ap chem 2020 exam (!!)

ksp never have denominators bc

the reactant is always solid, and solids arent included in equil. eqs

small ksp and large ksp meaning

small= not much dissolves
large= alot dissolves

solubility product

(Ksp) equilibrium constant, has only one value for a given solid at a given temp

solubility

an equilibrum position

solubility of compound

g/l

molar solubility

mol/l

calculating ksp

1.Write a balanced equation for dissolution process
2.Define change in concentration based on ICE
3.Plug results into Ksp expression and solve.

according to ksp

everything is a least a little soluble

comparing ksp value rules

1)Salts that produce the same # of ions can be compared to see which one is the most/least soluble.
2)If the salts break into different # of ions, you can't just look. Must calculate the solubility to determine which has the highest concentration of ions.

for solids as temp increases

solubility increases

ph and solubility

increase in solubility as solution becomes more acidic

common-ion effect

-If one of the ions in a solution equilibrium is already dissolved in the solution, the equilibrium will shift to the left and the solubility of the salt will decrease. Addition of ion (product) shifts equilibrium to the left (chateliers principle). commo

preciptation formation

-compare q and ksp values. Q= ion product
1. Not equilibrium or maybe it is. Not sure, use Q.
2. Need to find new [ ] after mixing and plug into Q
3. compare. -If Q = Ksp, the system is at equilibrium and the solution is saturated.
-If Q < Ksp, more solid

ksp represents

the minimum concentration of ions that will precipitate

ksp value table

Qualitiative Analysis for metallic elements

1. separate ions into broad
groups by solubility
2. specific ions are then
separated from the
group and identified
by specific tests

selective preciptation (mixtures of metal ions)

�Use a reagent whose anion forms a precipitate with only one or a few of the metal ions in the mixture.
�Example:
-Solution contains Ba2+ and Ag+ ions.
-Adding NaCl will form a precipitate with Ag+ (AgCl), while still leaving Ba2+ in solution.
-Use solubi

complex ion

a charged species composed of:
1. metallic cation
2. Ligands - Lewis bases that have a lone electron pair that can form a covalent bond with an empty orbital belonging to the metallic cation

Compounds that are insoluble in water, may be

soluble in a solvent that allows for the formation of a complex ion

common ligands

NH3, CN-, SCN-

coordination number

-the number of ligands attached to the cation
-2, 4, and 6 are the most common coordination numbers

complex ions and solubility

if reaction wants to go towards products

-Gibbs must be NEGATIVE!!!!

If the reaction wants to go towards reactants (in reverse)

-Gibbs must be POSITIVE

at equilibrum

delta g = 0

?Go is the

standard free-energy change
(i.e., for a reaction at standard conditions).

Under any other conditions... use

equilibrium point occurs

at the lowest value of free energy available to the reaction system

At equilibrium

DG = 0 and Q = K

table of standard delta g vs k

equations to use

equations to use #2

when to use what equations

�?G tells us spontaneity at current conditions. When will it stop?

�It will go to the lowest possible free energy which may be an equilibrium.
�At equilibrium ?G = 0, Q = K
�?G� = -RTlnK

temperature dependence of k

photoelectron spectroscopy

energy measurement of electrons emitted from solids, gases or liquids by the photoelectric effect, in order to determine the binding energies of electrons in a substance

photoelectric effect

The emission of electrons from a material when light of certain frequencies shines on the surface of the material

Electronegativity and trend

the ability of an atom to attract electrons when the atom is in a compound. increases to the right, decreases down

electronegativity on the periodic table explained

across a period: If the valence shell of an atom is less than half full, it requires less energy to lose an electron than to gain one. Conversely, if the valence shell is more than half full, it is easier to pull an electron into the valence shell than to

atomic radii trend

Decrease from left to right across a period, increase from top to bottom in a group or family

atomic radii trend explained

An atom gets larger as the number of electronic shells increase (explains down a group). there is a greater nuclear attraction as protons increase, meaning that the nucleus attracts the electrons more strongly, pulling the atom's shell closer to the nucle

ionization energy

The amount of energy required to remove an electron from an atom

ionization energy trend explained

Elements on the left side of the periodic table have low ionization energies because of their willingness to lose electrons and become cations. visa versa. There are more protons in atoms moving down a group (greater positive charge), yet the effect is to

electron affinty

the energy needed to remove an electron from a negative ion to form a neutral atom. (generally exothermic). opposite of electorngativity

Hyrdogen Bonding rules and why is it so strong?

N,O,F. the small size of the hydrogen atom allows the approach of dipoles to be very close and the polarity of the bond are also why its so strong.

4 types of solids

molecular, ionic, metallic, network covalent

molecular solid

a solid composed of molecules. held together by intermolecular forces (dipole-dipole and hydrogen bonding -polar and London dispersion-nonpolar)

ionic solid

a crystalline solid held together by ionic bonds. type of bonding: ionic-dipole.

metallic solid

metallic atoms, only has metals. type of bonds: metallic bonding/delocalized covalent.

network covalent

nonmetal atoms. type of bonds: covalent bonding (intra and inter are covalent)/directional covalent (leading to giant molecules). only 4 types: diamond, graphite, silicon dioxide (SiO2), silicon, sic

intramolecular forces

covalent, ionic, and metallic

intramolecular forces straight lowest to highest

non polar covalent, polar covalent, ionic, and metallic

melting point solids from highest to lowest

high melting point: ionic, network covalentmetallic is variable.low melting point: molecular

intermolecular forces weakest to strongest

LD, dipole-dipole, hydrogen bonding, and ion-dipole

vapor pressure, boiling point, and freezing points and intermolecular forces relationship

vapor pressure: The stronger the intermolecular forces, the lower the vapor pressurefreezing point: Molecules with stronger intermolecular forces are pulled together tightly to form a solid at higher temperatures, so their freezing point is higher. Molecu

heat of fusion

Amount of energy required to change a substance from the solid phase to the liquid phase.

heat of vaporization

The amount of energy required for the liquid at its boiling point to become a gas

vapor pressure of liquid in barometer

p vapor = p atmosphere - p Hh column

vapor pressure of a liquid

a measure of the force exerted by a gas above a liquid. this is when the system is at eqiliubrum.

vapor pressurs and temp

increases with temp

4 types of intermolecular forces (strongest to weakest)

ionic, covalent, dipole dipole (hydrogen bonding), and LDF

strong IMFS (BP and FP)

high BP, high MP/FP

LDF's increase with....

molar mass

look for to define molecular solids

diatomic elements, nonmetals, compounds formed from two or more nonmetals

characteristics of molecular solids

-soft (bc of weak IMFS)-low melting points-insulator and do not conduct electricity well-stronger intra than inter

Ionic solids characteristics

High melting points, high boiling points, and poor electrical conductivity in the solid state but high conductivity in the molten or aqueous state. brittle. low vapor pressure.Due to strong electrostatic interactions

greater the ions/charges

greater electrostatic forces.ex: cacl2 vs nacl

smaller the ions

greater the attraction.ex: kbr vs lif

metallic solid characteristics

malleability (able to be hammered/bent into a sheet), ductility (ability to be stretched into a wire) INSTEAD of brittle. melting points = high. conductor as a liquid and solid.

alloy

A mixture of metals. also have a sea of electrons. tend to be stronger than main metals so its less durable and malleable.

characteristics of network covalent solids

-very hard-high boiling points-not soluble, volatile, or conductive

how to find strongest LDF

which one has biggest radius/mass(more polarizable)

how to find strongest dipole-dipole

which is more polar?

how to find strongest hydrogen bonding

which has more h -bonds

how to find strongest ionic bonds

charge first / amount, then size

Condensation

achieves a dynamic equilibrium with vaporization in a close system

Electronegativity difference in bonding atoms = 0

covalent

Electronegativity difference in bonding atoms = intermediate

polar covalent

Electronegativity difference in bonding atoms = large

ionic

ionic bond vs covalent bond

Ionic- transfer of electronsCovalent- sharing of electrons to become more stable

bond energy

the amount of energy that will break a bond between two atoms

explain a graph of covalent bonds

as the atoms approach each other, the energy decreases until the distance reaches a good point and then begins to increase again due to repulsions.

polar covalent bond

A covalent bond in which electrons are not shared equally but one atom isn't strong enough to completely remove an electron from another.

lattice energy

the change in energy that takes place when separated gaseous ions are packed together to form an ionic solid. the energy released when an ionic solid forms from its ions.

enthalpy change (triangle H)

energy required (bonds broken) - energy released (bonds formed)

formal charge

number of valance electrons on the free atom - the number of valance electrons assigned to the atom in the molecule

assigned valance electrons

number of lone pair electrons + 1/2(number of shared electrons)

form vs breaking bonds

form = exothermic
break= endo

electron pair geometry

these are the 6 basic types (linear, trigonal planar, tetrehdral, etc)

molecular geometry

all of the types of shapes within the 6 types, this is the shape of the molecule or structure

Resonance

where there are multiple valid lewis strtuctures for a molecule

Resonence bonds vs single bonds

shorter and stronger

Resonence bonds vs double bonds

longer and weaker

sigma and pi bond rules

single bonds between atoms are always sigma bonds. Double bonds are comprised of one sigma and one pi bond. Triple bonds are comprised of one sigma bond and two pi bonds.

hybridization rules

when an atom is surrounded by 3 effective pairs: sp2
4:sp3
5:dsp3
6:d2sp3

real gases only behave ideally at

high temps and low pressure

problems with kinetic molecular theory

1. do occupy space
2. do experience forces

equation when ur given density, temp, and pressure

p / (t d) = p2 / (t2 d2)

for real gases, equation is P = nrt/ v-nb - n^2a/v^2. what does a and b represent?

both are constants. a is large when interparticle attractions are large (depends on size and polarity) . b is large for large gas particles.

Kinetic Molecular Theory

1. particles don't have volume
2. particles are in constant motion. the collisions of the particles with the walls of the container are the cause of the pressure
3. the particles are assumed to exert no forces on each other; they are assumed neither to at

what does kinetic molecular theory try to explain

the properties of an ideal gas, pv=nrt (ideal gas law) is a product of it

Ideal Gas Law

the relationship PV=nRT, which describes the behavior of an ideal gas

Solvent

A liquid substance capable of dissolving other substances. most in the soultion.

molality

moles of solute/kg of solvent

which solutions would have the same boiling point as .040 m C2h2 in water? a. .010 NA3Po4 2. .020 CaBr2 3. .020 m KCl 4. .020 m HF

a and c. bc .040 bc all ions dissociate, this increases stregth of intermolecular forces

freezing point depression

the difference in temperature between the freezing point of a solution and the freezing point of the pure solvent

boling point

the temperature at which the vapor pressure of the liquid is just equal to the external pressure

mass percent

mass of solute/ mass of solution

ppm, ppb, ppt

-parts per million
-parts per billion
-parts per trillion

1 ppm =

1 mg/L

molarity changes with

temperature, (volume changes with temp)

mololity doesn't change with

temperature (mass doesn't change with temp)

crytallization

opposite of solution process. goes from solution to solute and solvent.

when rates of solution and crytizialltion are equal

equilibrium is reached

Solubility

A measure of how much solute can dissolve in a given solvent in order to formed a saturated solution at a given temperature and pressure.

Saturated

soln is equal with undissolved solute, there is solid at the bottom

Unsaturated

more solute can be dissolved

Supersaturated

increase the concentration of (a solution) beyond saturation point. has clear appearance, but is very unstable. causes excess solute to crystallize. (too much dissolved)

as imfs between solute and solvent increase, solubility

increase

Miscible

Describes two liquids that are soluble in each other

as molar mass increases, polarity

decreases

pressure effect when dissolving gases

the dissolved gas is at equilibrium with the gas above the liquid and the gas in the liquid. if you increase pressure, the molecules dissolve faster and this equilibrum is disturbed. the system is now reaches a new equilibrium but with more gas molecules

does pressure have solubility effects on solids

no

temp effect on solid solubility

increases as other increases

tempt effect on gas solubility

rate decreases as temp increases, but we cant predict if temp effects how much can be dissolved

as pressure increases, gas solubility

increases

supersaturated vs saturated on a graph

solution formation

delta h/ step one: solute particles are separated from crystal lattice (intramolecular forces are broken). h is positive. this could be the lattice energy is compound is ionic.delta h/step two: solvent particles separate to make room for solute particles

if delta h of the solution is to positive

soln won't form

if molecules can attract each other vs can't

cant: h3 is small and negative
can: h3 is large and negative

Spontaneous processes occur

when h is negative, when disorder of system increases. disorder of system is more influnital to see a soln forms

chemical reaction vs solution formed

when evaporates, solid should be the same u started with

Roult's Law

x solvelent times pressure of solvent = vapor pressure soultion. only applies where solute doesn't contribute to the vapor pressure

ideal soltions that obey rheats law

low concentration, solute and solvent particles are silimar in size and IMFS

a nonvolatile solute ___ the vapor pressure of the solvent

lowers bc now there is more intermolecular forces to overcome

aquesous solution and pure water in close system

volume of water decreases, volume of solution increases. this is bc vapor pressure is less in pure water, so it evaporates faster. as this happens, the soultion absorbs the vapor to try and lower the vapor pressure

solutions

are called homogenous mixtures

chromatrography

-cant be separated by flirtation, take advangtage of imfs
-larger particles will move more slowly
-partiles with a polarity that matches the solvent, it will move further
-paper is nonpolar

chromatrography mobile and solvent phase

mobile phase moves through stationary phase" means

as it goes up the paper

solvent front

the furthest point reached by the solvent.

retention factor (R�)

distance travelled by sample/distance travelled by solvent

Distallation

-seperates parts of a mixture based on different boiling points
-heat mixture to a temp that is inbetween the two boiling points of the liquids
-lower boiling point will evaporate then condensate into separate tube

similar imfs =

soluble in each other

long wavelength

low frequency, low energy

photoelectric effect

1. only light at or above a threshold frequency will cause electrons to be ejected from a metal surface.
2. once threshold is reached, then the intensity (brightness) of the light determines how many electrons are admitted

emission of light within elements

-shows that elements only accept certain amounts of energy

Z (effective nuclear charge)

protons- number of inner electrons

According to the photoelectric effect

frequency of light determines if electrons are ejected or not and with what KE. brightness of light determines how many electrons are ejected

number of electrons ejected depends on light intensity so long as the light is

above a min energy ( this is also the ionization energy for a metal)

photon

a particle of light

Photoelectron

An electron released through the photoelectric effect. photon that hits metal and has enough energy to free electron = photoelectron

elctron configuration

shells: energy levels
subshells: s,p,d
orbitals: x,y,z

cation and anion size compared to parent atoms

cations: smaller than parent atoms. outermost electron is removed and repulsions between electrons are reduced
anions: larger. electrons are added and repluisons between electrons are increased.

Beer-Lambert Law

A=ebc. e and b are usually held constant. in such cases, shows that concentration is proportional to absorbance (direct relationship).

Fingerprints on the cuvettes will

cause the light to scatter resulting in less light passing through the sample. This will result in less light being detected and the instrument thinking there is greater absorbance than is really the case.

lower absorbance than actually is

distilled water

physical change

A change in a substance that does not involve a change in the properties but composition. only forces broken are intermolecular
ex: filration, chromatograpy, distlation

chemical change

A change in matter that produces one or more new substances

Titrant

standard solution of known concentration

Analyte

Substance being analyzed. concentration is unknown

titration

when equivalence point is reached when just enough titrant is added to react with all of the analyte. may be indicated by a color change. when moles of acid = moles of base basically

Conjugate Acid and Conjugate Base

Conjugate Acid = Forms when a base gains a proton (opposite of the Bronsted Lowery Base)
Conjugate Base = Forms when an acid loses a proton (opposite of a Bronsted Lowery Acid)
** Conjugate Acids and Bases are found in the products side of the reaction

oxidized means

losing electrons

Reduction means

the gain of electrons by a molecule

factors that govern rates of reaction

-concentration, (rate and concentration are directly related)
-temp (temp and rate are directly related)-catalyst, with a catalyst rate increases
-as reactant surface area increases, rate increases

catalyst

lowers activation energy and is not consumed

reaction rates units

usually Mol/l s

rate =

delta concentration / delta time. will be negative if reactant, postive if product.

how mole ratios relate to rates

for reaction n2 + 3h2, n2 goes down 1/3 as fast as h2.

rate constant k is not affected by

concentration but is affected by temp and catalysts

differential rate law

Describes how rate depends on concentration. r = k (concentration of products). for each type of differential rate law, there is a integrated rate law.

integrated rate law

describes how concentration depends on time. for each type of integrated rate law, there is a differential rate law

zero rate laws summary

rate law: rate =ki
ntegrated rate law: concentration of a = -kt + concentration of a at time = 0
straight line: concentration of a vs time
relationship of rate constant to slope of straight line: slope = -k
half life: t1/2 = concentration of A at time = 0

first order rate laws summary

rate= k(a)
integrated rate law: ln(a) = -kt + ln (a) at t = 0kt= ln ( a0)/(at)
straight line at: ln (a) vs t
relationship of rate constant to slope of straight line: slope = -k
half life: t1/2= .693 /k

second order rate laws summary

rate= k(a)^2
integrated rate law: 1/(a) = kt + 1/a(0 or at t = 0)kt= 1/(a)0 - 1/a(t)
straight line at: 1/a (a) vs relationship of rate constant to slope of straight line: slope = k
half life: t1/2= 1 /(k times a at t = 0)

reaction mechanisms

the process by which a reaction occurs, series of elementary steps. sometimes depends on temp. sum of elementary steps must equal overall balanced equation for reaction

intermediate

formed then consumed shortly after

molecularity

the number of particles/molecules colliding in an elementary step
unimolecular- 1, N2O4(g)?2NO2(g)
bimolecular -2, 2NOCl?2NO(g)+CO2(g)
termolecular - 3 (rare), 2NO(g)+O2(g)?2NO2(g)

subsitute for k (usually)

k2k1 / k-1

catalysts

speed up reaction without being used up. appear as an reactant in elementary step and are produced as a product. can be in rate law.

homogenous catalyst

same phase as reactanting molecules

hetergenous catalyst

are in a different phase as the reactanting molecules. usually solid catalysts in gaseous mixtures or liquid solutions first step

collision theory

1. molecules must collide to react
2. concentration affects rates because collisions are more likely
3. temp and rate are related
4. only a small number of collisions can react

collisions must have

sufficient energy to produce the reaction (must equal or exceed activation energy).

colliding particles

also must be correctly oriented to one another in order to produce an reaction

why are temp and rate directly related

kinetic energy increases with temp.

collision model

explains how rate is affected by concentration and temp. greater concentration, more collisions per sec. higher temperature: faster particles, more collisions per sec, activation energy is exceeded more often

delta e

has no effect on the reaction rate

reaction coordinate diagram

high point of diagram: transition state
species present at the transition state: activated complex
energy gap between reactants and activated complex: activation energy barrier

Maxwell-Boltzmann distribution

at higher temps, a larger population of molecules has higher energy. all have the same integral .

Maxwell-Boltzmann distribution explained

temp = avg kinetic energy. more molecules in lower temp will have low kinetic energy, more molecules with higher temp will have higher kinetic values. this explains height differences.

slowest step

has the highest activation energy

this reaction takes places in three steps, step 1

this reaction takes places in three steps, step 2

this is rate determing step

this reaction takes places in three steps, step 3

where are indemediates on graph

valleys

activated complex or transition state

peak

catalysis

the acceleration of a chemical reaction by a catalyst. not consumed during a reaction.

first step in hetergenous catalysis is

adsorption of reactant molecules onto mental surface, (active sites which are locations at which reactants attach to metal catalyst)

2nd step in hetergenous catalysis is

r bonds are broken or weakened, allowing p to form w a lower Ea, then, the products detach from catalysts

Enzyme

very large protein molecules. biological catalysts. very specific, for one reaction. names end in -"ase

substrates

substances/reactants that react at the active sites of enzymes

Catalysts and Rate

rate increases until the active sites of catalyst are filled. then the rate is independent of concentration

lock and key model

The model of the enzyme that shows the substrate fitting perfectly into the active site

reaction: a + b -----> C. give two ways rate law could be zero for a.

when concentrations of a don't change when initial rate changes. or when a isn't in the slowest elementary step.

k depends on

nature of reactants (state of matter, surface area,etc), temp, and order. order affects units, if catalyst is used.

k doesnt depend on

concentration of reactants

ur given activation energy and delta e for a reaction. what is the activation energy for the reverse reaction.

activation energy - delta e. draw a picture and go backwards.

do graphs start at 0

no, potential energy is never 0

given graph of energy and reaction coordinate, where is the overall activation energy for the reaction

the highest bump/ the highest activation energy. (only strong as ur weakest player)

increase in temp, increase in rate is

not a direct relationship, it is exponential

collision model

the kinetic energy of the molecules before collision is converted to potential energy as the molecules are distorted during a collision to break bonds and rearrange the atoms into the product molecules

why is k negative for 1st and 0 order but not 2nd

how to figure out order

given CONCENTRAION and intial rate, CONCENTRAION doubles, intial rate doubles = 1st, etc. INTIAL RATE DEPENDS ON CONCENTRAION

first order with respect to a and b products

means each reactant is 1st, overall is 2nd

delta h for reverse reaction

opposite of reaction

potential energy

energy due to position or composition. energy in stored bonds.

potential energy and bond strength

stronger or more stable the bond, the less potential energy there is between the bonded atoms. Strong bonds have low potential energy and weak bonds have high potential energy.

1st law of thermodynamics

the total energy of the universe is constant. so energy is reactants = energy in products. same things as law of conversation of energy

potential energy is due to

electrostatic particles between charged particles. this is related to the specific arrangements of atoms in the substance.

exothermic reactions

release energy from system to surroundings. energy is negative

endothermic reactions

absorb energy from surroundings. energy is positive

sign tells u

direction of heat

-q

heat is leaving system to surroundings

q

heat is entering system from surroundings

change in energy of internal system

delta e = e final - e initial

Examples of endothermic reactions

boiling, melting, sublimation (s to g)

examples of exothermic reactions

freezing, condensation, deposition (deposition occurs when molecules settle out of a solution. Deposition can be viewed as a reverse process to dissolution)

energy flow (endo vs. exo)

endo: energy flows into systeme
exo: energy flows out of system

enthalpy

flow of heat in a reaction

if a process occurs at a constant temp

the change in enthalpy of the system = the heat lost of gained by the system

enthalpy is an extensive process, MEANING

the amount of material affects its value

enthalpy of reaction

aka: heat of reaction. (products - reactants). kj/mol

Calorimetry

measurement of heat flow. based on the fact that heat released = heat absorbed. device used is called calorimeter

heat capacity

the number of heat needed to raise the temperature.

molar heat capacity

amount of heat to raise temp one mole of substance 1 k. units: j/mol c or k

specific heat capacity

amount of heat to raise temp one g of substance 1 k. units: j / g c or k

molar heat capacity (formula)

molar mass times specific heat

combustion reaction

produces CO2 and H2O. add o2

Coffe cup calorimeter, temp decreases

endo. vice versa

bond energy

the energy required to break a chemical bond. endothermic process.

bond enthalpy

the energy needed to break one mole of bonds in gaseous molecules under standard conditions. endothermic process

breaking bonds vs forming

break: endo
form: exo

how to find delta h through bond enthalpies

bond enthalpies of bonds broken - bond enthalpies of bonds formed. since required energy to form and then energy released, you can see the change in energy and the direction of which the energy is going.

enthalpy of formation

the enthalpy change for a reaction in which one mole of a compound is formed from its elements in their standard states

standard enthalpy

the enthalpy measured when everything is in its standard state (standard conditions)

standard enthalpy of formation

1 mole of compound is formed from substances in their standard states

why is heat of formation zero for any element in its standard state

because it takes no energy to form a naturally-occurring compound

delta of h =

heat of formation (products) - heat of formation (reactants)

heating curve and heat

right to left: -q
left to right: q
on flat line use q= +/- mcx
on increasing/decreasing line use q= mc delta t

formula for within a given state of matter

q= mc delta t

formula for changing states

q = +/- mcx (cx= heat of vaporization or heat of fusion)

molar heat of fusion

energy required to melt 1 mole of a substance (heat of fusion is per gram)

molar heat of vaporization

The energy required to vaporize one mole of a liquid. (heat of vaporization is per gram)

equations from first law of thermodynamics

q rxn = -q solution

why is liquid to solid exothermic

because the change is one where the matter loses heat

bond enthalpy vs bond energy (to find delta h)

bond energy ( total energy required + energy released)
bond enthalpy= (products - reactants)

how to know when to combine masses within calorimetry problems

reactions or solutions = add masses
piece of metal into water = don't add masses

thermal equilibrium

will eventually = the same kinetic energy

increase in pressure

shift to side with less moles

decrease in pressure

shift to side with more moles

increase in pressure

shift to side with less moles

decrease in pressure

shift to side with more moles

when is k constant

at any temp

reversible reaction

a chemical reaction in which the products re-form the original reactants

chemical equilibrium

In a chemical reaction, the state in which the rate of the forward reaction equals the rate of the reverse reaction, so that the relative concentrations of the reactants and products do not change with time. amts of p and r can be different at equilibrum

equilibrium graph

What is equal at equilibrium?

rates, concentration is not

what is determined by rates

concentrations and activation energy

what doesn't change at equilibrum

concentration

relationship between kc and kp

Kp=Kc(RT)^delta n

if you reverse reaction, k becomes

k1

rules about k

-only depends on reaction stoichometry, not its mechanism
-independent of initial concentrations
-unaffected by other substances, as long as they dont react with p and r
-varies with temp
-written without units
-NEVER includes pure solids or liquids

temp affects

rate

equilibrium concentrations don't

have to be equal, on k. unlimited combos for concentrations.

If K>1

products are favored and equilibrium lies to the right

if K<1

reactants are favored and equilibrium lies to left

k for forward and reverse reaction

are reciprocals

what does product favored mean

when equilibrium is achieved, most reactant has been converted to product

what does reactant favored mean

when equilibrium is achieved, very little reactant has been converted to product

if reaction involves pure solids or pure liquids

the concentration of the liquid doesn't change

reaction quotient

q, what you get when you plug the R and P amts, at any given time into the eq. -constant expression. tells you the direction the reaction will go to reach equilibrium

If Q>K

reaction shifts left.

If Q<K

reaction shifts right. As a system approaches towards equilibrium, Q approaches towards K.. Since Q<K, the reaction will shift to the right to reach equilibrium. visa versa

If Q=K

reaction is at equilibrium

shifts to left means

consumes products and forms reactants until equilibrium is achieved

shifts to right means

consumes reactants and forms products until equilibrium is achieved

can use ICE to solve for

1. equilibrium concentrations
2. k(c/p) at equilibrium

ICE table

Initial, Change, Equilibrium

if k<.01 (ice)

can ignore any -x in equation ( change value in concentrations)

if k>100 (ice)

most of reactant is converted to product, change value becomes insigificant.

5% rule

x should be 5% less than initial concentration

le chatelier

-when you take something away from a system, the system shifts in such a way to replace some what you've taken away.
-when you add something at equilibrium, the system shifts to use up what you've added

if u increase concentration of a reactant or product

-the system is now stressed on that side
-shifts away from the added reactant to relieve the stress

if u decrease concentration of a reactant or product

-the system now has less than it should on that side
-shifts towards the removed reactant to relieve the stress

even though water is a liquid, adding water can cause a shift in equilibrium. why?

it dilutes aqueous components

how to know which way it shifts in dilution

1. One way - plug in numbers to K expression for a normal concentration, then a lower concentration and see what happens
2. it will affect the side with more moles more than the other

change in pressure within a system

-look at number of moles
-if they add a gas, make sure it is in k expression, otherwise it has no effect

(pressure) same number of moles on each side

no effect

increase in pressure

Goes from side with HIGHER # of moles to LOWER # of moles

decrease in pressure

Goes from side with LOWER # of moles to HIGHER # of moles

change in pressure =

change in volume, inversely related (so relationships above are reverse)

p-v changes

does not change k, as long as temp isnt change

change in temp to a system

always results in shifts in eq. and change in k

for exothermic reactions

as t increases, shift left, k decreases
as t decreases, shift right, k increase

for endothermic reactions

as t increases, shift right, k increase
as t decreases, shift left, k decreases

things that have no effect (changes to a system):

- add inert gas
-add/remove a liquid or a solid
-add catalyst (although it speeds it up)

temp changes (more simplifed)

add heat/increase temp= towards
remove heat/decrease temp=away

decrease in concentration can be sometimes done

through making a solid (precipitate) or liquid (water)

common ion effect

a decrease in the solubility of an ionic compound caused by the addition of a common ion (le chatelier)