AP Chemistry Semester 1 Final

Laws of Thermodynamics

1. The energy contained within the universe is constant 2. The entropy of the universe is constantly increasing

oxoacid

1. an OH group bonded to an element that is not bound to other Oxygens (HOY acids) 2. an OH gropu bonded to an elemtn that IS bound to other Oxygens (HYOn) acids

rules for rates of effusion

1. gas particles that move faster will effuse at a faster rate 2. gas particles with smaller molar masses have higher avg. velocities

how to tell if a reaction is spontaneous

1. if ΔS(universe) > 0 2. Gibbs Free energy (ΔG < 0)

factors affecting solubility

1. structure (polar dissolves polar, non-polar dissolves non-polar) 2. temperature (rules vary, typically solubility increases as temperature increases) 3. Pressure (as it applies to gas-liquid solutions)

requirements for a reaction to take place

1. there must be a collision 2. the collision must occur with he correct orientation 3. the energy of the collision must be greater than or equal to the activation energy

conditions in which gases do NOT behave ideally

1. under high pressures (P > 5 atm) 2. at low temperatures (because gas particles now occupy a significant percent of the volume)

reaction intermediate

a substance that is produced and then consumed during the overall reaction

buffered solutions

a weak acid and its salt / a weak base and its salt buffers resist changes in pH

if ΔH is positive, ΔS is negative...

non-spontaneous at all temperatures

van't Hoff factor

number of ions within 1 formula unit of a strong electrolyte

n

number of moles*

atomic number

number of protons

mass number

number of protons + number of neutrons

if ΔH is positive, ΔS is positive...

spontaneous at high temperatures, non-spontaneous at low temperatures

if if ΔH is negative, ΔS is negative...

spontaneous at low temperatures, non-spontaneous at high temperatures

ΔH°(f)

standard enthalpy of formation ΔH°(rxn) = ∑nΔH°(f)(products) - ∑nΔH°(f)(reactants)*

STP

standard temperature and pressure; 0°C and 1atm

comparison of acids to conjugate bases and bases to conjugate acids

strong acids have weak conjugate bases, weak acids have strong conjugate bases; strong bases have weak conjugate acids, weak bases have strong conjugate acids

weak acid - strong base reactions

strong bases completely dissociate OH- rips H+ ion off of the weak acid this produces water and the conjugate base of the weak acid

conjugate acids / bases

the acid/base left over after a base/acid accepts/donates a proton

the common ion effect

the addition of a salt that dissociates completely, leaving a common ion to shift equilibrium (see "Acids and Bases IV" page 1)

activation energy

the amount of energy required for a reaction to take place

autoionization of water

H2O(l) + H2O(l) --> H3O+ + OH- K(w) = [H30+][OH-] = 1.0 x 10^-14* (see "Acids and Bases I" page 5)

the "Big Six" strong acids

HClO4, HI, HBr, HCl, HNO3, H2SO4

finding Δq(rxn) from Δq

the heat gained by the water = the heat lost in the reaction

solubility

the maximum molar concentration of formula units that will dissolve at a given temperature *a solution is at equilibrium when it is saturated

carboxylic acid strengths

the more electronegative R (any element) is, the stronger the acid

oxoacid strengths

HOY: acid strength increases as the electronegativity of Y increases (Y = an element) HYO: acid strength increases as more Oxygens are added to the Y element

entropy

S a measure of disorder of a system >S is favorable ΔS = S(products) - S(reactants)

percent dissociation

the percent of an acid that dissociates with H+ ions (multiply percent decimal by the concentration to find how much was dissociated)

limiting reactant

the reactant that is completely used up in a reaction; limits how far the reaction will proceed

excess reactant

the reactant that is not completely used up in a reaction

rate of the overall reaction is always equal to ____

the slowest elementary step

solubility and strong acids

the solubility of slightly soluble salts containing basic anions increases as [H+] increases equilibrium shifts to the right

solubility and weak acids

the solubility of slightly soluble salts containing basic anions increases as [H+] increases equilibrium shifts to the right

saturated solution

the solvent has dissolved the maximum amount of solute that it can at a certain temperature (some solid solute remains at the bottom)

K(a) vs. acid strength

the stronger the acid, the larger the K(a)

K(b) vs. base strength

the stronger the base, the larger the K(b)

half life

the time it takes for the concentration of a reactant to decrease by half

calorimetry

the use of a calorimeter to measure heat change, using: Δq = mcΔT* Δq - heat lost/gained by water m - mass of water c - specific heat capacity of water ΔT - temperature change of water

root mean square velocity

u(rms) = √(3RT/MM)* MM - molar mass (kg/mol)~

atomic mass unit (amu)

unit of measurement for weighing atoms (12 amu = Carbon-12 exactly)

finding pH after a titration

use moles of each, and subtract (if weak acids/bases are involved, use an ICE chart) (see "Acids and Bases V" pages 1-5)

carboxylic acids

weak organic acids (RCOOH / RCO2H, where R can be almost anything)

manipulating K(eq); reciprocal rule

when a reaction is reversed, the new K(eq) value is the inverse of the old K(eq) value

Le Chatelier's Principle

when a system at equilibrium is subjected to a stress (pressure, concentrations, temperature), the equilibrium will shift

factors that make ΔS positive

when changing from solid to liquid, liquid to gas when products have more particles than reactants when adding heat

non-ideal solutions

when intermolecular forces between solute and solvent molecules are very strong

ideal solutions

when the solute and the solvent have similar structures

manipulating K(eq); multiple equilibria rule

when two or more reactions are combined, the new K(eq) is the product of the K(eq) values from each reaction

relating ΔG° to Q

ΔG = ΔG° + RT lnQ* Q = e^(ΔG-ΔG°/RT)~ (must be in J, not kJ)~

finding ΔG (Gibbs Free Energy)

ΔG = ΔH - TΔS* (if ΔG <0, reaction is spontaneous) (must be in kJ, not J)

free energy change

ΔG°(f) from when one mole of a compound is made form its elements in their standard states.

ΔG°(rxn)

ΔG°(rxn) = ∑nΔG°(f)(products) - ∑nΔG°(f)(reactants)*

enthalpy of a reaction

ΔH = Hproducts - Hreactants

finding ΔH(rxn) from Δq(rxn)

ΔH = Δq/n (this finds the heat lost/gained per 1 mol)

average bond enthalpies

ΔH = ∑BE(bonds broken) - ∑BE(bonds formed)

finding ΔS(universe)

ΔS(universe) = ΔS(system) + ΔS(surroundings)

calculating ΔS

ΔS°(rxn) = ∑nS°(products) - ∑nS°(reactants)*

boiling point elevation

ΔT(b) = imK(b)* ΔT - change in temperature i - van't Hoff factor (number of ions within 1 formula unit of a strong electrolyte) m - molality K(b) - boiling point constant for the solvent

freezing point depression

ΔT(f) = imK(f)* ΔT - change in temperature i - van't Hoff factor (number of ions within 1 formula unit of a strong electrolyte) m - molality K(f) - freezing point constant for the solvent

atm vs. mm Hg (torr)

1 atm = 760 mm Hg (torr)*

Boyle's law

V1P1 = V2P2 (volume is inversely proportional to pressure)

Van der Waals equation

(P + n^2a/V^2)(V-nb) = nRT* a - constant b - constant

finding pH with the [H+] or [H3O+]

-log[H+] = pH* -log [H3O+] = pH

ΔH°(f) and ΔG°(f) for all elements

0

relating ΔG° to K(eq)

0 = ΔG° + RT lnK(eq) ΔG° = -RT lnK(eq) K(eq) = e^-ΔG°/RT

R constant

0.0821 L atm / mol K* 8.314 J / mol K*

2nd order integrated rate law

1/[A]t = kt + 1/[A]o* (written differently on formula sheet) [A]o = initial concentration (at t = 0) [A]t = concentration after some period of time

strong acid - strong base reactions

100% dissociation H+ and OH- ions form H2O the other parts act as spectator ions

finding [H+] or [H3O+] using pH

10^-pH = [H+] 10^-pH = [H3O+]

Avogadro's number

6.022 x 10^23* (number of particles in 1 mol of a substance)

acid dissociation constant

K(a) functions as other equilibrium constants

relationship between K(a) and K(b)

K(a) x K(b) = 1 x 10^-14*

base dissociation constant

K(b) functions as other equilibrium constants

equilibrium expression for conentrations

K(c)* K(c) = [products]^(coefficients) / [reactants]^(coefficients)

equilibrium constant

K(eq)

equilibrium expressoin for gases

K(p) K(p) = P(products)^(coefficients) / P(reactants)^(coefficients)

conversion between K(c) and K(p)

K(p) = K(c)(RT)^Δn* Δn= moles of gaseous products - moles of gaseous reactants

solubility product

K(sp) the equilibrium constant for ionic compounds in water

kinetic energy per gas particle

KE = 1/2 x mv^2* m - mass (kg)~

Raoult's Law: volatile solute

P(solvent) = X(solvent)P°(solvent) P(solute) = X(solute)P°(solute) P(solvent/solute) - partial pressure of solvent/solute X(solvent/solute) - mole fraction of the solvent/solute P°(solvent/solute) - vapor pressure of the solvent/solute

Raoult's Law: non-volatile solute

P(solvent) = X(solvent)P°(solvent) P(solvent) - partial pressure of solvent X(solvent) - mole fraction of the solvent P°(solvent) - vapor pressure of the solvent

combined gas law

P1V1/T1 = P2V2/T2*

ideal gas equation

PV=nRT*

reaction quotient (Q)

Q is compared to K(eq) to determine if a reaction is to the left/right of equilibrium (or at equilibrium) if Q > K(eq), reaction will proceed to the left (too many products) if Q < K(eq), reaction will proceed to the right (too many reactants) if Q = K(eq), the system is at equilibrium

1st order integrated rate law

Rate = ln([A]t) - ln([A]o) = -kt* [A]o = initial concentration (at t = 0) [A]t = concentration after some period of time

rate of disappearance of a single reactant

Rate(A) = —Δ[A] / Δt

rate of appearance of a single product

Rate(D) = Δ[D] / Δt

Charles' law

V1/T1 = V2/T2 (volume is proportional to temperature)

volume adjustment for gases under high pressure

V(ideal) = V(measured) - nb* n - moles of gas b - constant for the type of gas

mole fractions

Xa = (moles of one component in a mixture) / (sum of the moles of all components in a mixutre)*

mass percent (regarding atoms)

[ (# of atoms of an element) x (element's atomic mass) / (formula weight of compound) ] x 100

percent yield

[ (actual yield) / (theoretical yield) ] x 100

hetergeneous catalyst

a catalyst in a different phase as the reactants

homogeneous catalyst

a catalyst in the same phase as the reactants

spontaneous process

a process that proceeds without any assistance from outside the system; often exothermic ΔG < 0

acid-base titrations

acid and base mixed with an indicator to determine the concentration of acid/base

strong acid - weak base reactions

acid completely dissociates weak bases will accept protons from a strong acid weak bases may be nitrogen containing compounds such as NH3 or the conjugate bases of weak acids

calculating pH for a strong acid solution

acid dissociates completely; -log[H+] (see "Acids and Bases II" page 3)

calculating pH for a weak acid solution

acid does NOT dissociate completely; use ICE chart to find [H+] (see "Acids and Bases II" page 4)

binary acid strengths

acid strength increases when moving down a group (ex. HF < HCl)

Lewis acids and bases

acids: accept electron pairs (must have an empty orbital) bases: donate electron pairs (must have at least one lone pair)

Arrhenius acids and bases

acids: dissociate in water to produce H+ ions bases: dissociate in water to produce OH- ions

Brønsted-Lowry acids and bases

acids: donate protons bases: accept protons

gas phase acid-base reactions

all atoms are present in the net-ionic equation try to put together recognizable anions and cations; if not, just stick them together

if K(eq) = 1 (or close)

almost equal amounts of each, equilibrium lies in the middle

binary acid

an acid consisting of a single element with a single hydrogen atom

acid anhydrate reactions

anhydrates react with water or OH- to form an acid; oxidation numbers do not change

anions' effect on solutions

anions can make solutions basic

ΔG at equilibrium

at equilibrium, ΔG = 0

isotope

atoms of the same element with different numbers of neutrons; identical chemical behavior

average kinetic energy per mole

avg. KE per mol = 3/2 x RT*

anhydrates

binary compounds composed of non-metals; usually oxides

common characteristics of bases

bitter taste; feels slippery; found in cleaning products, blood

polyprotic acids

can donate more than one H+ in a solution (poly = multiple / protic = proton)

cations' effect on solutions

cations can make acidify solutions

bromthymol blue

changes color at a pH of 6.0-8.0 (common indicator)

phenolphthalein

changes color at a pH of 8.2-10.0 (common indicator

combustion analysis

combustion: yields CO2 + H2O; see "Stoichiometry" page 5

strong acid properties

completely dissociate in water (equilibrium lies far to the right)

volatile

easily evaporates

elastic collisions

energy is conserved in each collision

ΔH°

enthalpy at standard state (25°C, 1 atm)

decreasing concentration to one side of a reaction at equilibrium

equilibrium will shift towards that side

increasing concentration to one side of a reaction at equilibrium

equilibrium will shift towards the other side

increasing pressure on a system at equilibrium

equilibrium will shift towards the side with less moles

decreasing pressure on a system at equilibrium

equilibrium will shift towards the side with more moles

properties of gases

expand to fill the volume of their container, form homogeneous mixtures, low density, highly compressible, exert a pressure

osmotic pressure

experiment with a U-shaped tube with a semi-permeable membrane in the middle (only the solvent side can pass through); osmotic pressure makes the level of liquid on each side uneven (see "Solutions III" page 5)

how to choose a titration indicator

find the pH at the equivalence point (not always 7)

force attraction in real gasses

force is proportional to 1/d^6 d - distance between particles

empirical formula

formula for a molecule in the simplest ratio (ex. C8H12 would be C2H3)

weak acid - weak base reactions

generally do not go to completion (equil. arrows) acids and bases are mostly undissociated written as a proton transfer reaction

predicting possible precipitates

given a theoretical mixture, use find Q and compare it to K(sp) to predict if a precipitate will form (see "Solutions II" pages 2-3)

endothermic process

heat is absorbed; positive change in enthalpy

exothermic process

heat is released; negative change in enthalpy

comparing Q to K(sp)

if Q = K(sp), system is at equilibrium and is saturated if Q > K(sp), a precipitate will form; the reaction will proceed to the left if Q < K(sp), no precipitate forms; solution is unsaturated

Hess's Law

if a reaction is carried out in a series of steps, the overall change in enthalpy will be equal to the sum of the enthalpy changes for the individual steps. (modify the equation steps given so that you can cancel them out, while changing ΔH as well, then add them up) (see "Thermodynamics I" page 4)

manipulating K(eq); coefficient rule

if coefficients are changed by a factor of n, K(eq) is raised to the power of n

increasing temperature of a reaction at equilibrium

if exothermic: equilibrium will shift to the left if endothermic: equilibrium will shift to the right

decreasing temperature of a reaction at equilibrium

if exothermic: equilibrium will shift to the right if endothermic: equilibrium will shift to the left

ΔG's relationship with K(eq)

if ΔG° > 0, K(eq) < 1 if ΔG° > +20kJ, K(eq) <<1 if ΔG° < 0, K(eq) > 1 if ΔG° < -20kJ, K(eq) >>1

cation

ion with less electrons than usual (positive charge) (always written first in an ionic compound)

anion

ion with more electrons than usual (negative charge) (always written last in an ionic compound)

hyrdates

ionic compounds that trap water within their structures (ex. NaCl · H2O

salts with weak acids and weak bases

little to no effect on pH (remember this for quetion 4)

arrhenius equation

ln(k) = (-E(a)/R) x (1/T) + ln(A)* k - rate constant E(a) - activation energy A - constant

catalyst

lowers activation energy; not produced nor consumed in a reaction

overall order for a reaction

m + n (add the reaction orders)

mass percent (regarding solutions)

mass of component / total mass of solution

metal oxides

metal oxides are basic (remember this for question 4)

molality (m)

moles solute / kg solvent*

molarity (M)

moles solute / liters solution*

if K(eq) is > 10

mostly products, equilibrium lies far to the right

if K(eq) is < 0.1

mostly reactants, equilibrium lies far to the left

effusion

movement of gas particles through a very small hole (giggity) into a vacuum

diffusion

movement of one type of gas into another type of gas

finding averge atomic mass

multiply each isotope by the percent abundance, then add together (see "Atomic Theory I", page 2)

supersaturated solutions

occurs when the solvent has dissolved more solute than it should be able to at a certain temperature (done by carefully lowering the temperature of a saturated solution)

pOH

opposite of pH pH + pOH = 14* pOH = -log[OH-]* [OH-] = 10^-pOH

Henderson-Hasselbalch equation

pH = pK(a) + log([A-] / [HA])* pOH = pK(b) + log([HB+] / [B])* pK(a) - -log(K(a)) pK(b) - -log(K(b)) [A-] - molarity of conjugate base [HA] - molarity of weak acid [B] - molarity of weak base [HB+] - molarity of conjugate acid

buffered solutions and pH

pH = pK(a) when [A-] = [HA], as log(1) = 0

finding partial pressure using mole fraction

partial pressure = total pressure x mole fraction*

weak acid properties

partially dissociate (equilibrium lies to the left or near the middle)

pressure equation

pressure = force/area force - N area - m^2 pressure - Pa

Dalton's law of partial pressures

pressure exerted is equal to the sum of the partial pressure of each gas (P(total) = P(1) + P(2) + P(3) ... + P(n-1) + P(n))*

fractional crystallization

problems involving dissolving multiple salts, then cooling the solution in order to produce as much of one salt as possible (see "Solutions I" page 5)

rate law

rate = k[A]^m x [B]^n k - constant (different for each reaction) m - reaction order in terms of A n - reaction order in terms of B

rate of effusion

rate1/rate2 = √(MM2)/(MM1)*

if ΔG is zero...

reaction is at equilibrium

if ΔG is positive...

reaction is non-spontaneous, and can't happen (you can't go uphill on the graph)

if ΔG is negative...

reaction is spontaneous, and is moving towards equilibrium

finding half life for 1st order reactions

set [A]o equal to 1 M, and set [A]t equal to .5 M (see "Kinetics I" page 6)

items left out of the equilibrium expression

solids and liquids~

gas solubility and temperature

solubility of gases (mostly) decreases as temperature increases

miscible

soluble in all proportions

common characteristics of acids

sour taste; found in fruits (citrus)

if ΔH is negative, ΔS is positive...

spontaneous at all temperatures