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MCAT > Chem I: 11-12 > Flashcards

Chem I: 11-12 Flashcards

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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.

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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.

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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

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4
Q

solution

A

homogenous mixtures of 2+ substances that combine to form a single phase

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5
Q

relationship between mixtures and solutions

A

all solutions are considered mixtures, but not all mixtures are considered solutions

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6
Q

solute

A

dissolved in a solvent

ex: NaCl, NH3, CO2, glucose

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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

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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
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9
Q

if solvation is exothermic…

process is favored at ___ temperatures

A

new interactions are stronger than the original ones

low temp

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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

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11
Q

ideal solution

A

when enthalpy of dissolution is 0

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12
Q

spontaneous formation of solutions

exothermic vs endothermic

A

both can form spontaneously

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13
Q

at constant temp and pressure, entropy always ______ upon dissolution

A

increases

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14
Q

solubility

A

max amount of that substance that can be dissolved in a particular solvent at a given temp

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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
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16
Q

dilute

A

solution in which the proportion of solute to solvent is small

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17
Q

hydration

A

solvation in water

water molecules break ionic bonds

ions surrounded and stabilized by shell of solvent molecules

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18
Q

hydration rxn

NaCl (s) –>

A

Na+(aq) + Cl-(aq)

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19
Q

precipitation rxn

A

ions come together to form a solid that falls out of solution

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20
Q

precipitation rxn

NaCl(aq) + AgNO3(aq) –>

A

AgCl(s)

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21
Q

concentrated

A

solution in which the proportion of solute to solvent is small

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22
Q

sparingly soluble salts

A

solutes that dissolve minimally in the solvent

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23
Q

aqueous solution

A

solvent is water

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24
Q

H+ is never found alone in solution bc….

A

a free proton is difficult to isolate

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25
Q

most important solubility rules

A

all salts of Group I metals and all nitrate salts are soluble

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26
Q

7 general solubility rules

A
  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
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27
Q

complex ion

A

aka coordination compound

moleculae in which a cation is bonded to at lease one electron pair donor

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28
Q

ligands

A

electron pair donor molecule in complex ion

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29
Q

coordinate covalent bond

A

hold complex ions together

electron pair donor and electron pair acceptor from very stable lewis acid-base adducts

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30
Q

complex formation biological applications

A
  • active sites of proteins
    • iron cation in hemoglobin
  • coenzymes (vitamins) and cofactors
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31
Q

chelation

A

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

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32
Q

concentration

A

amount of solute dissolved in a solvent

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33
Q

percent composition by mass eq

A

mass of solute/mass of solution x 100%

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34
Q

mole fraction

eq

A

XA = moles of A/total moles of all species

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35
Q

sum of mole fractions in a system will always =

A

1

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36
Q

molarity eq

A

M = moles of solute/liters of solution

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37
Q

molality eq

A

m = moles of solute/kg of solvent

true only for dilute aqueous solutions

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38
Q

eq used to determin conc of a solution after dilution

A

MiVi=MfVf

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39
Q

saturation point

A

solution concentration is at its max value for the given temp and pressure

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40
Q

when solution is dilute, the thermodynamically favored process is ______.

initially, the rate of ______…..

A

dissolution

initially, the rate of dissolution will be greater than the rate of precipitation.

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41
Q

as solution becomes more concentrated and approaches saturation, the rate of dissolution ______, while the rate of precipitation _____.

A

lessens

increases

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42
Q

degree of solubility determined by:

A

relative changes in enthalpy and entropy associated with the dissolution of the ionic solute at a given temp and pressure

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43
Q

solubility product constant eq

A

Ksp = [An+]m[Bm-]n

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44
Q

Ksp dependent on

A

temperature

(and pressure for gases)

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45
Q

As temp increases, Ksp

A

increases for non gas solutes and decreases for gas solutes

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46
Q

ion product eq

A

IIP = [An+]m[Bm-]n

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47
Q

difference between ion product and Ksp

A

concentrations used in IP are concentrations of the ionic constituents at that given moment in time, which may differ from Ksp

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48
Q

IP < Ksp

A

unsaturated -> soln not yet at equilibirum

solute will continue to dissolve

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49
Q

IP = Ksp

A

saturated

solution is at equilibrium

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50
Q

IP > Ksp

A

supersaturated -> solution is beyond equilibrium

precipitation will occur

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51
Q

supersaturated soln

A
  • beyond equilibrium
  • thermodynamically unstable
  • any disturbance to soln will cause spontaneous precipitation of the excess dissolved solute
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52
Q

molar solubility

A

molarity of a solute in a saturated soln

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53
Q
A
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54
Q
A
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55
Q

every sparingly soluble salt of general formula MX will have Ksp=

A

Ksp = x2

x: molar solubility (assuming no common ion effect)

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56
Q

every sparingly soluble salt of general formula MX2 will have Ksp=

A

Ksp = 4x3

x: molar solubility (assuming no common ion effect)

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57
Q

every sparingly soluble salt of general formula MX3 will have Ksp=

A

Ksp = 27x4

x: molar solubility (assuming no common ion effect)

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58
Q

formation of complex ions _____ the solubility of salt in a soln

A

increases

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59
Q

formation or stability constant

A

Kf

for ocmplex ions

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60
Q
A
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61
Q

common ion effect

A

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

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62
Q

colligative properties

A

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

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63
Q

vapor pressure depression

A
  • 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
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64
Q

as conc of B increases, vapor pressure of A ______

A

decreases

as more solute is dissolved into solvent (as more B is dissolved into A), the vapor pressure of the solvent decreases

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65
Q

raoult’s law

eq

A

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

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66
Q

raoult’s law hold only when…

A

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

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67
Q

as vapor pressure decreases, boiling point ______

why?

A

increases

higher temp is required to match the atmospheric pressure

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68
Q

ideal solutions

A

solutions that obey raoult’s law

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69
Q
A
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70
Q

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

A

greater

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71
Q

boiling point

A

temp at which the vapor pressure of the liquid = the ambient (incident) pressure

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72
Q

extent to which BP of soln is raised

eq

A

ΔTb = iKbm

ΔTb : inc in BP

i: van’t hoff factor

Kb: proportionality constant (given)

m: molality

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73
Q

van’t hoff factor

A

number of particles into which a compound dissociates in a solution

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74
Q
A
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75
Q

freezing point depression

eq

A

ΔTf = iKfm

ΔTf : freezing point depression

i: van’t hoff factor

Kf: proportionality constant (given)

m: molality

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76
Q

freezing point depends on

A

concentration of particles, not identity

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77
Q
A
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78
Q

osmotic pressure

A

amount of pressure that must be applied to counteract the attraction of water molecules for the soln

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79
Q

osmotic pressure

eq

A

Π = iMRT

Π: osmotic pressure

i: van’t hoff factor

M: molarity

R: ideal gas constant

T: temp

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80
Q

how are molarity and molality related for water?

A

nearly equal at room temp

only bc 1 L soln is approx = 1 kg solvent for dilute solutions

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81
Q

how are molarity and molality related for solvents besides water?

A

differ significantly because their densities are not like water

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82
Q

solubility forward reaction

A

dissolution

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83
Q

solubility reverse reaction

A

precipitation

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84
Q

solving complex ion problems

A
  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
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85
Q

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).

A
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86
Q

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

(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).

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87
Q

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

A

(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)
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88
Q

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.

A
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89
Q

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

(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)

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90
Q

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

A

(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.

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91
Q

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

A

(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)

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92
Q

Does adding H+ to a solution of CaF2 result greater or less solubility of CaF2? Why?

A

Increased solubility because F- reacts with the H+, causing the solubility reaction to move forward according to Le Chatelier’s principle (removing products).

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93
Q

Does adding NH3 to a solution of AgCl result greater or less solubility of AgCl? Why?

A

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.

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94
Q

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

A

(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.

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95
Q

Which of the following is NOT a colligative property?

(A) Vapor pressure depression
(B) Boiling point depression
(C) Freezing point depression
(D) Osmotic pressure

A

(B) Boiling point depression

Colligative properties to know are: vapor pressure depression, boiling point ELEVATION, freezing point depression, and osmotic pressure.

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96
Q

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

(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
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97
Q

state functions

A
  • 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
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98
Q

process function

A

pathway taken from one equilibrium state to another

ex: work (W) and heat (Q)

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99
Q

state functions mnemonic

A

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
Q

standard conditions

A

25 deg C (298 K), 1 atm pressure, 1 M conc

101
Q

standard temperature and pressure (STP)

A

0 deg C (273 K), 1 atm pressure

102
Q

what are standard conditions used for?

A

kinetics, equilibrium, thermodynamics problems

103
Q

what is STP used for?

A

ideal gas calculations

104
Q

standard state

A

most stable and prevalent form of a substance under standard conitions

105
Q

standard states to know

A

H2(g), H2O(l), NaCl(s), O2(g), C(s)

106
Q

phase diagrams

A

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
Q

phase changes

A

solid, liquid, gas

reversible

exist at characteristic temps and pressures

108
Q

fusion

A

melting

solid to liquid

occurs at melting point

109
Q

freezing

A

crystallization or solidification

liquid to solid

occurs at freezing point

110
Q

vaporization

A

evaporation or boiling

liquid to gas

(have enough kinetic energy to leave liquid phase)

111
Q

condensation

A

gas to liquid

facilitated by lower temp or higher pressure

112
Q

evaporation: endo or exothermic?

what is heat source?

A

endothermicc

heat source is liquid water

113
Q

each time liquid loses high energy particle, temp of remaining liquid ___inc/dec___

A

decreases

114
Q

boiling

A
  • specific type of vaporization
  • occurs only under certain conditions
  • rapid bubbling of entire liquid with rapid release of liquid as gas particles
115
Q

boiling vs evaporation

A

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
Q

vapor pressure increases as temperature _______ because

A

increases because more molecules have sufficient kinetic energy to escape into the gas phase

117
Q

vapor pressure

A

the pressure that gas exerts over liquid at equilibrium

118
Q

the availability of energy microstates increases as the temperature of the solid __inc/dec__

A

increases

119
Q

how do amorphous solids (glass, plastic, chocolate) melt?

A

melt (or solidify) over larger range of temperatures due to their less ordered molecular structure

120
Q

sublimation

A

solid to gas (directly)

121
Q

deposition

A

gas to solid

122
Q

cold finger

A

may be used to purify a product that is heated under reduced pressure, causing it to sublime

123
Q

lines of equilibrium / phase boundaries

A

on a phase diagram

indicate the temp and pressure values for the equilibria

124
Q

phase diagrams

at which pressure and temp is the gas phase generally found?

A

high temp, low pressure

125
Q

phase diagrams

at which pressure and temp is the solid phase generally found?

A

high temp, high pressure

126
Q

phase diagrams

at which pressure and temp is the liquid phase generally found?

A

moderate temp, moderate pressure

127
Q

triple point

A

phase digram

point at which the 3 phase boundaries meet

temp and pressure at which the 3 phases exist in equilibrium

128
Q

critical point

A

phase diagram

phase boundary between the liquid and gas phases

temp and pressure which there is no distinction between the phases

129
Q

supercritical fluids

A

cannot distinguish between the phases

130
Q

temperatures above the critical point: liquid and gas phases are…

A

indistinguishable

131
Q

temperature

A

(T)

related to the avergae kinetic energy of the particles of a substnce

132
Q

when a subjects enthalpy increase, its temperature __inc/dec__

A

increases

133
Q

heat vs temperature

A

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
Q

heat

A

(Q)

process function

transfer of energy from one substance to another as a result in differences in temperature

135
Q

zeroth law of thermodynamics

A

objects are in thermal equilibrium only when their temperatures are equal

136
Q

first law of thermodynamics eq

A

ΔU = Q - W

U: internal energy

137
Q

endothermic

A

processes in which the system absorbs heat from surroundings

138
Q

endothermic

eq ΔQ

A

ΔQ > 0

139
Q

exothermic

A

processes in which the system releases heat into surroundings

140
Q

exothermic

eq ΔQ

A

ΔQ < 0

141
Q

enthalpy

A

ΔH

equivalent to heat (Q) under constant pressure

142
Q

calorimetry

A

process of measuring transferred heat

143
Q

heat (q) absorbed or released in a given process

eq

A

calorimetry

q = mcΔT

c: specific heat of substance

ΔT: change in temp

144
Q

specific heat

A

amount of energy to raise the temperature of one gram of a substance by one degree C or K

145
Q

specific heat of H2O (l)

A

c = 1 cal/gK

146
Q

heat capacity

A

mass x specific heat

147
Q

bomb calorimeter

A

aka decomposition vessel

constant volume

isolated system

148
Q

bomb calorimeter

eqs

A

no heat is exchange between calorimeter and rest of universe

ΔUsystem = -ΔUsurroundings

qsystem = -qsurrounds

qcold = -qhot

149
Q

phase change rxns and temp

A

phase change reactions do not undergo changes in temperature

cannot use q = mcΔT

150
Q

enthalpy/heat of fusion (ΔHfus)

used when

A

used to determine the heat transferred during phase change between solid-liquid

151
Q

when transitioning from solid to liquid, change in enthalpy is….

A

positive because heat must be added

152
Q

when transitioning from liquid to solid, the change in enthalpy is…

A

negative because heat must be removed

153
Q

which eq to use

liquid gas boundary

A

enthalpy of vaporization (ΔHvap)

q = mL

L: latent heat

154
Q

latent heat

A

enthalpy of isothermal process

155
Q

specific heat vs heat capacity

A

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
Q

constant volume vs constant pressure calorimetry

A

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
Q

endothermic eq

ΔHrxn

A

ΔHrxn > 0

158
Q

exothermic eq

ΔHrxn

A

ΔHrxn < 0

159
Q

standard enthalpy of formation

A

(ΔH°f)

enthalpy required to produce one mole of a compound from its elements in their standard states

160
Q

ΔH°f of element in standard state =

A

0

161
Q

standard enthalpy of rxn eq

A

ΔH°rxn = ΣΔH°f,products - ΣΔH°f,reactants

162
Q

hess’s law

A

enthalpy changes of reactions are additive

applies to any stat function, including entropy and free energy

163
Q

bond enthalpy

A

avg energy that is required to break a particular type of bond between atoms in gas phase

endothermic processes

164
Q

bond breakage is __endo/exothermic___

A

endothermic

positive H

165
Q

bond enthalpy eq

A

ΔH°rxn = ΣΔHbonds broken - ΣΔHbonds formed = total energy absorbed - total energy released

166
Q

standard heat of combustion

A

ΔH°comb

enthalpy change associated with the combustion of a fuel

167
Q

the larger the alkane reactant, the __more/less__ numerous the combustion products

A

more

168
Q

second law of thermodynamics

A

energy spontaneously disperses from being localized to becoming spread out if it is not hindered from doing so

169
Q

entropy

A

measure of spontaneous dispersal of energy at specific temperatures

how much energy is spread out, or how widely spread out energy becomes

170
Q

change in entropy eq

A

ΔS = Qrev/T

Qrev: heat that is gained or lost in a reversible process

171
Q

entropy units

A

J/molK

172
Q

when energy is distributed into a system at a given temperature, its entropy __inc/dec__.

A

increases

173
Q

when energy is distributed out of a system at a given temperature, its entropy __inc/dec__.

A

decreases

174
Q

ΔSuniverse =

A

ΔSsystem + ΔSsurroundings > 0

175
Q

gibbs free energy

A

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
Q

change in free energy

A

max amount of energy released by a process – occurring at constant temp and pressure – that is available to perform useful work

177
Q

gibbs free energy eq

A

ΔG = ΔH - TΔS

Goldfish are Horrible without (minus sign) Tartar Sauce

178
Q

TΔS represents

A

total amount of energy that is absorbed by a system when its entropy increases reversibly

179
Q

spontanous rxn eq

A

ΔG < 0

decrease in free energy

exergonic

180
Q

nonspontaneous rxn eq

A

ΔG > 0

movement away from equilibrium position

endergonic

181
Q

exergonic

A

spontaneous rxn

ΔG < 0

182
Q

endergonic

A

nonspontaneous rxn

ΔG > 0

183
Q

ΔG = 0

A

system is in equilibrium

184
Q

ΔG < 0

A

spontaneous

185
Q

ΔG > 0

A

nonspontaneous

186
Q

ΔG is temperature dependent when

A

ΔH and ΔS have the same sign

187
Q

ΔH > 0

ΔS > 0

outcome

A

spontaneous at high T

188
Q

ΔH > 0

ΔS < 0

outcome

A

nonspontaneous at all T

189
Q

ΔH < 0

ΔS > 0

outcome

A

spontaneous at all T

190
Q

ΔH < 0

ΔS < 0

outcome

A

spontaneous at low T

191
Q

ΔGºrxn =

A

ΔGºrxn = -RTlnKeq

192
Q

ΔGrxn =

A

ΔGrxn = ΔGºrxn + RTlnQ = RTln(Q/Keq)

Q = reaction quotient

193
Q

if Q/Keq < 1

A

reaction proceeds spontaneously forward

194
Q

if Q/Keq > 1

A

reaction proceeds un reverse direction spontaneously

195
Q

if Q/Keq = 1

A

reaction at equilibrium

196
Q

open system

A

can exchange both energy and matter with environment

197
Q

closed system

A

no exchange of matter with environment

198
Q

internal energy

A

U

sum of all the different interactions between and within atoms in a system

199
Q

ΔU in closed system =

A

ΔU = Q - W

200
Q

modified standard state

A

[H+] = 10-7 M

pH = 7

201
Q

reactions with more products than reactants have a more __neg/pos__ ΔG

A

negative

202
Q

reactions with more reactants than products have a more __neg/pos__ ΔG

A

positive

203
Q

why can heat be used as a measure of internal energy in living systems?

A

cellular environment has a relatively fixed volume and pressure, which eliminates work from our calculations of internal energy

204
Q

all spontaneous reactions are

A

irreversible

205
Q

What is the thermodynamic quantity that combines enthalpy and entropy? What is its units?

A

Gibbs free energy (ΔG) is the thermodynamic quantity that combines enthalpy and entropy. Its units are typically kJ/mol.

206
Q

Why is Gibbs free energy being a state function important in calculating ΔG and if a series of reactions is favorable?

A

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
Q

Would the formation of the product be favored or not if the ΔG value is positive?

A

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
Q

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

A

(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
Q

What is the difference between heat, temperature, and enthalpy?

A

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
Q

True or false? When looking up enthalpy values for a set of reaction species, the enthalpy depends on the concentration of the species.

A

True. The enthalpy depends on the concentration of the species.

211
Q

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

A

(B) Exothermic

The FORMATION of ZnO(s) is exothermic because it is the reverse reaction of the decomposing.

212
Q

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

A

(D) -1,002 kJ

The change in enthalpy is:
(ΔH products) - (ΔH reactants)

Thus the answer is -1,002 kJ.

213
Q

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

A

(B) First law of thermodynamics

The first law of thermodynamics states that the total amount of energy in the universe is constant.

214
Q

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

A

(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
Q

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.

A

True. Hess’s Law states that the energy change of a process is independent of the path that was taken to get there.

216
Q

Hess’s Law is true for variables such as enthalpy and entropy because these are _________ variables.

(A) Process
(B) State
(C) Fixed
(D) Variable

A

(B) State

Hess’s Law is true for variables such as enthalpy and entropy because these are state variables.

217
Q

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

A

(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
Q

True or false? A reaction can be both exergonic and endothermic.

A

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
Q

What is a reversible process vs. an irreversible process?

A

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
Q

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?

A

The equilibrium will favor the side with the least gas molecules.

221
Q

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

A

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
Q

For the following reaction, how will the reaction equilibrium be affected by an increase in volume?

H2O2(l) = H2(g) + O2(g)

A

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
Q

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)

A

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
Q

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

a) ores are mixtures, whereas alloys are solutions

225
Q

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) 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
Q

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

A

c) 1, with K2SO4 having the greater vant hoff factor

227
Q

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

a) [Ag+] is greater for pure water soon; [Cl-] is greater for NH3 soln

228
Q

a) decrease the temp
b) add CaCl2
c) add HCl
d) stir the soln

A

c) add HCl

229
Q

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

A

c) cobalt (II) carbonate will precipitate

230
Q

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

a) 1.8 x 10^05 g/L

231
Q

a) 2.3 x10^-5
b) 1.1 x 10^-5
c) 8 x 10^-8
d) 1.6 x 10^-7

A

d) 1.6 x 10^-7

232
Q

a) III only
b) I and II only
c) I, II, and III
d) I and III only

A

b) I and II only

233
Q

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

A

d) he did not account for stoichiometry in enthalpy calc

234
Q

a) switching to a larger container
b) adding catalyst
c) adding inert gas
d) adding CaCO3

A

a) switching to a larger container

235
Q
A

d

236
Q
A

b

237
Q
A

d

238
Q
A

c

239
Q
A

a

240
Q
A

d

241
Q
A

b

242
Q
A

b

243
Q
A

b

244
Q
A

d

245
Q
A

b

MCAT (26 decks)

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