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36 Cards in this Set

  • Front
  • Back
Principal quantum number
n, shell number: size and energy
greater n = greater distance from nucleus and higher energy level
Angular quantum number
l, shape of subshell
values 0-(n-1)
l=0: s, one orbital, 2 electrons
l=1: p, three orbitals, 6 electrons
l=2: d, five orbitals, 10 electrons
l=3: f, seven orbitals, 14 electrons
Magnetic quantum number
m(l), spatial orientation of orbital
values (-l to +l)
Spin quantum number
m(s): any electron can have only one of two values: -1/2 or +1/2
Pauli Exclusion Principle
No two electrons in an atom can have the same four quantum numbers. Each orbital can hold 2 electrons (one with positive, one with negative spin)
Precedence of oxidation states
1) standard state is zero
2) sum of states equal ion's charge
3) Group I and II
4) Fluorine
5) Hydrogen
6) Oxygen
7) Halogens
oxidation number of group I and group II metals
group I: +1, group 2: +2
oxidation number of fluorine
Fluorine: -1
oxidation number of hydrogen
Hydrogen: +1 when bonded to something more electronegative than carbon (FONClBrISCH)
Hydrogen: -1 when bonded to atom less electronegative than carbon
Hydrogen: 0 when bonded to C
oxidation number of oxygen
oxygen: -2
exception: H2O2 or Na2O2: oxygen: -2
oxidation number of halogens
halogens: -1
oxidation number of group VI
oxygen family: -2
describe metals
left of transition metals (except hydrogen), good conductors of heat and electricity
group 1: alkali metals
group 2A: alkaline earth metals-- very reactive
describe transition metals
partly filled or filled d orbitals, form coordination complexes which can be colored

elements in d4 and d9 promote electron from s orbital to achieve 5d or 10d configuration (more stable)

always ionize from s electrons first, after s electrons are all gone ionize from d electrons
describe nonmetals
right portion of periodic table, brittle in solid state, no metallic luster, poor conductors
group 7A: halogens
group 8A: noble/inert gases
describe metalloids
found along ladder, characteristics intermediate between metals and nonmetals
B, Si, Ge, As, Sb, Te, Po
atom that has all electrons spin paired
must be even number of electrons
leave no magnetic field, are repelled by externally produced magnetic fields
if electrons aren't all spin paired
attracted into magnetic fields
Types of bonding: intramolecular bonds
covalent: non-polar, polar, coordinate
Ionic bonding
happens between metals and nonmetals, large electronegativity difference which leads to transfer of electrons
polar covalent bonding
intermediate electronegativity difference leads to unequal sharing of electrons, element with higher electronegativity gets partial negative charge,
nonpolar covalent bonding
same electronegativity leads to equal sharing of electrons, happens mainly in diatomic molecules (H2, N2, O2, F2, Cl2, Br2, I2)
metallic bonding
occurs in metals and forms a sea of electrons, molecular orbitals are delocalized over entire crystal, mobility of electrons makes for good conductors
coordinate covalent bond
weak covalent bond between transition metal and ligand where one atom donates both of the shared electrons in a bond.

The molecule that donates the pair of electrons is known as a lewis base or ligand, molecule that accepts is known as a lewis acid
intermolecular forces
liquids and solids are held together by intermolecular forces that are weaker than ionic and covalent interactions.

ion-dipole forces>hydrogen bonding>dipole-dipole>London dispersion forces
ion-dipole forces
occur between an ion and a polar molecule

common example: solvation of ionic compounds (Na+ Cl- solvated by polar water molecule)
Hydrogen bonding
occurs between molecules that have hydrogen bound to a highly electronegative atom (H,F,O,N)

tend to have high boiling points

important in behavior of water, alcohols, amines, carboxylic acids
dipole-dipole forces
occurs between polar molecules
London forces
occurs between nonpolar and nonpolar molecules, creates transient dipoles in nonpolar molecules

Ex: F-F and F-F
Formal Charge
Lewis structure with smallest formal charge is preferred structure.

FC = (#VE) - (1/2 bonding e- + nonbonding e-)
VSEPR: 2 elecron pairs, 0 nonbonded pairs
linear, 180 degrees
VSEPR: 3 electron pairs, 0 nonbonded pairs
trigonal planar, 120 degrees
VSEPR: 4 electron pairs, 1 nonbonded pairs
trigonal pyramidal, 107 degrees
VSEPR: 4 electron pairs, 0 nonbonded pairs
tetrahedral, 109.5 degrees
VSEPR: 5 electron pairs, 0 nonbonded pairs
trigonal bipyramidal
VSEPR: 6 electron pairs, 0 nonbonded pairs