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CHEM 1310 - 2017 > Finals > Flashcards

Finals Flashcards

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1
Q

What is measured at constant pressure using a calorimeter?

A

delta H

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2
Q

Third Law of Thermodynamics

A

the entropy of a pure, perfectly crystalline substance at absolute zero (0K) is zero

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3
Q

Which are state functions and path functions?

A

state function: T, E, H, S, G, p, V, m, composition

path function: t,w and q

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4
Q

A reaction has a molar enthalpy (delta H) of formation of 50.63 kJ/mol. Under what conditions are the signs of delta E, w and q for this reaction?

A

*endothermic reaction
delta E = negative
w = negative
q = negative

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5
Q

How would you order the most stable to least stable to thermally decompose their elements under standard state conditions?

A
  • Entropy (S)
    1) most to least stable: solid, liquid, gas
    2) bigger the molar mass, the less stable it is
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6
Q

First Law of Thermodynamics

A

Energy may be transferred as work or heat, but no energy can be lost, nor can heat or work be obtained from nothing

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7
Q

Second Law of Thermodynamics

A

There is an increase in entropy (disorder) as a spontaneous reaction occur in the universe; some energy lost as heat

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8
Q

The delta T is dependent on:

A

 on the amount of heat transferred, q
 on the direction of heat flow
 inversely on the amount of material
 on the identity of the material.

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9
Q

Heat capacity (C)

A

quantity of energy required to increase the temperature of a sample by 1°C

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10
Q

C is dependent on:

A
  • mass of substance
  • type of material
  • state
  • temperature
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11
Q

Specific heat capacity (Cs)

A

amount of heat energy required to raise the temperature of one g of a substance by 1°C
Jg-1K-1

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12
Q

Work (w)

A

energy used to move an object against an opposing force

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13
Q

w is dependent on:

A
  • magnitude of the applied force

- displacement distance

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14
Q

As T increases, what happens to delta G

A

increases

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15
Q

Bond energy

A

energy required to break or make a bond
bond making= negative delta E
bond breaking= positive delta E

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16
Q

Exothermic vs Endothermic

A

Exothermic:if the reaction releases heat.  Endothermic: if the reaction absorbs heat.

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17
Q

Enthalpy (delta H)

A

heat transferred into or out of a system at constant pressure.

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18
Q

Heat of vaporization

A

heat required to convert liquid to gas

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19
Q

Hess’s Law

A

A change in any state function is independent of path.
Thus, the energy change in a chemical reaction is independent of the manner in which the reaction takes place.
The enthalpy change for any overall process is equal to the sum of enthalpy changes for any set of steps that leads from the starting materials to the products.

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20
Q

Standard state

A

most stable form of a substance at T = 25 °C and p = 1 bar, and 1 M if it is in solution

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21
Q

delta H is dependent on:

A

temperature, concentration and pressure

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22
Q

Spontaneous direction

A
  • preferred direction of reaction under specific conditions

- can only be reversed by the action of an outside force

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23
Q

What is the general tendency of heat?

A

Heat flows from high temperature to low temperature

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24
Q

Entropy (delta S)

A

the quantitative measure of dispersal resulting from an energy transfer
units: JK-1 or JK-1mol-1

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25
Q

As phases change from a solid to a liquid, and temperature increases, what happens to entropy?

A

entropy increases

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26
Q

delta S depends on:

A
  • temperature
  • phase changes (solid is most stable, gas is least satble)
  • molar mass (larger the molecule, greater the entropy
  • quantity of matter
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27
Q

can S ever be absolute zero?

A

no never at any temperature above 0K

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28
Q

Gibbs Free Energy (delta G)

A

Maximum useful work that can be obtained from a thermodynamic system at constant T, P.
Predicts in which direction a reaction is spontaneous

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29
Q

What does a negative and positive delta G indicate?

A

∆G > 0, process is reactant favoured

∆G < 0, process is product favoured

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30
Q

A reaction is spontaneous at all temperatures when ΔH°, ΔS° and ΔG° are?

A

ΔH° = -
ΔS° = +
ΔG° (high temperature)= -
ΔG° (low temperature)= -

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31
Q

A reaction is spontaneous at no temperatures when ΔH°, ΔS° and ΔG° are?

A

ΔH° = +
ΔS° = -
ΔG° (high temperature)= +
ΔG° (low temperature)= +

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32
Q

A reaction is spontaneous at high temperatures (entropy driven) when ΔH°, ΔS° and ΔG° are?

A

ΔH° = +
ΔS° = +
ΔG° (high temperature)= -
ΔG° (low temperature)= +

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33
Q

A reaction is spontaneous at low temperatures (enthalpy driven) when ΔH°, ΔS° and ΔG° are?

A

ΔH° = -
ΔS° = -
ΔG° (high temperature)= +
ΔG° (low temperature)= -

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34
Q

For a reaction to be thermodynamically favourable, what sign must be delta G?

A

negative delta G

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35
Q

The spontaneous direction of a phase change is dependent on:

A

pressure

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36
Q

Reaction mechanisms

A

the exact molecular pathway that starting materials follow on their way to becoming products

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37
Q

Reaction intermediates

A

 species that are not part of the reaction stoichiometry

 created in one step and consumed in a later step

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38
Q

Rate-determining step

A
  • the slowest elementary step in a mechanism
  • governs the rate of the overall chemical reaction because no net chemical reaction can go faster than its slowest step
  • reaction that requires the most energy to proceed (highest activation energy)
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39
Q

rate is dependent on:

A
  • concentration of a reactant
  • temperature
  • catalysts
    units: molL-1s-1
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40
Q

Rate law

A

the effect of concentration on the rate of a particular chemical reaction

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41
Q

The rate law of an elementary reaction can be found:

A

written directly from the stoichiometry of the reactants

*****ONLY FOR ELEMENTARY REACTIONS

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42
Q

An overall reaction’s rate law can be found by:

A

the rate determining step since it is the slowest reaction

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43
Q

Half-life

A

the time required for the reactant concentration to drop to one-half its original value, t1/2

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44
Q

What are the units of the zero, first and second order?

A

zero=Ms-1
first=s-1
second=M-1s-1

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45
Q

Mechanism

A
  • one or more elementary reactions describing how the chemical reaction occurs
  • sum of the individual steps in the mechanism must give the overall balanced chemical equation
  • reaction mechanism must be consistent with the experimental rate law
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46
Q

Activation energy

A

amount of energy to overcome an energy barrier

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47
Q

Activation energy is dependent on:

A

-temperature= as temperature increases, more reactants will have sufficient energy to overcome the activation barrier

48
Q

What must happen in order for a reaction to be successful?

A

 There must be sufficient collision energy  Reactants must be the correctly orientated
 The lower the probability of alignment, the lower the value of k and the slower the reaction.

49
Q

catalysts

A
  • substances that speed up the rate of a chemical reaction
  • not consumed in a chemical reaction
  • provide a different pathway with a lower activation energy.
  • have no effect on equilibrium
50
Q

Heterogeneous catalysts vs Homogeneous catalysts

A

Heterogeneous: catalyst is in a different phase than the reacting substance
Homogeneous: catalyst is in the same phase as the reacting substance

51
Q

Equilibrium constant (Keq)

A
  • related to the stoichiometry of the balanced net reaction.
  • applies only at equilibrium.
  • independent of initial conditions
52
Q

What happens to Keq when

A

 If you flip an equation, Keq (forward) = 1/Keq (reverse)
 If you add equations together,
Keq (overall) = Keq,1Keq,2
 If you multiply an equation by a coefficient n, then Keq,new = (old Keq)^n

53
Q

What is the relationship between Q and Keq?

A
  1. If Q< Keq, the reaction goes to the right to make products.
  2. If Q= Keq, the reaction is at equilibrium and there is no net change.
  3. If Q> Keq, the reaction goes to the left to make reactants.
54
Q

What is the relationship between Keq and temperature?

A

1) Keq of an exothermic reaction decreases with increasing temperature.
2) Keq of an endothermic reaction increases with increasing temperature.

55
Q

Le Chatelier’s principle

A
  • when a change is imposed on a system at equilibrium, the system will react in the direction that reduces the amount of change.
    ex) addition or removal of reactants, products and heat; in gases, addition or removal of volume and pressure
  • *****does not apply to the addition of catalysts. Catalysts only change the rate of the reaction, not the equilibrium conditions.
56
Q

Approximation method using ice tables

A

If the approximate value of x is less than 5% of the initial concentration, the approximation is valid!!

57
Q

Acid

A

donates a proton

58
Q

Base

A

accepts a proton

59
Q

Spectator ions

A

ionic species that undergo no significant reactions

60
Q

Monoprotic acids, Monoprotic bases, Polyprotic acids, Polyprotic bases

A

Monoprotic acids: can only donate one proton
Monoprotic bases: can only accept one proton
Polyprotic acids: can donate two or more protons
Polyprotic bases: can accept two or more protons

61
Q

Amphiprotic

A

molecules or ions which can behave either way (either an acid or a base)

62
Q

What happens to strong acids and baces

A

dissociate 100%

63
Q

***Strong acids

A 
HCl Hydrochloric acid 
HBr Hydrobromic acid 
HI Hydroiodic acid 
HNO3 Nitric acid 
HClO4 Perchloric acid 
H2SO4 Sulfur acid
64
Q

***Strong bases

A

LiOH Lithium hydroxide
NaOH Sodium hydroxide
KOH Potassium hydroxide
Any soluble hydroxide (Group 1 and 2)

65
Q

Hydronium ion

A

H3O+

66
Q

Keq of water pure water at 25 degrees celsius and 1 bar

A

1*10^-14

67
Q

Kw is dependent on:

A

equilibrium constant of pure water is dependent on temperature

68
Q

A change of pH by one unit results in what change in [hyrdonium ion]?

A

tenfold change

69
Q

All dissociations of weak acids and bases are determined by what, since it does not dissociate 100%?

A

equilibrium constants Ka and Kb

70
Q

Oxoacid

A

an acid that contains an inner atom bonded to a number of oxygen atoms and acidic OH groups

71
Q
  • Anions that are conjugate bases of weak acids make a solution ______.
  • Cations that are conjugate acids of weak bases make a solution _____.
A
  • Anions that are conjugate bases of weak acids make a solution basic.
  • Cations that are conjugate acids of weak bases make a solution acidic.
72
Q

The stronger the acid, the ____ the conjugate base.

A

weaker

73
Q

Rules on salts

A

 Cations are potential acids, anions are potential bases
 Anions of strong acids do not affect pH (Cl-, Br-, I-, NO3-, ClO4-)
 +1 and +2 metal cations do not affect pH  The anions of weak acids are weak bases and thus make the solution basic (RCOO-, CN-, F-, NO2-,…)
 The cations of weak bases are weak acids and thus make the solution acidic (NH4+, RNH3+)

74
Q

As acid strength increases, Ka value______ and the negative charge ______, polarity (electronegativity) _____, bond energy ______

A

increases
decreases
increases
weakens

75
Q

How does the structure of a molecule make a solution more acidic?

A
  • In order to donate a proton, a molecule must break a H–X bond.
  • Electronegative atoms withdraw electron density from O–H and thus weaken this bond, and thus increase acidity
76
Q

Indicator

A

an acid-base indicator changes colour based on pH and helps determine the endpoint of a reaction

77
Q

Endpoint

A

reaction is complete indicated by the physical change in colour of the solution

78
Q

Equivalence point / Stoichiometric endpoint

A

when the moles of acid = moles of base

79
Q

How do you choose the best indicator for your reaction?

A

choose one that gives an endpoint very close to the equivalence point

80
Q

As oxygen atoms are added to an acid, electronegativity ______ and acid strength ______

A

increases

increases

81
Q

Half-equivalence point

A

half of the weak acid has been converted to conjugate base. [HA] = [A–]

82
Q

In the titration of a polyprotic acid, what happens?

A

two equivalence points, the closer Ka1 and Ka2 are to each other, the less distinguishable the equivalence points

83
Q

Which Ksp values indicate a soluble, slightly soluble and insoluble compounds?

A
  • insoluble: Ksp &laquo_space;1
  • slightly soluble: (10^-5) < Ksp < (10^-2)
  • soluble: Ksp > (10^-2)
84
Q

Q vs Ksp

A

Q = Ksp: equilibrium solution is saturated
Q < Ksp not at equilibrium and not saturated Q > Ksp not at equilibrium and supersaturated
*If a solution is not at equilibrium, the reaction will shift either left or right until it is at equilibrium.

85
Q

Oxidation

A

becomes more positive by losing electrons

86
Q

Reduction

A

becomes more negative by gaining electrons

87
Q

Oxidizing agent

A

the compound being reduced

88
Q

Reducing agent

A

the compound being oxidized

89
Q

Galvanic cell

A

each half reaction is separated into half-cells connected by a wire and salt bridge

90
Q

Salt bridge

A

allows ion migration without free mixing of solutions

91
Q

Anode

A
  • negative
  • oxidation happens here
  • produces electrons
92
Q

Cathode

A
  • positive
  • reduction happens here
  • uses electrons
93
Q

Shorthand cell notation

A

Anode | Oxidation Reagents || Reduction Reagents | Cathode
double line=salt bridge
single lines=phase boundaries

94
Q

Standard electrical potential (Ecell)

A

-electrochemical cell potential under standard state conditions
-Electrical potential energy is converted to other forms of energy as the electrons flow through the external circuit
-The amount of potential energy depends on the two half reactions:
Potential = electrical potential energy difference

95
Q

Positive reduction potentials

A

Species that are reduced more easily than H3O+have more positive reduction potentials

  • Strong oxidizing agents
  • occurs at cathode as a reduction
96
Q

Standard Hydrogen electrode

A

2H30+ +2e = H2 + H2O

Standard Reduction Potential(Ecell)= 0V

97
Q

Negative reduction potentials

A
  • occurs at anode as oxidation

* strong reducing agent

98
Q

Standard cell potential

A

the difference between two standard reduction potentials

99
Q

Standard cell potential depends on:

A
  • concentrations
  • pressures
  • *** DOES NOT DEPEND ON STOICHIOMETRY, when multiplied by an integer, it remains unchanged
100
Q

Fuel cells

A
  • batteries in which reactants are constantly being added

- Anode and Cathode both Pt coated metal

101
Q

Electroplating

A

The process of depositing one metal on top of another, done usually for cosmetic reasons

102
Q

Battery

A

a galvanic cell that generates electrical current to power a practical device

103
Q

Alkaline Dry Cell Battery

A

oxidation of Zn under basic conditions with MnO2 and a passive graphite electrode

104
Q

Lead storage batteries

A

provides power in automobiles, Pb anode and PbO2 cathode

105
Q

Nickel-Cadmium Battery

A

Cd anode and NiO2 cathode

  • Rechargeable because the nickel and cadmium hydroxides adhere tightly to the electrodes.
  • Drawback=based on toxic heavy metal, also adds weight
106
Q

Lithium Ion Battery

A
  • Li ions are intercalated in both the anode and cathode materials
  • Li ions move from the anode to the cathode during discharge, and back during recharge.
  • Light weight, high power batteries
107
Q

Corrosion

A

the natural redox process that returns refined metals to their more stable oxides
ex) rust where Fe2+ is oxidized to Fe2O3

108
Q

How can you prevent or slow the process of corrosion?

A

“Galvanizing” an iron object with zinc metal prevents corrosion of the iron because the Zn is preferentiallty oxidized.

109
Q

Endothermic

A

delta E = negative
w = negative
q = negative
delta H = positive

110
Q

Exothermic

A

delta E = positive
w = positive
q = positive
delta H = negative

111
Q

A reaction is spontaneous, nonspontaneous or at equilibrium if delta G is:

A 
spontaneous = -
nonspontaneous = +
equilibrium = 0
112
Q

Which value of specific heat capacity will have the greatest temperature change?

A

lowest specific heat capacity value

113
Q
  1. Which of the following statements is true of a dynamic equilibrium process:

a) The rate of the forward reaction is equal to the rate of the reverse reaction
b) The concentrations (or pressures) of the products equal to that of the reactants
c) Neither the products or the reactants are completely used up
d) Two of the above statements about dynamic equilibrium are true
e) All three of the above statements about dynamic equilibrium are

A

a and c are both true

114
Q 
For the exothermic reaction: 
S (s) + 3/2 O2 (g) 	     SO3 (g)    

Which of the following will result in the formation of more products for this reaction?

(a) increase the pressure of the container
(b) add more SO3 to the reaction mixture
(c) add a catalyst
(d) heat the reaction mixture
(e) add N2 gas
A

a

115
Q
  1. Suppose we have a solution of AgBr at equilibrium. What will be the effect on this equilibrium when solid of Ag2CO3 is added:

a) Precipitate AgBr will form
b) More of the AgBr will dissolve
c) The pH of the solution will decrease
d) There will be no effect on equilibrium of the AgBr, since Ag2CO3 is a solid
e) The reaction mixture may explode…

A

a

116
Q

When the equilibrium constant is greater than 1, which side is favoured? What if it’s smaller than 1?

A

greater than 1= products are favoured

smaller than 1= reactants are favoured

CHEM 1310 - 2017 (2 decks)

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