102 final

solubility

max quantity of a substance that can dissolve, concentration of saturated solution

insoluble = less than 0.01 M at saturation

K_sp - solubility in terms of equilibrium

K “solubility product”

rate equation

Rate = k[A]²[B]

zero order half life

t ½ = [A]₀/2K

first order half life

t ½ = ln2/k

second order half life

1/k[A]₀

reaction order units

effect of reaction manipulation on K

double coefficients? K —> K²

multiply by 1/2? K—> K^1/2

reverse reaction? K —> 1/K

finding K_c from elemantary reactions

Kc = K₁*K₂

effect of an intert gas on K

no effect

inert = nonreactive, usually bc full valence shell (noble gases)

remember, only temp change affects K, not pressure volume or concentration of products or reactants

relation between acid/base strenght and conjugate strength

the stronger the acid, the weaker its conjugate base

P_ka, pH, and concentration of acid and base

when [acid] = [base], pH = p_pka (1/2 V at the equivalence point)

• ideal buffer

when [acid] < [base], pH > p_ka

when [acid] > [base] pH < pka

relationship between intermolecular forces and P_vap

weaker IMF = greater P_vap

change in mass of cathode v anode

anode loses mass

cathode gains mass

Q and K at standard conditions

Q = K

pH at equivalence point for

SA, SB

WA, SB

WB, SA

SA SB - 7

WA SB >7

• why? no more weak acid or strong base, just conjugate base and H2O

WB SA <7

at equilibrium, what is K and ΔG

K =1, ΔG = 0

what does it mean whne a solution is saturated?

it is at equilibrium

to find ksp, jsut plug in values

formation of soluble complex ion can…

increase solubility

Q

solubility concentration quotient, for any state that is not equilibrium

Qsp same as Ksp

when Q > K, should precipitate

Qsp < Ksp unsaturated solution

spontaneous reaction

process that proceeds without external energy input

• water flowing down hill

• salt dissolving in water

• iron rusting

• Zinc reacting with HCl

thermodynamics, system? surroundings?

study of energy transfer between the reaction system and the surroundings

system = chemicals reacting

surroundings = everything else (container, stir bar, bench, air, etc)

1st law of thermodynamics

energy cant be created/destroyed, only transferred

energy transfers as either…

work: action of force through a distance (piston)

• w = -PΔV

• 1 L*atm = 101.325 J

• for all reactions not in gas phase, work is negligible (considering enthalpy of rxn)

Heat. (q): temp difference between system and surroundings - joules

• heat released from syst, -q = exothermic

• heat absorbed by syst, +q = endothermic

E of system is U

• ΔU = q + w

enthalpy

state function fir the thermodynamics of a reaction, at constant pressure

ΔH > 0 = +q = endothermic

ΔH < 0 = -q = exothermic

irreversible processes

spontaneous, unidirectional

• burning paper

• scrambling an egg

when spontaneous process occurs, system and surroundings cannot return to exact original state

• all natural processes are in a sense irreversible — universe irreversibly changed

entropy

quantification of number of microscopic options (microstates) we have that fullfill certain criteria overall (macrostate)

count of how many possible microstates can fullfill a given macrostate

• higher entropy = more options for macrostates

• positive for processes that increase the number of possible microstates

• always increasing - surroundings = the universe

• positive — spontaneous

• extensive property - value depends on the size or amount of matter in a sample

must increase overall

microstate

a specific arrangement of locations and energies of atoms or molecules that make up a system

macrostate

macroscopic, thermodynamic state defined by parameters like pressure, temp, volume

when is it spontaneous (toward products)? ΔG and ΔS_sys

when ΔG is negative

when ΔS_sys positive

relationship between K, Q, G

redox reactions: reduction, oxidation, oxidizing agent, reducing agent

gain elections

lose electrons

what is reduced

what is oxidized

Van’t Hoff Equation: what does it show? what can be translated to?

shows dependence on temperature

assume H and S dont change much with temp

Equilibrium is compromise between minimum H and maximum S

oxidation state

charge that results when electrons of a bond are assigned to the more electronegative atom

charge an atom would possess if the bonding were ionic

• (#valence e on each free atom) - (#valence e in bond (?))

what is electricity

the movement of electrons

galvanic cell v. electrolytic cell

1. reaction produces electricity spontaneously (ΔG_rxn < 0)

2. run electricity through cell to produce a reverse reaction

electrode — cathode v anode

solid, surface of which is where redox reactions occur

cathode: where reduction occurs

• reduction — red colored nob

• where electrons are consumed, positive charge

anode: electrode where oxidation occurs

• a and o vowels

• where electrons enter, negative charge.

electrons move toward cathode to be consumed, current in wire = current via ions in salt bridge

salt bridge

two solutions connected via solution/gel through which ions can flow to maintain current

maintains ion balance

external circuit

metal leads (wires) connect cell to deliver electric current to external circuit'

drawing a cell diagram

single bar = boundary of phases

• solid electrode | aqueous ions

double bar = salt bridge, boundary of half reactions

• aqueous oxidant || aqueous reductant

oxidation on left, reduction on right

voltage

electrical potential difference

driving force of electricity, measured in volts (V)

• drives reactions forward, electromotive force (E_cell) = driving force, related to spontaneity

depends on K,T, and concentrations of products and reactants via Q

Ecell < E^ocell shift toward reactants

electrical energy units, measurement

U

measured in voltage (V) x charge (z)

charge units and measurement

z

measured in coulombs

represents charge transferred in cell reaction as written

Current

I

flow of charge over time

measured in amperes (A) = 1 C/s

Nernst equation

walther nernst

how concentration of reductant and oxidant, size of electrodes, concentration of ions, etc. affect voltage

Faraday’s constant

charge on proton x Avogadro’s number

E^o_cell

cell voltage at standard conditions

Q = 1

total cell potential, spontaneity

E^o_cell = E_red + E_ox

for the same half reaction, E^o_ox = -E^o_red

when E^o_cell > 0, ΔG < 0, therefore is spontaneous

reduction potentials,,,what does it mean

more positive = more powerful oxidizing agent (more easily reduced, readily accepts electrons)

more negative reduction potential = more powerful reducing agent (more easily oxidized)

What is the goal of kinetics?

to deduce the rate of reaction an d how it depends on the concentration of reactants

Instantaneous rate of concentration change

change in concentration at some particular instant of time

slope of tangent line at t

general rate equation

aA + bB —> cC + dD

rate of rxn = -1/a (Δ [A]/Δt) = -1/b (Δ[B]/Δt) = 1/d(Δ[D]/Δt)

rate law

rate of reaction = k[A]₀^x[B}₀^y

second order reaction rate law

rate of run 1/rate of run 3 = (k[A]₁/k[A]₃ )

rate equation for first order reaction + slope

- Δ[A]/Δt = k[A]

ln [A] = [A]₀ - kt

• y= b + mt, m= -k, b = [A]₀

• y-axis: ln [ ]

[A] = [A]_oe^-kt

rate equation + slope for 2nd order reaction

1/[A] = 1/[A]₀ + kt

y = b + mt, m=k

to graph:

y-axis = 1/[A]

x-axis = time

half life

time required for 50% of initial reactant to react

[A] = (1/2)ⁿ[A₀]

• n = number of half lives

nuclide

single type of nucleus, each element can have multiple (isotopes are a type)

band of stability

nuclei w/ 1:1 or higher ratio of neutrons to protons

• 1;1 stable below Z = 20

most stable isotopes, compare to periodic table mass

nuclear decay

spontaneous emission of particles

mass number

number of neutrons + protons

isobar

different elements with the same mass number

isotopes

atoms of an element with different number of neutrons and therefore different masses

ions

atoms can gain or lose electrons to change charge

radioactive decay

directly proportional to the amount of radiactive material (N) present: first order

ln N/N₀ = -(0.693/t_1/2)t

carbon-dating

carbon-14 used ofr radiodating archeological finds less than 30,000 yrs due to its ½ life:

t_1/2 = 5,730 yrs

half life of second order reactions

t_1/2 = 1/k[A]₀

mechanism

sequence of steps called “elementary reactions” that add up to the overall reaction

catalysts

substance that increases the rate of a chemical reaction w/o being consumed

enzymes = biological ones

lowers E_a

in beginning and end of reaction — is not consumed (?) used up then reformed in products of overall reaction

can be included in rate law

bimolecular interaction

elementary reactions w/ 2 reactants

monomolecular reactions

elementary reactions with only 1 reactant

ex/ radioactive decay

3 requirements for a reaction

rate of reaction = (collision frequency)(fraction of collisions w/ the required energy)(fraction of collisions in which molecules have the required relative orientation)

• ONLY collision frequency depends on the concentration of reactants

activation energy

amount of energy required for reactants to react

fraction of collisions w/ enough energy to react increases with temperature

activated complex

state with the least amount of additional energy needed to pass from reactants to products

cannot be isolated

as soon as have activated complex — have product

arrhenius equation

k = Ae^-E_a/RT

ln k₂/k₁ = E_a/R (ΔT/T₁T₂)

rate-determining step

overall rate = rate slowest step

equilibrium

rate forward reaction = rate reverse reaction

homogeneous catalyst

catalyst is in the same phase as the reaction mixture

heterogeneous catalyst

catalyst is in a different phase than the reaction mixture

enzyme

protein molecules that catalyze specific biochemical reactions

substrate

reactant molecule enzyme acts on

active site

region of protein where the reaction takes place

• only a small portion of the whole molecule

rate law of enzyme-catalyzed reaction

rate of reaction = R = Δ[P]/Δt = k[S]/[S] + K_m

[S] = concentration of substrate

[P] = concentration of product

k and K_M are constants

steady-state approximation

ES is consumed as fast as it is formed

k₁[E][S] = (k_-1 + k₂)[ES]

dynamic equilibrium

reactions aren’t being paused/frozen, no net change in [products] and [reactants] because being formed + reacted at same rate

for any reaction at chemical equilibrium…

[product]^x_eq /[reactant]^y_eq = K_c

k_c =. equil. constant in terms of concentrations

decay particles, ordered in terms of mass

alpha - emission of helium nucleus

• minus 4 at the top (mass number), 2 from atomic number

beta - emission of electron from the nucleus

gamma - high energy electromagnetic photon emission by nucleus - photon has no mass

positrons: anti-matter of electrons, particle with negligible mass and positive charge, beta + particles emitted by nucleus

Reaction quotient, Q_c. What does relationship between Q_c and k_c tell you about direction in which reaction should shift?

same expression as Kc, but for any state that is not at equilibrium

Q >K shift to reactants

Q<K shift to products

homogeneous catalyst

catalyst is in the same phase as the mixture

heterogeneous catalyst

catalyst is in a different phase than the reaction mixture

lineweaver-burk plot

When R = 1/2R_max, [S] = K_m

if 1/R is plotted against 1/[S], we get a straight line

Le Chatlier’s Principle

if reaction at equilibrium is subject to change in conditions that displaces it from equilibrium, then the reaction adjusts toward a new equilibrium state. The reaction proceeds in the direction that offsets the change in conditions

at equilibrium, high pressure and low temperature is best

what quantities affect equilibrium?

concentration of reactant or product

reaction volume or applied pressure

temperature — only temperature can change the value of k

if there is the same number of moles on both sides of the reaction, volume wont change equilibrium

affect of temperature on k

increase in temperature (exothermic) —> shift from right to left, k decreases

endothermic- left to right as sustem absorbs heat, increasing concentration of products

consider which direction reaction has to go in order to absorb heat

exo v. endothermic

q < 0 = exothermic

q > 0 = endothermic

heat capacity

amount of energy required to raise temperature of a sample by 1 K

Heat of fusion? Opposite?

q_fus = nΔH_fus

melting (solid to liquid)

freezing

heat of vaporization? opposite?

q_vap = nΔH_vap

liquid to gas

sublimation? opposite?

Energy required to melt 1 mole of a substance

q_vap = nΔH_sub

solid to gas

gas to solid

Variation of heat capacity (h2o)

C liquid > C solid > C gas

Intermolecular forces

ion-ion

• strongest

• not just 1 atom + 1 atom, expanding crystals of repeating units

• solid at room temp

dipole-dipole

• polar molecules attract one another via dipole-dipole

hydrogen bonds (subset of dipole-dipole)

• partial pos on hydrogen and partial neg on N,O,F

• hydrogen so small that partial pos is concentrated —> strogner than other partial positives

• why ice floats

london - dispersion forces

• momentary instantaneous asymmetry for weak dipoles

• everything has

stronger intermolecular forces = higher bp

= higher surface tension

= lower vapor pressure