AP Chemistry Study Guide

Inter/Intramolecular bonding:Strong Ionic (Cation + Anion)

Covalent (Electronegativity)Metallic (Fixed cation and delocalized electrons)Hydrogen (HN; HO; HF)VDW: Dipole-Dipole (Polar Dipoles)

Weak VDW: London-Dispersion Forces (Nonpolar)

Structure of solids:Ionic (Lattice) Network Covalent Amorphous

packing of cations high melting point dont have sharpinto spaces since covalent bonds melting pointshigh melting points must be broken formed due to

and low boiling points insoluble b/c of the sudden cooling ofconductive when covalent bonds liquiddissolved poor electrical melt over range ofhard but brittle conductor b/c no temperatures

mobile electrons

Molecular Molecular solids are characterized by relatively strong intramolecular bonds between the atoms that form the molecules and much weaker intermolecular bonds between these molecules. Because the intermolecular bonds are relatively weak, molecular solids are often soft substances with low melting points.

Metals:

- Delocalized electron create a sea of electrons between metallic cations- Properties of metallic solids:

are solids at room temperature - sea of electrons provide stability are ductile and malleable - sea of electrons forms crystal lattice structure, which

allows crystal lattice to slide conduct electricity in the solid and liquid phase - sea of electrons are mobile and

conduct electricity have high thermal conductivity - heat is carried through colliding sea of electrons are shiny - sea of electrons absorb and emit light over a wide wavelength range

Alloys:substitutional alloy when atoms of one metal are substituted for atoms of another metal with a similar atomic radius

- lower electrical and thermal conductivity because the substituted molecules interrupt the flow of the sea of electrons

- harder and stronger because its more difficult for a plane of atoms to slide past each other

- resulting solid remains malleable, ductile, with similar density

interstitial alloy the atomic radius of the solute

element must be less than about 60% of the atomic radius of the host metal- metallic lattice- decreased conductivity and increased strength- Resulting solid is more rigid less malleable/ductile

Solubility:Always soluble: alkali metal ions (Li+, Na+, K+, Rb+, Cs+), NH4+, NO3, ClO3, ClO4, C2H3O2Generally soluble: (mnemonics)Cl, Br, I Soluble except Hg22+, Ag+, Pb2+ (HAP)F Soluble except Ca2+, Ba2+, Sr2+, Pb2+ , Mg2+ (CBS airs at PM)SO42 Soluble except Pb2+, Ba2+, Sr2+, Ca2+ (PBS airs Cats)Generally insoluble:O2, OH Insoluble except alkali metals, and NH4+

Ca2+, Ba2+, Sr2+ (CBS) somewhat solubleCO32, PO43, S2, SO32, C2O42, CrO42

Insoluble except alkali metals and NH4+

Strong Acids:HCl hydrochloric acid HClO3 chloric acid HClO4 perchlorate acidHBr hydrobromic acidHI hydroiodic acid HIO4 periodic acidHNO3 nitric acidH2SO4sulfuric acid

Strong Bases:Alkali/Alkali Earth Metals + OHNH4OH = NH3

Enthalpy/Entropy/Gibbs:

In exothermic reactions heat energy is transferred from the system to the surroundings so the sign of H is negative because energy is lost. The energy of the products is less than that of the reactants.

In endothermic reactions heat energy is transferred from the surroundings to the system so the sign of H is positive because energy is gained. The energy of the reactants is less than that of the products.

E = E(final) E(initial) = q + w- E = energy

independent of the pathway- state function (Hesss Law)

- q = heat q is positive in endothermic reactions q is negative in exothermic reactions

- w = work w is negative if the system does work w is positive if work is done on the system

- Both w and q are pathway dependentEntropy:- Tends towards disorder- State function

2nd Law of Thermodynamics: Any spontaneous process increases disorder Spontaneous = requires no external energy

Gibbs Free Energy:- Amount of energy available for work- State FunctionH S G + always negative negative at lower temperatures+ + negative at higher temperatures+ never negative

Gasses/Gas Laws:Ideal gas law: P V = n R T

Limitations- works best with gasses at low pressures and high temperatures- relies on Kinetic Molecular Gas Theory

Gaseous molecules have no volume

Collisions between gaseous molecules are elastic (no energy is gained or lost) Gaseous molecules have no attraction with one another Average kinetic energy is related to the temperature of the molecules.

Boyles Law Charles Law Gay-Lussacs Law Avogadros LawP1V1 = P2V2 V1/T1 = V2/T2 P1/T1 = P2/T2 V1/n1 = V2/n2

Combined Gas LawP1V1/T1 = P2V2/T2

Oxidation/Reduction:

OIL RIG = Oxidation: electron is lost Reduction: electron is gained

redox - where electrons are exchanged

remember that O will always have a charge of 2-+ Voltage = + G and a spontaneous reactionelectrolysis - the breaking apart of molecules by the use of electricity

Rate:

rate1 = k[A]m[B]n Where m and n can only be determined experimentally

rate2 = k[A]m[B]n

Basic Kinetics: Determined by slowest reaction

- requiring most activation energy require this much energy to break and form bonds (KE)

Rate Laws on Graphs: Disappearance and Appearance should be proportional

- ex. NO2 disappears at 2 mol/sec while O2 appears at a rate of 1 mole/sec

Nuclear Chemistry:Alpha Decay Alpha particle = helium-4 (top 4 bottom 2)

Two protons, two neutrons bound together Low energy and relatively heavy Easily stopped by a piece of paper

Beta Decay Beta particle = electron (top 0 bottom -1) Higher energy but stopped with aluminum foil or skin

Gamma Decay High frequency electromagnetic energy Gamma symbol (top and bottom both 0) Often when electrons transition from excited state to ground state

Energy can be released into visible or UV light, X rays, or Gamma rays Can occur simultaneously with other reactions Can penetrate very deep (dangerous)

E = mc2 Where mass and energy can be used interchangeably- mass defect = difference between total mass of individual

nucleons and nucleus as a whole

Acids/Bases:

M1V1 = M2V2 Where 1 is acid and 2 is base

Arrhenius Theory: Acids are substances that produce hydrogen ions in solution Bases are substances that produce hydroxide ions in solution Neutralization occurs when hydrogen ions and hydroxide ions react to produce water

Brnsted-Lowry Theory An acid is a proton (hydrogen ion) donor A base is a proton (hydrogen ion) acceptor

Hydroxide still base because they accept hydrogen ions Conjugates

Strong [acid/base] > Weak conjugate Weak [acid/base] > Strong conjugate

Lewis Theory An acid is an electron pair acceptor A base is an electron pair donor Electrons come from hydrogen ion and bonds with other ions