a dash
Bond dissociation enthalpy
amount of energy required to dissociate a hydrogen molecule H2 to two separate hydrogen atoms
Unshared pairs
Valence electrons not involved in bonding
Octet rule
in forming compounds, elements gain, lose, or share electrons to achieve a stable electron configuration characterized by 8 valence electrons
Double bond
bond formed by the sharing of 4 electrons between 2 atoms
Triple bond
bond formed by the sharing of 6 electrons between 2 atoms
Polar covalent
if one atom has a greater tendency to attract electrons toward itself than the other, the electron distribution if polarized and the bond is described at this
Electronegativity
tendency of an atom to attract the electrons in a covalent bond toward itself
Electrostatic potential map
uses the colors of the rainbow to show charge distribution
Bond dipole moments
exists whenever opposite charges are separated from each other; direction of this is toward the more electronegative atom
Dipole moment mu
product of the amount of the charge e multiplied by the distance d between the centers of charge (mu = e x d)
debye (D)
unit customarily used for measuring dipole moments; 1D = 1 x 10^-18 esu cm
Formal charges
the charge, either positive or negative, of an atom calculated by subtracting the number of valence electrons in the neutral atom a number equal to the sum of its unshared electrons plus half the electrons in its covalent bonds
Molecular formula
tells us which atoms and how many of each are present in a compound
Connectivity
order in which atoms are connected
Isomers
different compounds that have the same molecular formula
Constitutional isomers
isomers that differ in connectivity; AKA structural isomers
Stereoisomers
isomers that differ in arrangement of atoms in space
Resonance
when 2 or more Lewis structures that differ only in the distribution of electrons can be written for a molecule, no single Lewis structure is sufficient to describe its true electron distribution
Resonance hybrid
the collection of Lewis structures that, taken together, represent the electron distribution in a molecule
Contributing structures
the various resonance structures that can be written for a molecule
Localized
Lewis formulas show electrons as this; they either are shared between 2 atoms in a covalent bond or are unshared electrons belonging to a single atom
Delocalized
electrons shared by several nuclei
Curved arrows
writing the various Lewis formulas that contribute to a resonance hybrid can be made easier by using these to keep track of delocalized electrons; converts 1 Lewis structure to another by moving electron pairs using these
Condensed formulas
leave out some, many, or all of the covalent bonds and use subscripts to indicate the number of identical groups attached to a particular atom
Bond-line formula
AKA carbon skeletal diagram; we assume that there is a carbon atom at every vertex and at the end of a line
Heteroatoms
atoms that are neither carbon nor hydrogen
Solid wedge
represents a bond that projects toward you
Dashed wedge
represents a bond that points away from you
Simple line
represents a bond that lies in the plane of the paper
Valence shell electron-pair repulsion (VSEPR) model
rests on the idea that an electron pair, either a bonded pair or an unshared pair, associated with a particular atom will be as far away from the atom’s other electron pairs as possible
Tetrahedral angle
permits four bonds to be maximally separated; characterized by angles of 109.5 degrees; ex. water, ammonia, methane
Molecular dipole moment
resultant of all of the individual bond dipole moments of a substance
Double-barbed arrow
shows the movement of a PAIR of electrons, either a bonded pair or a lone pair
Single-barbed (fishhook) arrow
shows the movement of 1 electron
Acid
substance that ionizes to give protons when dissolved in water
Base
ionizes to give hydroxide ions
Acidity constant
equilibrium constant Ka; measures the strength of a weak acid
pKa
-log (base 10) Ka
Bronsted-Lowry acid
proton donor
Bronsted-Lowry base
proton acceptor
Conjugate acid
base and this differ by a single proton
Conjugate base
acid and this differ by a single proton
Oxonium ion
systematic name for the conjugate acid of water (H3O+)
Hydronium ion
common name for H3O+
Basicity constant Kb
Bronsted-Lowry approach involving conjugate relationships between acids and bases make this unnecessary
Inductive effects
structural effects that are transmitted through bonds
Strong acid
one that is stronger than H3O+
Weak acid
one that is weaker than H3O+
Strong base
one that is stronger than OH-
Lewis acid
electron-pair acceptor
Lewis base
electron-pair donor
Lewis acid/Lewis base complex
species that results by covalent bond formation between a Lewis acid and a Lewis base
Substitution
one atom or group replaces another in a reaction
Nucleophiles
nucleus seekers; Lewis bases that use an unshared pair to form a bond to some other atom
Electrophiles
electron seekers; Lewis acids
