General Chemistry: Periodic Trends And Bonding

Front 1

Principal quantum number

Back 1

n, shell number: size and energy
greater n = greater distance from nucleus and higher energy level

Front 2

Angular quantum number

Back 2

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

Front 3

Magnetic quantum number

Back 3

m(l), spatial orientation of orbital
values (-l to +l)

Front 4

Spin quantum number

Back 4

m(s): any electron can have only one of two values: -1/2 or +1/2

Front 5

Pauli Exclusion Principle

Back 5

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)

Front 6

Precedence of oxidation states

Back 6

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

Front 7

oxidation number of group I and group II metals

Back 7

group I: +1, group 2: +2

Front 8

oxidation number of fluorine

Back 8

Fluorine: -1

Front 9

oxidation number of hydrogen

Back 9

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

Front 10

oxidation number of oxygen

Back 10

oxygen: -2
exception: H2O2 or Na2O2: oxygen: -2

Front 11

oxidation number of halogens

Back 11

halogens: -1

Front 12

oxidation number of group VI

Back 12

oxygen family: -2

Front 13

describe metals

Back 13

left of transition metals (except hydrogen), good conductors of heat and electricity
group 1: alkali metals
group 2A: alkaline earth metals-- very reactive

Front 14

describe transition metals

Back 14

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

Front 15

describe nonmetals

Back 15

right portion of periodic table, brittle in solid state, no metallic luster, poor conductors
group 7A: halogens
group 8A: noble/inert gases

Front 16

describe metalloids

Back 16

found along ladder, characteristics intermediate between metals and nonmetals
B, Si, Ge, As, Sb, Te, Po

Front 17

diamagnetic

Back 17

atom that has all electrons spin paired
must be even number of electrons
leave no magnetic field, are repelled by externally produced magnetic fields

Front 18

paramagnetic

Back 18

if electrons aren't all spin paired
attracted into magnetic fields

Front 19

Types of bonding: intramolecular bonds

Back 19

ionic
covalent: non-polar, polar, coordinate
metallic

Front 20

Ionic bonding

Back 20

happens between metals and nonmetals, large electronegativity difference which leads to transfer of electrons

Front 21

polar covalent bonding

Back 21

intermediate electronegativity difference leads to unequal sharing of electrons, element with higher electronegativity gets partial negative charge,

Front 22

nonpolar covalent bonding

Back 22

same electronegativity leads to equal sharing of electrons, happens mainly in diatomic molecules (H2, N2, O2, F2, Cl2, Br2, I2)

Front 23

metallic bonding

Back 23

occurs in metals and forms a sea of electrons, molecular orbitals are delocalized over entire crystal, mobility of electrons makes for good conductors

Front 24

coordinate covalent bond

Back 24

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

Front 25

intermolecular forces

Back 25

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

Front 26

ion-dipole forces

Back 26

occur between an ion and a polar molecule

common example: solvation of ionic compounds (Na+ Cl- solvated by polar water molecule)

Front 27

Hydrogen bonding

Back 27

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

Front 28

dipole-dipole forces

Back 28

occurs between polar molecules

Front 29

London forces

Back 29

occurs between nonpolar and nonpolar molecules, creates transient dipoles in nonpolar molecules

Ex: F-F and F-F

Front 30

Formal Charge

Back 30

Lewis structure with smallest formal charge is preferred structure.

FC = (#VE) - (1/2 bonding e- + nonbonding e-)

Front 31

VSEPR: 2 elecron pairs, 0 nonbonded pairs

Back 31

linear, 180 degrees

Front 32

VSEPR: 3 electron pairs, 0 nonbonded pairs

Back 32

trigonal planar, 120 degrees

Front 33

VSEPR: 4 electron pairs, 1 nonbonded pairs

Back 33

trigonal pyramidal, 107 degrees

Front 34

VSEPR: 4 electron pairs, 0 nonbonded pairs

Back 34

tetrahedral, 109.5 degrees

Front 35

VSEPR: 5 electron pairs, 0 nonbonded pairs

Back 35

trigonal bipyramidal

Front 36

VSEPR: 6 electron pairs, 0 nonbonded pairs

Back 36

octahedral