1. S° increases as temperature increases
2. S° increases as a substance changes from solid to liquid to gas 
3. S° of a dissolved solid or liquid is greater than the S° of the pure solute. This is not necessarily the case for ionic solutions since the water molecules may inhibit ions from an increased microstate. 
4. The S° of a gas decreases when it dissolves in an aqueous solution. The entropy of a gas increases however when dissolved in another gas due to the separation and mixing of gas molecules. 
5. Heavier molecules have larger S° values.
6. For allotropes the entropy is higher for for more atoms that allow more freedom of motion.
7.  S° increases with chemical complexity.
8. The physical state of a compound dominates the complexity.

1. Write the Ksp equation for the ions based on their coefficients when dissolved.
2. Locate the Ksp value and set it equal to ionic equation.
3. Plug in molarity of medal ion. 
4. Solve for hydroxide ion.

1. Write out the Ksp equation.
2. Find concentration of OH^- via 14-pH=pOH, 10^-pOH=[OH^-].
3. Determine Ksp (given in appendix).
4. Solve for ionic medal.

Percent concentration left in solution:

[(Current concentration)/(initial concentration)](100)

The equilibrium constant for a solid dissolving into ions.

Ksp = products/reactants

Since reactants are solids and solids are not included in equilibrium calculations there is no denominator and on products are multiplied and raised by their coefficient.

High spin vs. low spin complexes

High spin complexes: created by weak crystal fields, such complex ions contain the least amount of unpaired electrons.

Low spin complexes: created by strong crystal fields, such complex ions contain the maximum amount of unpaired electrons.

1. Concentration of reactants: the more frequently reactants collide the faster a reaction occurs; therefore, an increase in concentration will increase the number of collisions between reactants which will increase the rate of the reaction.
2. Physical state of reactants: the more exposed reactants are to one another the faster the reaction due to a higher number of collisions. For example, wood chips burn faster than a wooden log because there is more surface area exposed in the mulch. 
3. Temperature of environment: increasing temperature increases the number of collisions between reactants; however, more importantly, increased temperature increases the kinetic energy of the molecules that collide. Kinetic energy is of significant importance because reaction are dependent on kinetic energy. If the kinetic energy of a collision is not adequate, chemical reactions between reactants may not occur.
4. The use of a catalyst.

Reaction rate

The rate law expresses the rate as a function of reactant concentrations, product concentrations , and temperature. The components of the rate equation include the rate constant and the reaction order. The rate law is deduced by calculating the initial rate of a reaction, using initial rates from several experiments to find the reaction orders, using the values obtained from the initial rate and and reaction order to find the rate constant. 

A constant specific for a given reaction and is not changed by the stage in the reaction.

Reaction orders

Define how the rate is affected by reactant concentration. The reaction orders are not derived from a reactions stoichometry (coefficients) but from rate data. Reaction orders cannot be negative or fractions; they are zero or positive integers. 

1. A change in color due to the production or elimination of color.
2. A change in pressure if there is an increase or decrease in gaseous moles of a substance. 
3. A change in conductivity when a reactions conductivity changes as it proceeds.

1. Cu: [Ar]4s¹3d¹⁰

1. If the precipatite produces hydroxide (when in aqueous form) than the hydroxide will react with the hydronium to produce water. The reaction will proceed to the right (more precipitate will dissolve) because solute evolved into water.
2. If a component of the of the precipitate is a weak base in aqueous form than the weak base will react with hydronium to form water. The reaction will proceed to the right as more product evolved (more precipitate will dissolve).
3. If the dissolved precipitate is not a weak base (eg. strong base) the reaction will not be effected when hydronium is dissolved.

1. Write out the reaction when the precipitate dissolves.
2. Write the Ksp equation
   • Treat coffecients by multiplying and raising to the power by the coefficient.
   • Let solute components equal S.
3. Solve for S (solubility).

1. Write out the solubility.
2. Set up a reaction table with S as unknown and ion already present added to unknown.
3. Find theoretical concentrations at equilibrium.
4. Solve for S (solubility).

1. Convert solubility to molar solubility.
2. Write out reaction table based on coefficients.
3. plug into Ksp (Ksp = [2A]^2[B]) reaction equation.
4. Solve for Ksp.

1. First order integral: x
2. Second order integral: ln(x)
3. Third order integral: 1/x

1. Soluble:
   1. Group 1 alkali metals are soluble.
   2. Ammonia (NH4) is soluble.
   3. Nitrates, acetates, and chlorates are soluble.
   4. Br, Cl, F, and I are soluble except when mixed with Ag, Pb, Cu, and Hg.
   5. All sulfates (SO4) are soluble except when Ba, Sr, Ca, Ag, and Pb.
2. Insoluble compounds:
   1. All Hydroxides are insoluble except when bonded to group 1 metals or large group two metals (Ca and bigger).
   2. Phosphates, carbonates are insoluble except when bonded with group 1 metals and NH4. 
   3. Sulfides are insoluble except when bonded to group 1, 2 metals, or NH4. 
3. Remember:
   1. The larger the charge of the ions in an ionic compound, the less soluble the compound.
   2. Cues for soluble compounds: group 1 metals, 2 metals, NH4, nitrate, and sulfate, halogens (Br, I, Cl, F), acetates, perchlorates.
   3. Insoluble: Hydroxides, carbonate, phosphate, sulfide, Pb, Ag, Cu, Hg.

1. Strong acids:
   1. Halogens HBr, HCl, HI (excluding HF which is weak).
   2. Oxoacids: 2 or more oxygens than hydrogens (H2SO4, HNO3, HPO3, HClO4). 
2. Weak acids:
   1. Acids where the hydrogen is not bonded to the oxygen.
   2. Oxoacids with 1 oxygens/ 1 hydrogen ratio.
   3. Carboxylic  acids.
3. Strong bases:
   1. Hydroxide compounds.
   2. Compounds containing oxygen with a negative two charge.
4. Weak bases:
   1. Compounds with electron rich nitrogens (example: NH3 (ammonia) and amines.

([A]₀-[A]_t)/[A]₀

k=Ae^-(Ea/RT)

Ea/RT: gives the number of molecules at a given temperature that will have enough energy to react. 

A: the "pre-exponential factor"and accounts for the fact that some molecules won't form products upon interaction (for example incorrect orientation upon collision). The pre-exponential symbol expresses the number of molecules that receive an adequate amount of activation energy but do not proceed the hump to and evolve into products. 

Decay, of an unstable nucleus, by emitting radiation.

Nucleons (23.1)

A nuclear with a specific number of two types of nucleons. When an element has two isotopes, each isotope is a unique number of neutrons but the same number of protons. 

Atoms with a characteristic number of protons but a different number f neutrons. 

1. Electrons involved in chemical reactions vs. protons, neutrons and subatomic particles involved in nuclear reactions.
2. In chemical reactions substances are changed; in nuclear reactions atoms are changed.
3. Chemical reactions are accompanied by relatively small changes in energy vs. nuclear reactions are accompanied by relatively large quantities of energy.
4. Chemical reaction rates are influenced by temperature, catalysts, concentration, and the compound at hand. Nuclear reaction rates are influenced mainly by the number of nuclei and only on rare occasions by element/compound at hand. 

• An acid is any substance that has H⁺ ions in it yields hydronium in water.
• A base is a substance that has OH in its formula and dissociates in water yields OH. 

Kw = [H30+][OH-] = 1e-14

pKw = pH+pOH = 14

Arrhenius acid base definition: an acid is any molecule that contains a hydrogen ion and dissociates ions into H+ in aqueous solutions. A base is any compound containing a OH- complex that dissociates into OH- ions when in aqueous solution.

Bronstead lowry definition: An arrhenius and Bronstead lowry acid are the same. A Bronstead Lowry base accepts H+ in aqueous solutions and other situations.

Lewis acid base definition: An acid is any compound that donates accepts a lone pair of electrons. A base is any compound that donates a pair of lone electrons. 

HS- and H2S

NH3 and NH4