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Thermodynamics Flashcards

1
Q

Heat

A

(q) amount of thermal energy in transit between a system and its surroundings due to a temperature difference

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2
Q

Thermal energy

A

energy that arises from kinetic energy of particles

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3
Q

Second Law of Thermodynamics

A

energy and matter naturally tend to disperse uniformly

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4
Q

Clausius Theorem

A

heat does not spontaneously pass from a colder object to a hotter body

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5
Q

thermal equilibrium

A

energy and temperature of both objects are equal and temperature change of both objects is 0 so no net flow of thermal energy

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6
Q

endothermic

A

system experiences an increase in thermal energy

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7
Q

exothermic

A

system experiences a decrease in thermal energy

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8
Q

Entropy

A

(S: J/K) amount of energy in a system that is unavailable to do work. Also a measure of the number of ways a system can be arranged under constraints such as temperature and pressure. extensive property

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9
Q

S_universe =

A

S_surroundings + S_system

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10
Q

Change in Entropy of Surroundings

A

Change S_universe = Change S_surroundings + Change S_system where for a reversible process: change in S= q_rev/T

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11
Q

spontaneous reaction

A

increases the entropy of the universe

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12
Q

non-spontaneous

A

decrease the entropy of the universe

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13
Q

change in entropy of the universe is greater than 0

A

spontaneous reaction; favored process

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14
Q

change in entropy of the universe is less than 0

A

nonspontaneous reaction; disfavored

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15
Q

change in entropy of the universe is equal to 0

A

at equilibrium; neither favored or disfavored

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16
Q

Heat capacity

A

(C: J/K) objects that are heated or cooled without a change in phase experience a change in temperature–physical and extensive property of matter

q=CΔT

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17
Q

molar heat capacity

A

(C_m: J/mol K) amount of heat required to change one mole of a substance one Kelvin

q=nC_mΔT

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18
Q

Specific Heat Capacity

A

(c: J/gK) amount of heat required to change one mole of a substance one Kelvin

q=mcΔT

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19
Q

heat capacity changes with

A

phase that the object or substance is given in

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20
Q

Molar Heat Capacity–> Specific Heat Capacity

A

c(X) = C(X) M(X)^-1

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21
Q

Specific Heat Capacity–> Molar Heat Capacity

A

C(X)=c(X) M(X)

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22
Q

Work

A

(w) form of energy transfer

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23
Q

Internal Energy

A

(U) energy of a system necessary too bring the system from its standard (internal) state to its present internal state of interest

  • accounts for the kinetic and potential energy microscopically. That is, it accounts for the kinetic energy of molecules (motion, vibrations, and rotations) as well as potential energy arising from chemical bonds and intermolecular forces.
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24
Q

Change in internal energy

A

ΔU=q+w

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25
Total Energy
accounts for internal energy of the system and the macroscopic forms of kinetic and potential energy including the kinetic energy of an entire system and the potential energy of an entire system
26
Total Energy Equation
E=U+T+V
27
Change in Internal Energy can be determined by
measuring the heat absorbed or released by the system as well as the work performed on or by the system
28
q>0
+ heat absorbed by system (endothermic) U increases
29
q<0
- heat released by system (exothermic) U decreases
30
w>0
+ work performed on the system U increases
31
w<0
- work performed by the system U decreases
32
Pressure-Volume Work
(isobaric process) work that occurs when the volume of a system changes against a constant external pressure
33
Pressure-Volume Work Equation
w=-P_extΔV (when the system is increasing in volume) w=P_extΔV (when the system is compressed ΔV becomes negative making the work energy quantity positive)
34
Constant-Volume Conditions
(isochoric process) change in volume is 0 and no pV-work is performed SO change in internal energy is directly related to the quantity of heat that was transferred (ΔU=qv+wv but wv=0 so ΔU=qv)
35
enthalpy
(H) describe the energy change under constant-pressure conditions H=U+V AND ΔH=qp where qp is change in heat under constant-pressure conditions
36
Bomb calorimetry
technique used to measure changes in internal energy of a system under constant-volume conditions
37
state functions
quantities that only depend on the thermodynamic state of a system (EX: mass, energy, entropy)
38
Path functions
quantities that depend on "path taken" between two thermodynamic states. Include heat and work
39
molar enthalpy of reaction or molar heat of reaction
(ΔrHm in kJ mol–1) enthalpy change per mole of reaction associated with a chemical reaction
40
Standard State
reference point used to simplify thermodynamic evaluations: denoted with a Plimsoll (⦵) or degree symbol (°)
41
Standard State Gases
hypothetical state a gas would have as a pure substance obeying the ideal gas equation at standard state pressure
42
liquids standard state
state of a pure substance as a liquid under standard state pressure
43
solids standard state
state of pure, crystalline substance under standard state pressure
44
solutes standard state
hypothetical state a solute would have in an ideal solution with a standard amount concentration c°, of exactly 1 mol L–1
45
Standard state of pressure
p°, is defined to be a pressure of 1 bar or 105 Pa.
46
standard enthalpy of formation
ΔfH° in kJ mol–1), also called the standard heat of formation, of a compound is the change in enthalpy during the formation of exactly 1 mole of a substance from its constituent elements in their reference state, with all substances in the standard state.
47
Standard heat of formation for substances in their pure, elemental state
0
48
Standard enthalpy of reaction
(ΔrH° or ΔrH⦵ in kJ mol–1) the difference in the total standard molar product and total standard molar reactant enthalpies of formation
49
Standard enthalpy of reaction represents:
Standard enthalpy of reaction represents reactants and products in the standard state where the reactants are unmixed (i.e. in pure, isolated form) and the products are unmixed. Therefore, the a standard enthalpy of a reaction compares the state of separated/isolated reactants to the state of separated/isolated products.
50
Equation for transforming enthalpy determined at 298.15 K to other temperatures
check website
51
Enthalpy and Internal Energy are related by
ΔH=ΔU+Δ(pV)
52
Constant Pressure Conditions gives the following relationship between enthalpy and internal energy
ΔH= ΔU+ pΔV
53
Ideal gas law
pV=nRT R=molar gas constant (8.134 J/molK)
54
Ideal gas law gives relationship between enthalpy and internal energy
Δ(pV)=Δ(nRT)
55
change in moles of gas
sum of products-sum of reactants
56
Calorimetry
technique used to determine the energy evolved from a chemical reaction or physical process by measuring changes in temperature.
57
Constant Pressure Calorimetry
measures enthalpy changes (Δ.H)
58
Where does heat exchange occur in calorimetry
within the system no work or heat exchange with the surroundings
59
Constant Volume Calorimetry
used to measure internal energy changes (ΔU)
60
Hess's Law
total enthalpy change of a complete set of reactions is independent of the sequence of steps taken

Chemistry (4 decks)

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