Chem 2 Final

Front 1

dipole-dipole attractions

Back 1

between polar molecules

Front 2

hydrogen bonding

Back 2

strong dipole-dipole, H + N,O, or F

Front 3

london forces

Back 3

increases as size of atoms increase, longer chain of atoms. small compact molecular shapes, weaker.

Front 4

surface tension

Back 4

molecules at the surface feel fewer attractions than within the liquid, causes interior molecules to have lower potential energy When IM forces are large, surface tension is large

Front 5

wetting

Back 5

the attractive forces between liquid and surface must be about as strong as between molecules of the liquid.

Front 6

viscosity

Back 6

increases with increasing IM attractions and molecular size

Front 7

evaporation/sublimation

Back 7

molecules with large KE escape, leaving avg. KE lower in liquid/solid left behind. substances with weaker IM forces evaporate more quickly

Front 8

vapor pressure

Back 8

independent of container size, surface area, amount of liquid. strong IM forces, low VP. VP increases as temp rises

Front 9

boiling point

Back 9

point at which vapor pressure equals atmospheric pressue

Front 10

endothermic phase changes

Back 10

melting, evaporation, sublimation

Front 11

exothermic phase changes

Back 11

freezing, condensation

Front 12

leChatelier's principle

Back 12

increase in temp shifts eq. in direction of endothermic change. increase in pressure shifts eq. in direction that favors a decrease in volume

Front 13

critical point

Back 13

above this point, no amt. of pressure can create a liquid phase

Front 14

As strengths of IM forces increases, there is an increase in

Back 14

surface tension, viscosity, molar heat of vaporization, molar heat of sublimation, normal boiling point, critical temp

Front 15

As strengths of IM forces increases, there is a decrease in

Back 15

rate of evaporization, vapor pressure

Front 16

What are the four crystalline types?

Back 16

ionic, molecular, covalent, metallic

Front 17

gas solubility as a function of temp and pressure

Back 17

gases become less soluble in water as temp increases, more soluble at higher pressure

Front 18

Henry's law

Back 18

Cgas=kHPgas
Cgas-concen. gas in soln.
kH-proportionality constant
Pgas-partial pressure

Front 19

Molality

Back 19

moles of soluter per kilogram solvent

Front 20

Raoult's law (vapor-pressure concentration law)

Back 20

P(solution)=X(solvent)*P˚(solvent)
P(soln.)-VP of soln.
P˚(solvent)-VP of solvent
x(solvent)-mole fraction

Front 21

solutes on freezing/boiling point

Back 21

Δt=k(b)m
Δt=k(f)m
m-molal conc.
kb-boiling pt. elevation constant.

Front 22

Five factors that control rate of reaction

Back 22

1.chemical nature(bonds, etc.)2.ability of reactants to meet
3.concentrations of reactants
4.temperature
5.catalysts

Front 23

Arrhenius equation

Back 23

k=Ae^(-Ea/RT)
A-proportionality constant
Ea-activation energy
k-rate constant

Front 24

plot of lnk vs 1/T

Back 24

straight line with slope at -Ea/R

Front 25

adsorption

Back 25

when a substance is bound to a surface

Front 26

concentration vs. time (first order reaction)

Back 26

ln([A]o/[A]t)=kt

Front 27

concentration vs. time (second order reaction)

Back 27

(1/[B]t)-(1/[B]o)=kt

Front 28

half-life (first order reaction)

Back 28

t1/2=ln2/k

Front 29

half-life (second order reaction)

Back 29

t1/2=1/k[B]o

Front 30

relating Kc and Kp

Back 30

Kp=Kc(RT)^Δngas

Front 31

stressed equilibrium

Back 31

decrease in volume, shifts eq. toward side with fewer gas molecules. increase in temp. shifts eq. in direction that absorbs heat. catalysts have no effect on eq.

Front 32

binary acid periodic trend

Back 32

the acids become stronger as the location of atom x in the periodic table moves from left to right in a period or from top to bottom in a group

Front 33

oxoacids periodic trend

Back 33

increase in strength from bottom to top, increase from left to right, increase with number of oxygen atoms

Front 34

Lewis acid/base

Back 34

acid-accepts electron pair
base-donates electron pair

Front 35

acidic cations

Back 35

cations of groups 1A and 2A except Be do not have high enough charge densities to be acidic

Front 36

acid ionization constant

Back 36

Ka=[H+][A-]/[HA]
pKa=-logKa

Front 37

Henderson-Hasselbalch

Back 37

pH=pKa-log([HA]init./[A-]init.)

Front 38

common ion effect

Back 38

a salt is less soluble in a solution that contains one of its ions rather than water

Front 39

first law of thermodynamics

Back 39

internal energy may be transferred as heat or work but it cannot be created or destroyed ΔE=q+w

Front 40

q

Back 40

heat
q>0, heat is absorbed
q<0, heat is released

Front 41

w

Back 41

w>0, work is done on a system
w<0, work is done by the system.

Front 42

pressure-volume work

Back 42

w=-PΔV
P-external pressure on system
ΔE=q-PΔV

Front 43

enthalpy

Back 43

the heat constant of a system.
ΔH=ΔE+PΔV.
ΔH>0, endothermic
ΔH<0, exothermic

Front 44

ideal gas law

Back 44

V=nRT/P
Δngas=(ngas)products-(ngas)reactants
ΔH=ΔE+ΔngasRT

Front 45

entropy(s)

Back 45

# of equivalentways the E of a system can be distributed.
any event that has an increase in s of system will have a tendency to occur spontaneously. Endothermic reactions increase s.

Front 46

volume on entropy

Back 46

for gases, entropy increases with increasing volume

Front 47

temp. on entropy

Back 47

the higher the temp., the larger is the entropy

Front 48

physical state and entropy

Back 48

entropy increasing moving from solid to liquid to gas.

Front 49

predicting spontaneity from standard cell potentials

Back 49

in galvanic cell, the calculated cell potential for the spontaneous reaction is always positive. If negative, it is spontaneous in the reverse direction

Front 50

free energy changes from cell potentials

Back 50

ΔG=-nFEcell
F=coulombs/mol e-
F=96,500C/1mol e-

Front 51

equilibrium constant from cell potential

Back 51

Ecell=RT/nFlnKc
R=8.314
F=9.65e4C per mol e-
n=# mols e- transferred

Front 52

nernst equation

Back 52

Ecell=Ecell-RT/nFlnQ

Front 53

standard reduction potentials table

Back 53

oxidizing agents to left of double arrows. best oxidizing agents are at top, also have greatest tendency to occur as reductions.

Front 54

predicting redox reactions by reduction potentials

Back 54

the half reaction with the more positive reduction potential always takes place as written (as a reduction), while the other runs in reverse (oxidation)

Front 55

reduction potential

Back 55

measure of the tendency of a given half reaction to occur as a reduction. when two half cells are connected, the one with the larger reduction potential aquires e- from the half cell with the lower reduction potential

Front 56

standard reduction potential

Back 56

Ecell=(standard reduction potential of substance reduced)-(standard reduction potential of substance oxidized)

Front 57

a negative reduction potential means

Back 57

that the substance is not as easily reduced as H+

Front 58

predicting the sign of entropy change

Back 58

gaseous products larger than gaseous reactants, positive entropy change. less gases on reactant side, negative entropy change. reactions that increase number of particles in the system tend to have a postive entropy change

Front 59

three factors that influence spontaneity

Back 59

enthalpy change, entropy change, temp

Front 60

second law of thermodynamics

Back 60

whenever a spontaneous event takes place in our universe, the total entropy of the universe increase.
ΔHsystem-TΔSsystem<0

Front 61

free energy (G)

Back 61

ΔG=ΔH-tΔS at constant T and P.
can only be spontaneous if it is accompanied by a decrease in the free energy of the system.

Front 62

ΔH<0
ΔS>0

Back 62

the process will be spontaneous at any temp

Front 63

ΔH>0
ΔS<0

Back 63

the process is not spontaneous at any temp

Front 64

ΔH>0
ΔS>0

Back 64

spontaneous at high temps, endothermic reaction

Front 65

ΔH<0
ΔS<0

Back 65

spontaneous at low temps, exothermic reaction

Front 66

third law of thermodynamics

Back 66

at absolute zero the entropy of a perfectly ordered pure crystalline substance in zero

Front 67

standard entropy change

Back 67

ΔS=(sum of S of products)-(sum of S reactants)

Front 68

standard entropy of formation

Back 68

the value of ΔS for the formation of one mole of a substance from its elements in their standard states

Front 69

standard free energy change (ΔG)

Back 69

ΔG=ΔH-(298.15K)ΔS
ΔG=(sum of ΔGfproducts)-(sum of ΔGf reactnts)

Front 70

thermodynamically reversible process

Back 70

a process that occurs by an infinite number of steps during which the driving force for the change is just barely bigger than the force that resists the change

Front 71

Is there a limit to the amount of the available energy in a reaction that can be harnessed as useful work?

Back 71

the maximum amount of energy produced by a reaction that can be theoretically harnessed as work is equal to ΔG

Front 72

What is ΔG when a system is at equilibrium?

Back 72

ΔG=0
Gproducts=Greactants
no work can be done by a system at equilibrium.
ΔH=TΔS

Front 73

relating the reaction quotient to ΔG

Back 73

ΔG=ΔG˚+RTlnQ
lnQ=ln of reaction quotient, partial pressures or molar concentration.

Front 74

thermodynamic equilibrium constants

Back 74

ΔG˚=-RTlnK
at equil. Q=K
Kp=gases
Kc=soln.

Front 75

bond energy

Back 75

the amount of energy needed to break a chemical bond to give electrically neutral fragments

Front 76

atomization energy (ΔHatom)

Back 76

the amount of energy needed to rupture all the chemical bonds in 1 mol of gaseous molecules to give gaseous atoms as products

Front 77

the formation of the CH3OH molecule from gaseous atoms...

Back 77

is exothermic because energy is always released when atoms become joined by covalent bonds

Front 78

galvanic cell

Back 78

an electrochemical cell in which a spontaneous redox reaction produces electricity

Front 79

cathode/anode (galvanic cell)

Back 79

cathode: reduction (electron gain) occurs
anode: oxidation (electron loss) occurs
cations move toward the cathode, anions toward the anode

Front 80

metallic conduction

Back 80

conduction of electrical charge by the movement of electrons

Front 81

volt

Back 81

unit of electromotive force on emf in joules per coulomb

Front 82

electrolytic conduction

Back 82

the transport of electrical charge by ions

Front 83

cell potential

Back 83

the emf of a galvanic cell when no current is drawn from the cell

Front 84

The type of solid generally characterized by low melting point and extremely low conductivity is:

Back 84

molecular

Front 85

When a nonvolatile solute is added to a volatile solvent, what happens to VP, BP and FP?

Back 85

vapor pressure decreases, boiling point increases and freezing point decreases

Front 86

A solid that has a high melting point, great hardness, poor electrical conduction, and the particles at the lattic points are connected by covalent bonds

Back 86

covalent network solid

Front 87

when the phase diagram for a substance has a solid-liquid phase boundary line that has a negative slope, the substance

Back 87

can go from solid to liquid within a small temp range via the application of pressure

Front 88

a catalyst acts by

Back 88

diminishing the activation energy of a reaction

Front 89

for a system whose equilibrium constant is relatively small...

Back 89

equilibrium lies to the left

Front 90

what are the units on the rate constant for the following rate law: rate=k[D][x]^2[E]^0

Back 90

L^2mol^-2s^-1

Front 91

HA is a weak acid. What is the equilibrium that corresponds to Kb for A-?

Back 91

A-(aq)+H20(l)->HA(aq)+OH-(aq)

Front 92

Ka on acid strength

Back 92

the smaller Ka, the weaker the acid

Front 93

bronsted/lowry acid/base

Back 93

acid-proton donor
base-proton acceptor

Front 94

endothermic reaction

Back 94

heat + products -> reactants
heat absorbing reaction.
products have more potential energy on reaction coordinate.