Chem I: 11-12 Flashcards
(245 cards)
1
Q
When you add a solution of NaCl to a solution of AgNO3, why is it that the precipitate is AgCl but not NaNO3 in terms of Ksp?
A
Because AgCl has a lower Ksp value, meaning it is more likely to form a precipitate at a lower concentration.
2
Q
Which of the following is NOT a salt solubility rule in water?
(A) All group 1 and ammonium salts are soluble
(B) All nitrate, perchlorate and acetate salts are soluble
(C) All carbonate and phosphate salts are soluble
(D) All silver, lead and mercury salts are insoluble, except for their nitrates, perchlorates and acetates
A
(C) All carbonate and phosphate salts are soluble
Carbonate and phosphate salts are typically INSOLUBLE, unless they are bound to group 1 or ammonium salts.
3
Q
a) all nitrate, perchlorate, and acetate salts are soluble
b) all carbonate and phosphate salts are soluble
c) all silver, lead and mercury salts are insoluble, except for their nitrates, perchlorates, and acetates
d) all group 1 and ammonium salts are soluble
A
b) all carbonate and phosphate salts are soluble
4
Q
solution
A
homogenous mixtures of 2+ substances that combine to form a single phase
5
Q
relationship between mixtures and solutions
A
all solutions are considered mixtures, but not all mixtures are considered solutions
6
Q
solute
A
dissolved in a solvent
ex: NaCl, NH3, CO2, glucose
7
Q
solvent
A
component of solution that remains in same phase after mixing
if the substances are already in same phase, the solvent is the component present in greater quantity
8
Q
solvation
A
aka dissolution
- electrostatic interaction between solute and solvent molecules
- breaking intermolecular interactions between solute and solvent molecules and forming new intermolecular interactions between them
9
Q
if solvation is exothermic…
process is favored at ___ temperatures
A
new interactions are stronger than the original ones
low temp
10
Q
if solvation is endothermic…
process is favored at ___ temperatures
A
new interactions are weaker than the original ones
high temp –> since new interactions weaker, energy needed to facilitate their formation
11
Q
ideal solution
A
when enthalpy of dissolution is 0
12
Q
spontaneous formation of solutions
exothermic vs endothermic
A
both can form spontaneously
13
Q
at constant temp and pressure, entropy always ______ upon dissolution
A
increases
14
Q
solubility
A
max amount of that substance that can be dissolved in a particular solvent at a given temp
15
Q
saturated
A
- when max amount of solute has been added –>> dissolved solute is in equilibrium with its undissolved state
- if more solute is added, it will not dissolve
- rates of dissolution and precipitation are equal
16
Q
dilute
A
solution in which the proportion of solute to solvent is small
17
Q
hydration
A
solvation in water
water molecules break ionic bonds
ions surrounded and stabilized by shell of solvent molecules
18
Q
hydration rxn
NaCl (s) –>
A
Na+(aq) + Cl-(aq)
19
Q
precipitation rxn
A
ions come together to form a solid that falls out of solution
20
Q
precipitation rxn
NaCl(aq) + AgNO3(aq) –>
A
AgCl(s)
21
Q
concentrated
A
solution in which the proportion of solute to solvent is small
22
Q
sparingly soluble salts
A
solutes that dissolve minimally in the solvent
23
Q
aqueous solution
A
solvent is water
24
Q
H+ is never found alone in solution bc….
A
a free proton is difficult to isolate
25
most important solubility rules
all salts of Group I metals and all nitrate salts are soluble
26
7 general solubility rules
1. salts containing ammonium (NH4+) and group I cations are water soluble
2. salts containing nitrate (NO3-) and acetate (CH3COO-) anions are water soluble
3. halides (excluding fluorides) are water soluble
1. exceptions: Ag+, Pb2+
4. salts containing sulfate (SO42-) are water soluble
1. exceptions: Ca2+, Sr2+, Ba2+, Pb2+
5. all metal oxides insoluble
1. exceptions: alkali metal, ammonium, CaO, SrO, BaO
6. all hydroxides insoluble
1. exceptions: alkali metal, ammonium, Ca2+, Sr2+, Ba2+
7. all carbonates, phosphates, sulfides, and sulfites are insoluble
1. exception: alkali metals and ammonium
27
complex ion
aka coordination compound
moleculae in which a cation is bonded to at lease one electron pair donor
28
ligands
electron pair donor molecule in complex ion
29
coordinate covalent bond
hold complex ions together
electron pair donor and electron pair acceptor from very stable lewis acid-base adducts
30
complex formation biological applications
* active sites of proteins
* iron cation in hemoglobin
* coenzymes (vitamins) and cofactors
31
chelation
central cation bonded to same ligand in multiple places
requires large organic ligands that can double back to form a second or third bond with the central cation
32
concentration
amount of solute dissolved in a solvent
33
percent composition by mass eq
mass of solute/mass of solution x 100%
34
mole fraction
eq
XA = moles of A/total moles of all species
35
sum of mole fractions in a system will always =
1
36
molarity eq
M = moles of solute/liters of solution
37
molality eq
m = moles of solute/kg of solvent
true only for dilute aqueous solutions
38
eq used to determin conc of a solution after dilution
MiVi=MfVf
39
saturation point
solution concentration is at its max value for the given temp and pressure
40
when solution is dilute, the thermodynamically favored process is \_\_\_\_\_\_.
initially, the rate of \_\_\_\_\_\_.....
dissolution
initially, the rate of dissolution will be greater than the rate of precipitation.
41
as solution becomes more concentrated and approaches saturation, the rate of dissolution \_\_\_\_\_\_, while the rate of precipitation \_\_\_\_\_.
lessens
increases
42
degree of solubility determined by:
relative changes in enthalpy and entropy associated with the dissolution of the ionic solute at a given temp and pressure
43
solubility product constant eq
Ksp = [An+]m[Bm-]n
44
Ksp dependent on
temperature
| (and pressure for gases)
45
As temp increases, Ksp
increases for non gas solutes and decreases for gas solutes
46
ion product eq
IIP = [An+]m[Bm-]n
47
difference between ion product and Ksp
concentrations used in IP are concentrations of the ionic constituents at that given moment in time, which may differ from Ksp
48
IP \< Ksp
unsaturated -\> soln not yet at equilibirum
solute will continue to dissolve
49
IP = Ksp
saturated
solution is at equilibrium
50
IP \> Ksp
supersaturated -\> solution is beyond equilibrium
precipitation will occur
51
supersaturated soln
* beyond equilibrium
* thermodynamically unstable
* any disturbance to soln will cause spontaneous precipitation of the excess dissolved solute
52
molar solubility
molarity of a solute in a saturated soln
53
54
55
every sparingly soluble salt of general formula MX will have Ksp=
Ksp = x2
x: molar solubility (assuming no common ion effect)
56
every sparingly soluble salt of general formula MX2 will have Ksp=
Ksp = 4x3
x: molar solubility (assuming no common ion effect)
57
every sparingly soluble salt of general formula MX3 will have Ksp=
Ksp = 27x4
x: molar solubility (assuming no common ion effect)
58
formation of complex ions _____ the solubility of salt in a soln
increases
59
formation or stability constant
Kf
for ocmplex ions
60
61
common ion effect
presence of common ion results in a reduction in molar solubility of the salte
has no effect on the value of the solubility product constant itself
62
colligative properties
physical properties of solns that are dependent on the conc of dissolved particles
NOT on the chemical identity of the dissolved particles
include: vapor pressure depression, BP elevations, freezing point depression, osmotic pressure
63
vapor pressure depression
* raoult's law
* as solute is added to a solvent, vapor pressure of solvent decrease proportionately
* presence of solute molecules can block the evaporation of solvent molecules, but not their condensation --\> reduces vapor pressure
64
as conc of B increases, vapor pressure of A \_\_\_\_\_\_
decreases
as more solute is dissolved into solvent (as more B is dissolved into A), the vapor pressure of the solvent decreases
65
raoult's law
eq
PA = XA P°A
PA: vapor pressure of A when solutes present
XA: mole fraction of A
P°A: vapor pressure of A in pure state
66
raoult's law hold only when...
the attraction between the molecules of the different components of the mixture is equal to the attraction between the molecules of any one component in its pure state
67
as vapor pressure decreases, boiling point \_\_\_\_\_\_
why?
increases
higher temp is required to match the atmospheric pressure
68
ideal solutions
solutions that obey raoult's law
69
70
when a nonvolatile solute is dissolved into a solvent to create a soln, the MP of thee soln will be ______ than that of the pure solvent
greater
71
boiling point
temp at which the vapor pressure of the liquid = the ambient (incident) pressure
72
extent to which BP of soln is raised
eq
ΔTb = iKbm
ΔTb : inc in BP
i: van't hoff factor
Kb: proportionality constant (given)
m: molality
73
van't hoff factor
number of particles into which a compound dissociates in a solution
74
75
freezing point depression
eq
ΔTf = iKfm
ΔTf : freezing point depression
i: van't hoff factor
Kf: proportionality constant (given)
m: molality
76
freezing point depends on
concentration of particles, not identity
77
78
osmotic pressure
amount of pressure that must be applied to counteract the attraction of water molecules for the soln
79
osmotic pressure
eq
Π = iMRT
Π: osmotic pressure
i: van't hoff factor
M: molarity
R: ideal gas constant
T: temp
80
how are molarity and molality related for water?
nearly equal at room temp
only bc 1 L soln is approx = 1 kg solvent for dilute solutions
81
how are molarity and molality related for solvents besides water?
differ significantly because their densities are not like water
82
solubility forward reaction
dissolution
83
solubility reverse reaction
precipitation
84
solving complex ion problems
1. write the normal solubility eq with the salt and the one that makes the complex ions
1. if has mole greater than 1 --\> 2x not just x
2. add them together and multiply the Ks
3. ICE
85
Other equilibrium constants tend to follow the mass-action ratio. Write out this ratio of products and reactants that is equal to Keq or Q (if the reactants and products are not at equilibrium).
86
Given that the ion product, Qsp, is less than Ksp, which of the following will occur?
(A) More salt will dissolve
(B) No net change in salt dissolving
(C) Less salt will dissolve
(D) Not enough information given
(A) More salt will dissolve
Qsp and Ksp have the same relationship as Q and Keq. If there are less products than the equilibrium suggests should exist (Qsp \<\< Ksp), then more products will form (i.e. salt will dissolve).
87
Calculate the Ksp of PbCl2 (MM = 278.1) assuming that .14 grams of PbCl2 enters solution upon addition of 14.98 grams of PbCl2 to 147 mL of water?
(A) 3.54 ⋅ 10^-9
(B) 7.98 ⋅ 10^-8
(C) 1.60 ⋅ 10^-7
(D) 3.45 ⋅ 10^-6
(C) 1.60 ⋅ 10^-7
.14 grams PbCl2 ⋅ 1 mol PbCl2 / 278.1 g PbCl2 = approx. 5⋅10^-4 moles PbCl2 (actual: 5.03 ⋅ 10^-4)
5 ⋅ 10^-4 moles / .147 L = approx. 3.5 ⋅ 10^-3 M PbCl2 (actual: 3.42 ⋅ 10^-3)
```
[Pb2+] = approx. 3.5 ⋅ 10^-3 M
[Cl-] = approx. 7 ⋅ 10^-3 M
```
```
Ksp = [Pb2+][Cl-]^2
Ksp = (3.5 ⋅ 10^-3)(7 ⋅ 10^-3)^2
Ksp = approx. 1.75 ⋅ 10^-7 (actual: 1.60 ⋅ 10^-7)
```
88
Two common methods of defining the amount of solute in a solvent are molarity (M) and molality (m). Write out the definitions (equations) for each.
89
What is the concentration of Cu2+ at equillibrium upon the addition of Cu(OH)2 (Ksp = 2.2 ⋅ 10^-20)?
(A) 1.8 ⋅ 10^-7 M
(B) 6.9 ⋅ 10^-7 M
(C) 1.8 ⋅ 10^-8 M
(D) 6.9 ⋅ 10^-8 M
(A) 1.8 ⋅ 10^-7 M
Ksp = [Cu2+][OH-]^2
2.2 ⋅ 10^-20 = x \* ^(2x)^2
2.2 ⋅ 10^-20 = 4x^3
approx. 5 ⋅ 10^-21 = x^3
x = approx. 2 ⋅ 10^-7 (actual: 1.8 ⋅ 10^-7)
90
Suppose that a partially soluble salt dissolving is exothermic. According to Le Chatelier's principle, which of the following would happen if the mixture was heated?
(A) There is no observable change of the salt
(B) There is net dissolving of the salt
(C) There is a net precipitation of the salt
(D) Not enough information is given
(C) There is a net precipitation of the salt
Because heat was added, the system will want to remove heat. If the salt dissolving was exothermic, then the salt precipitating would need to be endothermic and would occur.
91
PbCl2 (Ksp = 1.60 ⋅ 10^-5) is added to a 8.79 ⋅ 10^-2 M solution of KCl. What is the final concentration of Pb2+?
(A) .00021 M
(B) .0021 M
(C) .021 M
(D) .21 M
(B) .0021 M
Ksp = [Pb2+][Cl-]
1.60 ⋅ 10^-5 = (x)(8.79 ⋅ 10^-2 + 2x)^2 but 2x can be approximated as 0 as compared to the concentration of Cl-:
1.60 ⋅ 10^-5 = (x)(8.79 ⋅ 10^-2)^2
1.60 ⋅ 10^-5 = (x)(approx. 81 ⋅ 10^-4)
(1.6 ⋅ 10^-5)/(8.1 ⋅ 10^-3) = x
x = approx. .2 ⋅ 10^-2 M (actual: .21 ⋅ 10^-2 M)
92
Does adding H+ to a solution of CaF2 result greater or less solubility of CaF2? Why?
Increased solubility because F- reacts with the H+, causing the solubility reaction to move forward according to Le Chatelier's principle (removing products).
93
Does adding NH3 to a solution of AgCl result greater or less solubility of AgCl? Why?
Increased solubility because NH3 reacts with the Ag to form a complex ion, causing the solubility reaction to move forward according to Le Chatelier's principle.
94
Solutions are often thought of strictly as involving liquids as the solvent. Which of the following is the general term for a solution of a solid in a solid?
(A) Solenoid
(B) Colloids
(C) Ore
(D) Alloy
(D) Alloy
The majority of elements that are solids at common temperatures are metals. On a side note, ore is more of a mixture because the different compounds can be extracted rather easily, not truly being dissolved.
95
Which of the following is NOT a colligative property?
(A) Vapor pressure depression
(B) Boiling point depression
(C) Freezing point depression
(D) Osmotic pressure
(B) Boiling point depression
Colligative properties to know are: vapor pressure depression, boiling point ELEVATION, freezing point depression, and osmotic pressure.
96
Of the following colligative properties, which of the following does NOT have the van't Hoff factor in its equation?
(A) Vapor pressure depression
(B) Boiling point elevation
(C) Freezing point depression
(D) Osmotic pressure
(A) Vapor pressure depression
Vapor pressure depression is actually modeled by Raoult's law, which states PA = XA \* PAo.
```
PA = Vapor pressure of solvent A when solutes are present
XA = Mole fraction of solvent A
PAo = Vapor pressure of solvent A in its pure state
```
97
state functions
* describe the physical properties of a system in an equilibrium state
* independent of the process of the system -\> how the system got to its current equilibrium
* not independent of other state functions
98
process function
pathway taken from one equilibrium state to another
ex: work (W) and heat (Q)
99
state functions mnemonic
When I'm under **pressure** and feeling **dense,** all I want to do is watch **TV** and get **HUGS**
pressure, density, temp, volume, enthalpy (H), internal energy (U), gibbs free energy, entropy (S)
100
standard conditions
25 deg C (298 K), 1 atm pressure, 1 M conc
101
standard temperature and pressure (STP)
0 deg C (273 K), 1 atm pressure
102
what are standard conditions used for?
kinetics, equilibrium, thermodynamics problems
103
what is STP used for?
ideal gas calculations
104
standard state
most stable and prevalent form of a substance under standard conitions
105
standard states to know
H2(g), H2O(l), NaCl(s), O2(g), C(s)
106
phase diagrams
graphs that show the standard and nonstandard states of matter for a given substance in an isolated system, as determined by temperatures and pressures
show the temps and pressures at which phases will be in equilibrium
when a substance will be thermodynamically stable in a particular phase
107
phase changes
solid, liquid, gas
reversible
exist at characteristic temps and pressures
108
fusion
melting
solid to liquid
occurs at melting point
109
freezing
crystallization or solidification
liquid to solid
occurs at freezing point
110
vaporization
evaporation or boiling
liquid to gas
(have enough kinetic energy to leave liquid phase)
111
condensation
gas to liquid
facilitated by lower temp or higher pressure
112
evaporation: endo or exothermic?
what is heat source?
endothermicc
heat source is liquid water
113
each time liquid loses high energy particle, temp of remaining liquid \_\_\_inc/dec\_\_\_
decreases
114
boiling
* specific type of vaporization
* occurs only under certain conditions
* rapid bubbling of entire liquid with rapid release of liquid as gas particles
115
boiling vs evaporation
while evaporation can happen in all liquids at all temps, boiling can only occur above the BP of a liquid and involves vaporization through the entire volume of the liquid
116
vapor pressure increases as temperature _______ because
increases because more molecules have sufficient kinetic energy to escape into the gas phase
117
vapor pressure
the pressure that gas exerts over liquid at equilibrium
118
the availability of energy microstates increases as the temperature of the solid \_\_inc/dec\_\_
increases
119
how do amorphous solids (glass, plastic, chocolate) melt?
melt (or solidify) over larger range of temperatures due to their less ordered molecular structure
120
sublimation
solid to gas (directly)
121
deposition
gas to solid
122
cold finger
may be used to purify a product that is heated under reduced pressure, causing it to sublime
123
lines of equilibrium / phase boundaries
on a phase diagram
indicate the temp and pressure values for the equilibria
124
phase diagrams
at which pressure and temp is the gas phase generally found?
high temp, low pressure
125
phase diagrams
at which pressure and temp is the solid phase generally found?
high temp, high pressure
126
phase diagrams
at which pressure and temp is the liquid phase generally found?
moderate temp, moderate pressure
127
triple point
phase digram
point at which the 3 phase boundaries meet
temp and pressure at which the 3 phases exist in equilibrium
128
critical point
phase diagram
phase boundary between the liquid and gas phases
temp and pressure which there is no distinction between the phases
129
supercritical fluids
cannot distinguish between the phases
130
temperatures above the critical point: liquid and gas phases are...
indistinguishable
131
temperature
(T)
related to the avergae kinetic energy of the particles of a substnce
132
when a subjects enthalpy increase, its temperature \_\_inc/dec\_\_
increases
133
heat vs temperature
heat is a specific form of thermal energy transferred between objects as a result of differences in their temperatures
temperature is an indirect measure of a system that looks at the average kinetic energy of particles in a system
134
heat
(Q)
process function
transfer of energy from one substance to another as a result in differences in temperature
135
zeroth law of thermodynamics
objects are in thermal equilibrium only when their temperatures are equal
136
first law of thermodynamics eq
ΔU = Q - W
U: internal energy
137
endothermic
processes in which the system absorbs heat from surroundings
138
endothermic
eq ΔQ
ΔQ \> 0
139
exothermic
processes in which the system releases heat into surroundings
140
exothermic
eq ΔQ
ΔQ \< 0
141
enthalpy
ΔH
equivalent to heat (Q) under constant pressure
142
calorimetry
process of measuring transferred heat
143
heat (q) absorbed or released in a given process
eq
calorimetry
q = mcΔT
c: specific heat of substance
ΔT: change in temp
144
specific heat
amount of energy to raise the temperature of one gram of a substance by one degree C or K
145
specific heat of H2O (l)
c = 1 cal/gK
146
heat capacity
mass x specific heat
147
bomb calorimeter
aka decomposition vessel
constant volume
isolated system
148
bomb calorimeter
eqs
no heat is exchange between calorimeter and rest of universe
ΔUsystem = -ΔUsurroundings
qsystem = -qsurrounds
qcold = -qhot
149
phase change rxns and temp
phase change reactions do not undergo changes in temperature
cannot use q = mcΔT
150
enthalpy/heat of fusion (ΔHfus)
used when
used to determine the heat transferred during phase change between solid-liquid
151
when transitioning from solid to liquid, change in enthalpy is....
positive because heat must be added
152
when transitioning from liquid to solid, the change in enthalpy is...
negative because heat must be removed
153
which eq to use
liquid gas boundary
enthalpy of vaporization (ΔHvap)
q = mL
L: latent heat
154
latent heat
enthalpy of isothermal process
155
specific heat vs heat capacity
specific heat (c) is energy required to raise the temp of one gram of a substance by 1 degree celsius
heat capacity (mc) is the product of mass and specific heat
156
constant volume vs constant pressure calorimetry
constant volume - as reaction proceeds, the temp of the contents is measured to determine the heat
constant volume - heat measured indirectly by assessing temp change in water bath around the vessel
157
endothermic eq
ΔHrxn
ΔHrxn \> 0
158
exothermic eq
ΔHrxn
ΔHrxn \< 0
159
standard enthalpy of formation
(ΔH°f)
enthalpy required to produce one mole of a compound from its elements in their standard states
160
ΔH°f of element in standard state =
0
161
standard enthalpy of rxn eq
ΔH°rxn = ΣΔH°f,products - ΣΔH°f,reactants
162
hess's law
enthalpy changes of reactions are additive
applies to any stat function, including entropy and free energy
163
bond enthalpy
avg energy that is required to break a particular type of bond between atoms in gas phase
endothermic processes
164
bond breakage is \_\_endo/exothermic\_\_\_
endothermic
positive H
165
bond enthalpy eq
ΔH°rxn = ΣΔHbonds broken - ΣΔHbonds formed = total energy absorbed - total energy released
166
standard heat of combustion
ΔH°comb
enthalpy change associated with the combustion of a fuel
167
the larger the alkane reactant, the \_\_more/less\_\_ numerous the combustion products
more
168
second law of thermodynamics
energy spontaneously disperses from being localized to becoming spread out if it is not hindered from doing so
169
entropy
measure of spontaneous dispersal of energy at specific temperatures
how much energy is spread out, or how widely spread out energy becomes
170
change in entropy eq
ΔS = Qrev/T
Qrev: heat that is gained or lost in a reversible process
171
entropy units
J/molK
172
when energy is distributed into a system at a given temperature, its entropy \_\_inc/dec\_\_.
increases
173
when energy is distributed out of a system at a given temperature, its entropy \_\_inc/dec\_\_.
decreases
174
ΔSuniverse =
ΔSsystem + ΔSsurroundings \> 0
175
gibbs free energy
measure of the change in the enthalpy and the change in entropy as a system undergoes a prcoess
indicates whether a reaction is spontaneous or not
176
change in free energy
max amount of energy released by a process -- occurring at constant temp and pressure -- that is available to perform useful work
177
gibbs free energy eq
ΔG = ΔH - TΔS
**G**oldfish are **H**orrible without (minus sign) **T**artar **S**auce
178
TΔS represents
total amount of energy that is absorbed by a system when its entropy increases reversibly
179
spontanous rxn eq
ΔG \< 0
decrease in free energy
exergonic
180
nonspontaneous rxn eq
ΔG \> 0
movement away from equilibrium position
endergonic
181
exergonic
spontaneous rxn
ΔG \< 0
182
endergonic
nonspontaneous rxn
ΔG \> 0
183
ΔG = 0
system is in equilibrium
184
ΔG \< 0
spontaneous
185
ΔG \> 0
nonspontaneous
186
ΔG is temperature dependent when
ΔH and ΔS have the same sign
187
ΔH \> 0
ΔS \> 0
outcome
spontaneous at high T
188
ΔH \> 0
ΔS \< 0
outcome
nonspontaneous at all T
189
ΔH \< 0
ΔS \> 0
outcome
spontaneous at all T
190
ΔH \< 0
ΔS \< 0
outcome
spontaneous at low T
191
ΔGºrxn =
ΔGºrxn = -RTlnKeq
192
ΔGrxn =
ΔGrxn = ΔGºrxn + RTlnQ = RTln(Q/Keq)
Q = reaction quotient
193
if Q/Keq \< 1
reaction proceeds spontaneously forward
194
if Q/Keq \> 1
reaction proceeds un reverse direction spontaneously
195
if Q/Keq = 1
reaction at equilibrium
196
open system
can exchange both energy and matter with environment
197
closed system
no exchange of matter with environment
198
internal energy
U
sum of all the different interactions between and within atoms in a system
199
ΔU in closed system =
ΔU = Q - W
200
modified standard state
[H+] = 10-7 M
pH = 7
201
reactions with more products than reactants have a more \_\_neg/pos\_\_ ΔG
negative
202
reactions with more reactants than products have a more \_\_neg/pos\_\_ ΔG
positive
203
why can heat be used as a measure of internal energy in living systems?
cellular environment has a relatively fixed volume and pressure, which eliminates work from our calculations of internal energy
204
all spontaneous reactions are
irreversible
205
What is the thermodynamic quantity that combines enthalpy and entropy? What is its units?
Gibbs free energy (ΔG) is the thermodynamic quantity that combines enthalpy and entropy. Its units are typically kJ/mol.
206
Why is Gibbs free energy being a state function important in calculating ΔG and if a series of reactions is favorable?
Gibbs free energy being a state function means that ΔG values for different reactions can be added together to see if the overall reaction is favorable or not. It allows for coupling of reactions.
207
Would the formation of the product be favored or not if the ΔG value is positive?
The formation of the product would not be favored if delta G is positive. A positive ΔG means the reaction would require huge amounts of energy to form the product.
208
Suppose that a reaction has ΔH= -77 kJ and ΔS= -0.48 kJ. At what temperature will it change from spontaneous to non-spontaneous?
(A) 47 K
(B) 160 K
(C) 243 K
(D) 321 K
(B) 160 K
```
ΔG = ΔH - TΔS
0 = (-77) - T(-0.48)
77 = 0.48T
T = about 150K (actual 160.4K)
```
At 160 Kelvin, this reaction will change from spontaneous to non-spontaneous. Because this transition is when ΔG = 0, if we substitute in the equation that ΔG = 0, we could solve for T.
209
What is the difference between heat, temperature, and enthalpy?
Heat is the transfer of energy due to change in temperature.
Temperature is the measure of the average kinetic energy of molecules.
Enthalpy is referred to as the heat transfer from the perspective of the system during reactions.
210
True or false? When looking up enthalpy values for a set of reaction species, the enthalpy depends on the concentration of the species.
True. The enthalpy depends on the concentration of the species.
211
When 1.0 mole of ZnO(s) decomposes, the ΔH = 348 kJ/mol of heat energy. This tells you that the formation of ZnO(s) is:
(A) Endothermic
(B) Exothermic
(C) In equilibrium
(D) Endergonic
(B) Exothermic
The FORMATION of ZnO(s) is exothermic because it is the reverse reaction of the decomposing.
212
What is the change in enthalpy for the following reaction: 2Mg + O2 -\> 2MgO, if ΔH Mg = 0 kJ, ΔH O2 = 0 kJ, and ΔH MgO = -501 kJ/mol?
(A) 1,002 kJ
(B) 501 kJ
(C) -501 kJ
(D) -1,002 kJ
(D) -1,002 kJ
The change in enthalpy is:
(ΔH products) - (ΔH reactants)
Thus the answer is -1,002 kJ.
213
The laws of thermodynamics dictate transformations of energy from one form to another. Which law of thermodynamics states that the total amount of energy in the universe is constant?
(A) Zeroth law of thermodynamics
(B) First law of thermodynamics
(C) Second law of thermodynamics
(D) Third law of thermodynamics
(B) First law of thermodynamics
The first law of thermodynamics states that the total amount of energy in the universe is constant.
214
The laws of thermodynamics dictate transformations of energy from one form to another. Which law of thermodynamics states that the entropy of a system approaches some constant value as its temperature approaches absolute zero?
(A) Zeroth law of thermodynamics
(B) First law of thermodynamics
(C) Second law of thermodynamics
(D) Third law of thermodynamics
(D) Third law of thermodynamics
The third law of thermodynamics states that the entropy of a system approaches some constant value as its temperature approaches absolute zero.
215
True or false? Hess's Law states that the energy change of a process is independent of the path that was taken to get there.
True. Hess's Law states that the energy change of a process is independent of the path that was taken to get there.
216
Hess's Law is true for variables such as enthalpy and entropy because these are _________ variables.
(A) Process
(B) State
(C) Fixed
(D) Variable
(B) State
Hess's Law is true for variables such as enthalpy and entropy because these are state variables.
217
At a constant pressure, the change in enthalpy is equal to what?
(A) The temperature change
(B) The heat added to the system
(C) The molarity change
(D) The disorder added to the system
(B) The heat added to the system
At a constant pressure, the change in enthalpy is equal to the heat added to the system.
218
True or false? A reaction can be both exergonic and endothermic.
True. A reaction can be both exergonic (spontaneous) and endothermic (requiring heat). However, these reactions are quite rare, and must have a high increase in entropy. A common example would be an instant icepack!
219
What is a reversible process vs. an irreversible process?
A reversible process is one in which a reaction can go forwards and backwards without losing any energy, right near equilibrium in an ideal world.
An irreversible process is one in which a reaction can't go forwards and backwards without losing any energy.
220
When the pressure on a gaseous reaction system at equilibrium is increased, the equilibrium will shift in which direction according to Le Chatlier's principle?
The equilibrium will favor the side with the least gas molecules.
221
For the following reaction, how will the reaction equilibrium be affected by an increase in temperature? H2O2(l) -\> H2(g) + O2(g), delta H = 187 kJ
Since delta H is a positive value it indicates that it is an endothermic reaction and the energy is on the reactant side because it is being absorbed. Thus an increase in the temperature of the system causes equilibrium to shift towards the right, towards the products H2(g) and O2(g) because this causes the reaction to shift away from the heat to balance out the reaction.
222
For the following reaction, how will the reaction equilibrium be affected by an increase in volume?
H2O2(l) = H2(g) + O2(g)
The reaction will shift to the right, thus increasing the products. H2O2 is a liquid, so it cannot change its volume. H2 (g) and O2 (g) can both take advantage of the increased volume by expanding and spreading out the molecules. Thus an increase in H2O2 will shift the reaction to the right towards H2 (g) and O2 (g).
223
If volume is decreased for the following reaction system at equilibrium, what color will you expect the mixture to become?
N2O4(g) (colorless) + heat = 2 NO2(g) (pink)
You would expect the mixture to become more colorless because there are less gas molecules on that side of the equation. Shifting the equilibrium to the side with less gas molecules will relieve the stress of decreasing volume (which inherently means increasing pressure).
224
a) ores are mixtures, whereas alloys are solutions
b) ores are found only below sea level and are mined for, whereas alloys are found above sea level
c) ores are solutions found in nature, whereas alloys are man made solutions
d) ores strictly describe metals that are trapped in igneous rock, and are a subclass of alloys
a) ores are mixtures, whereas alloys are solutions
225
the reaction quotient, Q, tends to change so it approaches the equilibrium constant, K. which of the following statements is not true when Qsp \> Ksp?
a) A common laboratory technique to purify compounds is to allow Q to be greater than K, followed by extracting that desired salt while the impurities remain soluble.
b) This can be caused by saturating a solution while the solvent is chilled, followed by letting the solution heat up.
c) Q can approach K by having some of the salt recrystallize, or precipitate out
d) This can be caused by a solution having some of its solvent evaporate
a) A common laboratory technique to purify compounds is to allow Q to be greater than K, followed by extracting that desired salt while the impurities remain soluble
226
a) 1, with NaCl having the greater vant hoff factor
b) 3, with K2SO4 having the greater vant hoff factor
c) 1, with K2SO4 having the greater vant hoff factor
d) 5, with K2SO4 having the greater vant hoff factor
c) 1, with K2SO4 having the greater vant hoff factor
227
a) [Ag+] is greater for pure water soon; [Cl-] is greater for NH3 soln
b) [Ag+] and [Cl-] are greater for NH3 soln
c) [Ag+] and [Cl-] are greater for pure water soln
d) [Cl-] is greater for the pure water soln; [Ag+] is greater for NH3 soln
a) [Ag+] is greater for pure water soon; [Cl-] is greater for NH3 soln
228
a) decrease the temp
b) add CaCl2
c) add HCl
d) stir the soln
c) add HCl
229
a) neither copper (II) carbonate nor cobalt (II) carbonate will precipitate
b) copper (II) carbonate will precipitate
c) cobalt (II) carbonate will precipitate
d) both copper (II) carbonate and cobalt (II) carbonate will precipitate
c) cobalt (II) carbonate will precipitate
230
a) 1.8 x 10^05 g/L
b) 2.1 x 10^-10 g/L
c) 1 x 10^10 g/L
d) 9 x 10^-6 g/L
a) 1.8 x 10^05 g/L
231
a) 2.3 x10^-5
b) 1.1 x 10^-5
c) 8 x 10^-8
d) 1.6 x 10^-7
d) 1.6 x 10^-7
232
a) III only
b) I and II only
c) I, II, and III
d) I and III only
b) I and II only
233
a) he did not account for the air pressure in the lab
b) he did not account for the conc of the compounds used in lab
c) he had no theoretical mistakes, and physical mistakes probably caused these
d) he did not account for stoichiometry in enthalpy calc
d) he did not account for stoichiometry in enthalpy calc
234
a) switching to a larger container
b) adding catalyst
c) adding inert gas
d) adding CaCO3
a) switching to a larger container
235
d
236
b
237
d
238
c
239
a
240
d
241
b
242
b
243
b
244
d
245
b
