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unit 11 Flashcards

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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
Q

enthalpy (H)

A

measures heat exchanged under conditions of constant pressure for a system and is a state function

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26
Q

(change in H) measures

A

heat exchanged under conditions of constant pressure

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27
Q

endothermic reaction

A

Positive ∆H - heat (thermal energy) flows into the system from the surroundings

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28
Q

exothermic reaction

A

Negative ∆H – the system releases heat to the surroundings

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29
Q

internal energy of a system =

A

kinetic energy + potential energy

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30
Q

what is the source in an exothermic chemical reaction?

A

potential energy

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31
Q

Under normal circumstances, chemical PE arises from

A

electrostatic attractions between protons and electrons of the atoms and molecules in the system (chemical bonds)

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32
Q

what happens to bonds in an exothermic reaction?

A

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

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33
Q

what happens to bonds in an endothermic reaction?

A

as some bonds break and form, nuclei and electrons reorganize into an arrangement with a higher PE

absorbs energy in the process

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34
Q

enthalpy of reaction or heat of reaction (enthalpy change for a chemical reaction, ∆Hrxn)

A

∆Hrxn = Hproducts - Hreactants

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35
Q

Thermochemical equation

A

balanced chemical equations that show the associated ∆H

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36
Q

how is enthalpy an extensive property?

A

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.

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37
Q

relationship between ∆H in a reaction and its reverse reaction

A

equal in magnitude but opposite in sign to ∆H for the reverse reaction

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38
Q

what does ∆H for a reaction depend on?

A

states of reactants and products so it is important to specify states of the products and reactants

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39
Q

equation for heat generated or absorbed during the reaction

A

qsoln = msoln x Cs, soln x ∆T

40
Q

qrxn and qsoln relationship

A

qrxn = -qsoln

41
Q

equation of heat of reaction for a specific amount of reactants that reacted

A

qrxn = qp = ∆Hrxn

42
Q

how many quantitative relationships are there between a chemical reaction and ∆Hrxn? what are they?

A

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
Q

Hess’s law

A

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
Q

how can ∆Hrxn be determined?

A

experimentally using calorimetry or inferentially through Hess’s law. It can also be determined from standard enthalpies of formation

45
Q

how many enthalpy relative standards are there?

A

three

46
Q

standard state of enthalpy relative

A

-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
Q

Standard enthalpy change (∆Ho)

A

∆H for a process when all reactants and products are in their standard states

48
Q

Standard enthalpy (heat) of formation (∆Hof): for a pure compound

A

∆H when one mole of the compound forms from its constituent elements in their standard states

49
Q

Standard enthalpy (heat) of formation (∆Hof): For a pure element in its standard state

A

∆Hof = 0. Allows us to measure ∆H relative to those of pure elements in their standard states. It is measured in kJ/mol

50
Q

how can we calculate the standard enthalpy change for a reaction?

A

Hess’s law and tabulated ∆Hof values

51
Q

equation to use for enthalpies of formation to calculate enthalpies of reaction

A

∆Horxn = Σ np ∆Hof (products) - Σ nr ∆Hof (reactants)

np = stoichiometric coefficients of products
nr = stoichiometric coefficients of reactants
Σ = sum of

52
Q

entropy

A

The driving force behind physical and chemical changes is related to the dispersion (spreading out) of energy

53
Q

goal of thermodynamics

A

predict spontaneity

54
Q

Spontaneous process (thermodynamically favored) definition

A

The direction in which and the extent to which a process proceeds

occurs without ongoing outside intervention

55
Q

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:

A

All exothermic reactions would be spontaneous

All endothermic reactions would be nonspontaneous

56
Q

does entropy increase or decrease in endothermic reactions?

A

increases

57
Q

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.

A

more

58
Q

second law of thermodynamics

A

Processes that increase entropy are thermodynamically favored (spontaneous)

Processes that decrease entropy are not thermodynamically favored (nonspontaneous)

59
Q

entropy equation

A

∆S = Sfinal - Sinitial

60
Q

entropy determines

A

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
Q

Entropy of matter increases as

A

it changes state from a solid to a liquid and from a liquid to a gas

62
Q

Energy in a molecular solid consists largely of

A

the vibrations between its molecules

63
Q

As a vapor, the energy in the substance can take the form of

A

straight-line motions (translational energy) and rotations of the molecules (rotational energy)

64
Q

are vapors or solids more dispersed?

A

vapors

65
Q

Gibbs free energy

A

maximum amount of energy available from any chemical reaction

66
Q

what 2 forces drive chemical reactions?

A

Enthalpy (∆Ho) which represents the change in internal potential energy of the atoms

A drive to increase the entropy (∆So) of the system

67
Q

equation for Gibbs free energy change

A

∆G = ∆H - T∆S

H is enthalpy
T is kelvin temperature
S is entropy

68
Q

Gibbs free energy change and thermodynamic favorability relation

A

high ∆G, low thermodynamic favorability

69
Q

Gibbs free energy is also called _ and is analogous to

A

chemical potential, mechanical potential

70
Q

chemical systems tend toward

A

lower potential energy
lower Gibbs free energy

71
Q

A decrease in Gibbs free energy (∆G < 0) corresponds to

A

thermodynamically favored process

72
Q

an increase in Gibbs free energy (∆G > 0) corresponds to

A

non-thermodynamically favored process

73
Q

Gibbs free energy: ∆H negative, ∆S positive

A

-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
Q

Gibbs free energy: ∆H positive, ∆S negative

A

-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
Q

Gibbs free energy: ∆H negative, ∆S negative

A

-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
Q

gibbs free energy: ∆H positive, ∆S positive

A

-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
Q

does spontaneity depend on temp when ∆H and ∆S have opposite signs?

A

no

78
Q

does spontaneity depend on temp when ∆H and ∆S have the same signs?

A

yes

79
Q

Standard entropy change of a reaction ((∆Sorxn) definition

A

change in entropy for a process in which all reactants and products are in their standard states

80
Q

Standard entropy change of a reaction ((∆Sorxn) formula

A

∆Sorxn = Soproducts - Soreactants

81
Q

how can you define a relative zero for enthalpy?

A

assign a value of zero to the standard enthalpy of formation (∆Hof) for an element in its standard state

82
Q

3rd law of thermodynamics

A

The entropy of a perfect crystal at absolute zero (0 K) is zero

83
Q

how is entropy an extensive property?

A

depends on the amount of the substance

84
Q

units of standard molar entropy

A

joules / mol ∙ K (J / mol ∙ K)

85
Q

rank gases, liquids, and solids from most to least standard entropy

A

gas
liquid
solid

86
Q

entropy and molar mass relation

A

entropy increases with increasing molar mass (as long as the substances being compared are in the same state)

87
Q

allotropes

A

In the same physical state, some elements can exist in two or more forms

88
Q

entropy and molecular complexity relation

A

For a given state of matter, entropy generally increases with increasing
molecular complexity

89
Q

dissolution and entropy relation

A

dissolution of a crystalline solid usually results in increased entropy. The energy of the solid becomes dispersed throughout the solution

90
Q

how do you calculate ∆Sorxn?

A

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
Q

how can you calculate ∆Gorxn?

A

subtracting the tabulated free energies of formation of the reactants from the free energies of formation of the products

92
Q

Free energies of formation (∆Gof)

A

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
Q

Free energies of formation (∆Gof) formula

A

∆Gorxn = Σ np ∆Gf (products) - Σ nr ∆Gf (reactants)

94
Q

how are the standard free energy change (∆Gorxn) and the equilibrium constant (K) related?

A

-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
Q

equation for the relationship between ∆Gorxn and K

A

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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