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unit 11 Flashcards

1
Q

thermochemistry

A

study of the relationships between chemistry and energy

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2
Q

heat

A

flow of energy caused by temp differences

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3
Q

thermal energy

A

energy associated with the temp of an object (type of kinetic energy due to motions of particles in a substance)

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4
Q

potential energy

A

energy associated with position or composition of an object

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5
Q

chemical energy

A

energy associated with relative positions of electrons and nuclei in atoms and molecules (a form of potential energy)

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6
Q

joule

A

SI unit of energy (also equal to 1kg · m2/s2). Often expressed as kilojoules (kJ)

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7
Q

thermodynamics

A

study of energy and its interconversions

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8
Q

1st law of thermodynamics

A

law of energy conservation-total energy of the universe is constant

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9
Q

internal energy of a system =

A

KE + PE of all particles of a system

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10
Q

internal energy change (regular) =

A

E(products) - E(reactants)

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11
Q

what does internal energy depend on?

A

the state of the system, not how the system arrived at that state

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12
Q

internal energy change (in a chemical system) =

A

E(products) - E(reactants)

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13
Q

E(system) =

A

E(surroundings)

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14
Q

when energy is released in a chemical reaction, E is (+/-)

A

negative

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15
Q

when energy is absorbed in a chemical reaction, E is (+/-)

A

positive

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16
Q

Which statement is true of the internal energy of a system and its surroundings during an energy exchange with a negative ∆Esys?

-The internal energy of the system increases and the internal energy of the surroundings decreases

-The internal energy of both the system and surroundings increases

-The internal energy of both the system and surroundings decreases

-The internal energy of the system decreases and the internal energy of the surroundings increases

A

-The internal energy of the system decreases and the internal energy of the surroundings increases

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17
Q

temperature definition

A

measure of the thermal energy in a sample of matter

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18
Q

thermal energy always flows from

A

matter at higher temps to matter at lower temps

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19
Q

thermal equillibrium

A

heat will be transferred from a hotter object to a cooler object, until they both reach the same temperature

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20
Q

q =

A

C x (change in)T

q: heat
C: heat capacity

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21
Q

heat capacity

A

measure of the system’s ability to absorb thermal energy without undergoing a large change in temp

quantity of heat required to change the system temp by 1 degree C

extensive property (depends on amount of matter being heated)

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22
Q

units of heat capacity

A

J / C
(obtained by solving C = q x (change in T))

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23
Q

specific heat capacity (Cs)

A

measure of intrinsic capacity of a substance to absorb heat

amount of heat needed to raise the temp of 1g added of the substance by 1 degree C (J/g x degree C)

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24
Q

equation relating heat added to amount of substance and temperature increase

A

q = m x Cs x (change in T)

q: amount of heat (j)
m: mass (g)
Cs: specific heat capacity (J/g x degrees C)
T: temp change (degrees C)

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25
enthalpy (H)
measures heat exchanged under conditions of constant pressure for a system and is a state function
26
(change in H) measures
heat exchanged under conditions of constant pressure
27
endothermic reaction
Positive ∆H - heat (thermal energy) flows into the system from the surroundings
28
exothermic reaction
Negative ∆H – the system releases heat to the surroundings
29
internal energy of a system =
kinetic energy + potential energy
30
what is the source in an exothermic chemical reaction?
potential energy
31
Under normal circumstances, chemical PE arises from
electrostatic attractions between protons and electrons of the atoms and molecules in the system (chemical bonds)
32
what happens to bonds in an exothermic reaction?
bonds break and new bonds form nuclei and electrons reorganize into a system with lower PE as the molecules rearrange, their PE converts into thermal energy, the heat emitted in the reaction
33
what happens to bonds in an endothermic reaction?
as some bonds break and form, nuclei and electrons reorganize into an arrangement with a higher PE absorbs energy in the process
34
enthalpy of reaction or heat of reaction (enthalpy change for a chemical reaction, ∆Hrxn)
∆Hrxn = Hproducts - Hreactants
35
Thermochemical equation
balanced chemical equations that show the associated ∆H
36
how is enthalpy an extensive property?
magnitude of ∆H is proportional to the amount of reactant consumed in the process of the reaction. Although chemical equations are usually written with whole-number coefficients, thermochemical equations sometimes use fractions.
37
relationship between ∆H in a reaction and its reverse reaction
equal in magnitude but opposite in sign to ∆H for the reverse reaction
38
what does ∆H for a reaction depend on?
states of reactants and products so it is important to specify states of the products and reactants
39
equation for heat generated or absorbed during the reaction
qsoln = msoln x Cs, soln x ∆T
40
qrxn and qsoln relationship
qrxn = -qsoln
41
equation of heat of reaction for a specific amount of reactants that reacted
qrxn = qp = ∆Hrxn
42
how many quantitative relationships are there between a chemical reaction and ∆Hrxn? what are they?
three. If a chemical equation is multiplied by a factor, ∆Hrxn is multiplied by the same factor. If a chemical equation is reversed, ∆Hrxn changes sign Hess's law
43
Hess's law
If a chemical equation can be expressed as the sum of a series of steps, then ∆Hrxn for the overall equation is the sum of the ∆Hrxn for each step
44
how can ∆Hrxn be determined?
experimentally using calorimetry or inferentially through Hess’s law. It can also be determined from standard enthalpies of formation
45
how many enthalpy relative standards are there?
three
46
standard state of enthalpy relative
-Gas: the pure gas at 1 atm pressure -Liquid or solid: pure substance in its most stable form at 1 atm pressure and the temperature of interest (usually 25o C) -Substance in solution: a 1M concentration
47
Standard enthalpy change (∆Ho)
∆H for a process when all reactants and products are in their standard states
48
Standard enthalpy (heat) of formation (∆Hof): for a pure compound
∆H when one mole of the compound forms from its constituent elements in their standard states
49
Standard enthalpy (heat) of formation (∆Hof): For a pure element in its standard state
∆Hof = 0. Allows us to measure ∆H relative to those of pure elements in their standard states. It is measured in kJ/mol
50
how can we calculate the standard enthalpy change for a reaction?
Hess's law and tabulated ∆Hof values
51
equation to use for enthalpies of formation to calculate enthalpies of reaction
∆Horxn = Σ np ∆Hof (products) - Σ nr ∆Hof (reactants) np = stoichiometric coefficients of products nr = stoichiometric coefficients of reactants Σ = sum of
52
entropy
The driving force behind physical and chemical changes is related to the dispersion (spreading out) of energy
53
goal of thermodynamics
predict spontaneity
54
Spontaneous process (thermodynamically favored) definition
The direction in which and the extent to which a process proceeds occurs without ongoing outside intervention
55
A chemical system can proceed in the direction of lowest enthalpy and could be a candidate for a chemical potential. However, if this were the case:
All exothermic reactions would be spontaneous All endothermic reactions would be nonspontaneous
56
does entropy increase or decrease in endothermic reactions?
increases
57
Does a state in which a given amount of energy is more highly dispersed (or more highly randomized) have more or less entropy than a state in which the same energy is more highly concentrated.
more
58
second law of thermodynamics
Processes that increase entropy are thermodynamically favored (spontaneous) Processes that decrease entropy are not thermodynamically favored (nonspontaneous)
59
entropy equation
∆S = Sfinal - Sinitial
60
entropy determines
the direction of physical and chemical change (proceeds in a direction that increases entropy) In thermodynamically favored processes, ∆S is positive. In non-thermodynamically favored processes ∆S is negative.
61
Entropy of matter increases as
it changes state from a solid to a liquid and from a liquid to a gas
62
Energy in a molecular solid consists largely of
the vibrations between its molecules
63
As a vapor, the energy in the substance can take the form of
straight-line motions (translational energy) and rotations of the molecules (rotational energy)
64
are vapors or solids more dispersed?
vapors
65
Gibbs free energy
maximum amount of energy available from any chemical reaction
66
what 2 forces drive chemical reactions?
Enthalpy (∆Ho) which represents the change in internal potential energy of the atoms A drive to increase the entropy (∆So) of the system
67
equation for Gibbs free energy change
∆G = ∆H - T∆S H is enthalpy T is kelvin temperature S is entropy
68
Gibbs free energy change and thermodynamic favorability relation
high ∆G, low thermodynamic favorability
69
Gibbs free energy is also called _ and is analogous to
chemical potential, mechanical potential
70
chemical systems tend toward
lower potential energy lower Gibbs free energy
71
A decrease in Gibbs free energy (∆G < 0) corresponds to
thermodynamically favored process
72
an increase in Gibbs free energy (∆G > 0) corresponds to
non-thermodynamically favored process
73
Gibbs free energy: ∆H negative, ∆S positive
-Reaction is exothermic (∆H < 0) -change in entropy is positive (∆S > 0) -Change in free energy is negative at all temperatures -reaction is favored at all temperatures
74
Gibbs free energy: ∆H positive, ∆S negative
-Reaction is endothermic (∆H > 0) -change in entropy is negative (∆S < 0) -Change in free energy is positive at all temperatures -reaction is not favored at all temperatures
75
Gibbs free energy: ∆H negative, ∆S negative
-Reaction is exothermic (∆H < 0) -change in entropy is negative (∆S < 0) -Sign of the change in free energy depends on temperature -the reaction is spontaneous at low temperature -nonspontaneous at high temperature
76
gibbs free energy: ∆H positive, ∆S positive
-Reaction is endothermic (∆H > 0) -change in entropy is positive (∆S > 0) -Sign of the change in free energy depends on temperature -reaction is nonspontaneous at low temperature and -spontaneous at high temperature
77
does spontaneity depend on temp when ∆H and ∆S have opposite signs?
no
78
does spontaneity depend on temp when ∆H and ∆S have the same signs?
yes
79
Standard entropy change of a reaction ((∆Sorxn) definition
change in entropy for a process in which all reactants and products are in their standard states
80
Standard entropy change of a reaction ((∆Sorxn) formula
∆Sorxn = Soproducts - Soreactants
81
how can you define a relative zero for enthalpy?
assign a value of zero to the standard enthalpy of formation (∆Hof) for an element in its standard state
82
3rd law of thermodynamics
The entropy of a perfect crystal at absolute zero (0 K) is zero
83
how is entropy an extensive property?
depends on the amount of the substance
84
units of standard molar entropy
joules / mol ∙ K (J / mol ∙ K)
85
rank gases, liquids, and solids from most to least standard entropy
gas liquid solid
86
entropy and molar mass relation
entropy increases with increasing molar mass (as long as the substances being compared are in the same state)
87
allotropes
In the same physical state, some elements can exist in two or more forms
88
entropy and molecular complexity relation
For a given state of matter, entropy generally increases with increasing molecular complexity
89
dissolution and entropy relation
dissolution of a crystalline solid usually results in increased entropy. The energy of the solid becomes dispersed throughout the solution
90
how do you calculate ∆Sorxn?
subtract the standard entropies of the reactants multiplied by their stoichiometric coefficients from the standard entropies of the products multiplied by their stoichiometric coefficients **standard entropies are always nonzero at 25o C
91
how can you calculate ∆Gorxn?
subtracting the tabulated free energies of formation of the reactants from the free energies of formation of the products
92
Free energies of formation (∆Gof)
change in free energy when 1 mol of a compound in its standard state forms from its constituent elements in their standard states. The free energy of formation of pure elements in their standard states is zero
93
Free energies of formation (∆Gof) formula
∆Gorxn = Σ np ∆Gf (products) - Σ nr ∆Gf (reactants)
94
how are the standard free energy change (∆Gorxn) and the equilibrium constant (K) related?
-K becomes larger as ∆Gorxn becomes more negative -If reactants undergo a large negative ∆Gorxn as they become products, then K is large with products strongly favored at equilibrium -If reactants undergo a large positive ∆Gorxn as they become products, then K is small with reactants strongly favored at equilibrium
95
equation for the relationship between ∆Gorxn and K
ΔGorxn = - RT ln K When K < 1, ln K is negative and ΔGorxn is positive. When K > 1, ln K is positive and ΔGorxn is negative. When K = 1, ln K is zero and ΔGorxn is zero. The reaction is at equilibrium.

ap chem (11 decks)

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