Chemistry Cool

When excited electrons return to ground state

energy is released

Bohr’s theory

Electrons travel in fixed-distance orbits around the nucleus

Electrons have only specific

discrete energies

Bond formation

Always releases energy.

Bond breaking

Always takes energy.

Lewis theory of bonding

A chemical bond involves atoms sharing valence electrons.

Covalent bond

A chemical bond where electrons are shared between atoms.

Crystalline lattice

Material with a long-range

Amorphous lattice

Material with no overall order in the extended structure.

Metals tend to

Lose electrons.

Nonmetals tend to

Gain electrons.

Electronegativity

The ability of an atom to attract electron density from other atoms in a molecule.

Difference in electronegativity > 2.0

Mostly ionic bond.

Difference in electronegativity 0.5 - 2.0

Polar covalent bond.

Difference in electronegativity < 0.5

Non-polar covalent bond.

Molecular dipole moment

Dipole associated with the entire molecule (vector addition).

Polar molecule

Has a net direction electrons are moving (net dipole moment > 0).

Nonpolar molecule

Has a vector total (net dipole moment) equal to 0.

Ion-dipole force

Attraction between a charged ion and a molecule with a permanent dipole.

Hydrogen bonding

Attraction between H (bonded to N

Dipole-dipole force

Attraction between two molecules with permanent dipoles.

London forces (Dispersion forces)

Forces arising from a polarizable electron cloud forming a temporary dipole; all molecules exhibit this.

NH+ 4

Ammonium

CH3CO- 2

Acetate

NO- 3

Nitrate

CO2- 3

Carbonate

SO2- 4

Sulfate

PO3- 4

Phosphate

ClO- 3

Chlorate

OH-

Hydroxide

Binary ionic compound

Composed of a metal and a nonmetal (Ex: LiCl

Binary covalent compound

Composed of two nonmetals or metalloids (Ex: CO

Naming binary covalent compounds rule

Change the second element’s name to end in “ide”.

SF6

Sulfur hexafluoride

N2O5

Dinitrogen pentoxide

HCl

Hydrogen monochloride (or hydrochloric acid in solution)

CO

Carbon monoxide

NO2

Nitrogen dioxide

N2O

Dinitrogen monoxide

Naming acids without oxygen

Add the prefix hydro- and the suffix -ic acid.

Naming acids with oxygen

if name ends in -ite

Naming acids with oxygen

if name ends in -ate

HF

Hydrofluoric acid

HCN

Hydrocyanic acid

H2S

Hydrosulfuric acid

H2SO3

Sulfurous acid

HNO2

Nitrous acid

H2SO4

Sulfuric acid

HNO3

Nitric acid

VSEPR Method

Valence Shell Electron Pair Repulsion method for determining molecular shape.

Steric number

The number of electron groups around the central atom (lone pairs and multiple bonds each count as one group).

PBr3 shape

Pyramidal.

Synthesis reaction

Two or more substances react to form a single new substance (Two+ reactants

Decomposition reaction

A single compound breaks down into two or more simpler products (One reactant

Single replacement reaction

One element replaces a second element in a compound (Use activity series for metals).

Double replacement reaction

A chemical change involving an exchange of positive ions between two compounds.

Combustion reaction

A compound reacts with oxygen (O2) producing energy; often produces CO2 and H2O.

Precipitation reaction

Solid formation from mixing two or more solutions.

Precipitate

The solid that forms in a precipitation reaction.

Dissociation

Breaking apart in solution (e.g.

Soluble in water

All sodium (Na+)

Soluble in water

All acetates and nitrates.

Soluble in water (Halides exception)

Halides of lead(II)

Soluble in water (Sulfates exception)

Sulfates of calcium