1. Balance the equation
2. Write the equilibrium expression
3. List the initial concentratsions
4. Calculate Q and determine the shift to equilibrium
5. Define equilibrium concentrations
6. Substitute equilibrium concentrations into equilibrium expression and solve
7. Check calculated concentrations by calculating K

Concentration: the system will shift away from the added component.

Temperature: K will change depending upon the temperature (treat the enery change as a reactant).

Pressure: a) adding of inert gas does ont affect the equilibrium position. b) decreasing the volume shifts the equilibrium toward the side with fewer moles.

Arrhenius acid

Arrhenius base

Bronsted-Lowry acid

Bronsted-Lowry base

produces H+ in aqueous solutions.

produces OH- in aqueous solutions.

H+ donor

H+ acceptor

Ionization equilibrium lies far to the right (products).

Yields a weak conugate base

K = large

Completely breaks apart in water

Ionization equilibrium lies far to the left (reactants).

Weaker the acid, the stronger its conjugate base.

Does not break apart much.

HCl _{hydrocloric acid}

HBr _{hydrobromic acid}

HI

HClO_{4 perchloric acid}

HNO_{3 nitric acid}

H₂SO_{4 sulfuric acid}

LiOH

NaOH

KOH

RbOH

CsOH

Ca(OH)₂

Sr(OH)₂

Ba(OH)₂

has a reaction with itself

Amphoteric: can behave as either an acid or base

H₂O + H₂O --> H₃O⁺ + OH^-

Kw = 1.0 x 10^-14

Acidic

Basic

Neutral

H⁺ > OH^{-      pH < 7}

OH^- > H^{+     pH > 7}

H⁺ = OH^{-       pH = 7}

Larger Ka value = stronger acid

Smaller pKa value = stronger acid

will have H+ bound to F N or O atoms that can be donated

pH = -log[H+]

pH decreases as H+ increases

The acid(base) dominates pH (pOH) if:

• its concentration is greater than 10^-6M
• (pure water would donate 10^-7M H+)
• its Ka(Kb) is larger than Kw

[H₃O⁺]/[HA] x100%

more concentrated   more diluted

<---acid conc -----

----percent dissociation-->

<---[H+] conc-------

Product of the ionization constants of an acid conjugate base pair is the ionization constant of water

Ka x Kb = Kw

pKa + pKb = pKw

Larger Kb = Stronger base

Smaller pKb = Stronger base

will have N or O atoms that have lone pair electrons that can attract H+

pOH = -log [OH-]

pOH > 7 acidic

pOH < 7 basic

pOH = 7 neutral

an acid that contatins more than one ionizable H atom per molecule

the first proton is the easiest to remove (largest Ka)

After that there is an anion that the proton would be removed from.

It gets progressively more difficult to removed succeeding protons

cations may act as acids in water

except: Li+, Na+, K+, Rb+, Cs+, Ca₂+, Sr₂+

(strong bases with hydroxide removed)

these are the conjugate acids of strong bases and are such super weak acids that they will not act as an acid in water

these are pH neutral

Anions may act as bases in water

Except: Cl-, Br-, I-, NO₃-, HSO₄-, ClO₄-, BrO₄-, IO₄-

(strong acids without H+)

these are the conjugate bases of strong acids and are such super weak bases that they will not act as a base in water

these are pH neutral

Factors effecting acid strength:

bond polarity (high is good)

bond strength (low is good)

• Contains the group H-O-X (x=everything else)
• for a given series the acid strength increases with an increase in the number of oxygen atoms attached to the central atom.
• The greater the ability of X to draw electrons toward itself, the greater the acidity of the molecule.
• the greater the electron density, the easier the proton will leave. The stronger the acid.

Molecules containing the grouping H-O-X can behave as acids

• the acid strength increases with increasing electron-withdrawing ability of X. This enables the H+ to be released. Ex. HClO₄
• If X has a very low electronegativity, the OH- can be lost and the solution will be basic Ex. NaOH

Lewis Acid

Lewis Base

electron pair acceptor

electron pair donor

Common Ion Effect

The suppression of the ionization of a weak electrolyte caused by the addition of an ion that is also a product of the ionization equilibrium of the weak electrolyte.

Shift in equilibrium position that occurs because of the addition of an ion already involved in the equilibrium reaction

an application of Le Chatelier's principle

the strong acid causes less of the weak acid to dissociate than would be expected

resist a change in pH

they are weak acids or bases containing a common ion

after addition of strong acid or base, deal with stoichiometry first, then the equilibrium

• buffers contian relatively large amounts of weak acid and corresponding conjugate base
• added H+ reacts to completion with the conjugate base
• Added OH- reacts to completion with the weak acid
• The pH in the buffered solution is deteremined by the ratio of the concentrations of the weak acid and weak base. As long as this ratio remains virtually constant. This will be the case as long as the concentrations of the buffering materials (HA and A- or B and BH+) are large compared with amounts of H+ or OH- added

the amount of acid or base that a buffer can neutralize before its pH changes appreciable

when the ratio n_base/n_acid is close to 1, the buffer has its maximum buffer capacity.

Plotting the pH of the solution being analyzed as a function of the amount of titrant added

equivalence (stoichiometric) point- point in the titration when enough titrant has been added to react exactly with the substance in solution being titrated

a substance whose color depends on the pH of the solution to which it is added.

The indicator is a weak acid. It has a different color than its conjugate base.

The equivalence point is not necessarily the same as the end point (but they are ideally as close as possible)

the color change will become apparent when about 10% of the initial from of the indicator has been converted

increase the number, more it breaks apart

decrease the number, less it breaks apart

lower ksp more solid that will dissolve

mixing two solutions of ions

Q>Ksp : precipitation occurs and will continue until the concentrations are reduced to the point that they satisfy Ksp: too many products

Q< Ksp : no precipitation occurs: too many reactants, make more products

• if the anion of the solid is a good base, the solubility is greatly increased by acidifying the solution
• in cases where the anion is not sufficiently basic, the ionic solid often can be dissolved in a solution containing a ligand that forms stable complex ions with its cation.

a change that occurs in a system left to itself; once started no external action is necessary to make the process continue.

exothermic heat given off; combustion

one that occurs without outside intervention

lets us predict whether a process will occur but gives no information about the amount of time required for the process

thermodynamic property related to the degree of disorder in a system

nature tends toward disorder.

increase if:

• liquisd are formed from solids
• gases are formed from either solids or liquids
• the number of molecules of gas increases as a result of a chemical reaction
• the temperature of a substance increases

a gas expands into a vacuum because the expanded state has the highest positional probability of states available to the system

S solid < S liquid << S gas

the change in positional entropy is dominated by the relative numbers of molecules of gaseous reactants and products

in any spontaneous proces there is always an increase in the entropy of the universe

the entropy of the universe is increases

delta S_univ> 0

where delta S_univ = delta S_sys + delta S_surr

for a spontaneous process

the sign of delta S_surr depends on the direction of the heat flow.

The magnitude of delta S_surr depends on the temperature

delta S_surr = q_rev/T = delta H/T

q_{rev = heat gained in a reversible process -Joules}

T _{= temperature (kelvin)}

the change to the system is the reverse of the change to the surroundings

delta S_{surr =} delta H/T = -delta H _sys/T

delta G = delta H - TdeltaS (from the standpoint of the system)

a process (at constant T, P) is spontaneous in the direction in which free energy decreases

a [-] delta G means [+] deltas S_univ

delta G negative spontaneous

delga G positive nonspontaneous

delta G zero equilibrium

• the entropy of a perfect crystal at 0K is zero
• because S is explicitly known (=0) at 0K, S values at other temps can be calculated

The more complex the molecule, the larger the entropy

delta G = delta G zero + RT ln(Q)

R gas law constant 8.3145 J/k*mol

T: temperature in Kelvin

Q: reaction quotient (in partial pressures)

Delta G zero: free energy change at the standard state

Q = K

maximum possible useful work obtainable from a process at constant temperature adn pressure is equal to the change in free energy.

Wmax = deltaG

• Achieving the maximum work available from a spontaneosu process can occur only via a hypothetical pathway. Any real pathway wastes energy.
• All real processes are irreversible
• First law: cant have more than start with
• Second law: b/c disorder (entropy)

Electrochemistry

involves a transfer of electrons from the reducing agent to the oxidixing agent.

oxidation

reduction

reducing agent

oxidizing agent

loss of electrons

gain of electrons

electron donor

electron acceptor

Write a separate reduction, oxidation reactions:

For each half reaction:

• Balance elements except H, O
• balance O using H₂O
• balance H using H+
• Balance charge using electrons
• multiply by interger to equalize electron count
• add half reactions
• check that elements and charges are balanaced

• balance as in acid
• Add OH- that equals H+ ions (both sides)
• Form water by cominbing H+, OH-
• Check elements and charges for balance

a device in which chemical energy is changed into electrical energy

this is done with a oxidation-reduction (redox) reaction

anode

cathode

the electrode where oxidation occurs

the electrode where reduction occurs

Voltaic Cell: cathode reaction

Voltaic Cell: anode reaction

mass increases, electrode larger

mass decrease, electrode smaller

the unit of electrical potential defined as one joule of work per coulomb of charge transferred

an instrument that measures cell potential by drawing electrical current through a known resistance

a galvanic cell consists of an oxidizing agent in one compartment that pulls electrons through a wire from a reducing agent in the other compartment

the pull or driving force on the electrons is called the cell potential (Ecell) or emf of the cell.

unit of electrical potential is the volt V- 1 joule of work per coulomb of charge transferred

all half reactions are given as reduction processes in standard tables

when a half reaction is reversed, the sign of Ezero is reversed

when a half reaction is multiplied by an integer, E zero remains the same

a galvanic cell runs spontaneously in the direction that gives a positive value for E zero cell

anode components are listed on the left. Cathode components are listd on the right.

anode and cathode are separated by double vertical lines. a phase difference is indivated by a single vertical line.

work is never the maximum possible if any current is flowing

in any real, spontaneous process some energy is always wasted- the actual work realized is always less than the calculated maximum

directly related to the free energy difference between the reactants and the products in the cell

delta G zero = -nFEzero

n=number of moles of electrons

F=Faraday= 96,485 coulombs per mole of electrons