Thermochemistry Flashcards

1
Q

The law of conservation of energy

A

Energy can neither be created nor destroyed, but that all thermal, chemical, potential, and kinetic energies are interconvertible

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2
Q

Isolated Systems

A

no exchange of energy/matter with the environments (bomb calorimeters)

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3
Q

Closed Systems

A

can exchange energy but not matter with the environment

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4
Q

Open Systems

A

can exchange both energy and matter with the environment

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5
Q

Isothermal

A

temperature remains constant

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6
Q

Adiabatic

A

no heat exchange occurs

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7
Q

Isobaric

A

pressure of a system remains constant

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8
Q

Isovolumetric

A

aka isochoric

volume remains constant

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9
Q

Heat

A

the transfer of thermal energy from one object to another

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10
Q

Endothermic

A

Reactions that absorb thermal energy

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11
Q

Exothermic

A

Reactions that release thermal energy

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12
Q

Constant Volume and Constant Pressure Calorimetry

A

Used to indicate conditions under which the heat flow is measured

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13
Q

Equation to calculate heat absorbed or released

A

q=mcΔT

m= mass
c= specific heat
ΔT= change in temperature
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14
Q

State Functions

A

described by the macroscopic properties of the system, and the magnitudes of these properties depend on only the initial and final states of the system, not the path.
Include: pressure, density temperature, volume, enthalpy, internal energy, free energy, and entropy

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15
Q

Enthalpy

A

(H) used to express the heat changes at constant pressure

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16
Q

Standard heat of formation

A

the enthalpy change that would occur if one mole of a compound was formed directly from its elements in their standard states

17
Q

Standard heat of reaction

A

the hypothetical enthalpy change that would occur if one mole of a compound was formed directly from its elements

18
Q

Hess’s Law

A

states that enthalpies of reactions are additive

the reverse of any reaction has an enthalpy of the same magnitude as that of the forward reaction (sign is opposite)

19
Q

Bond Dissociation Energy

A

an average of the energy required to beak a particular type of bond in one mole of gaseous molecules

20
Q

Bond Enthalpy

A

the standard heat of reaction can be calculated using the values of bond dissociation energies of particular bonds

21
Q

Entropy

A

(S) the measure of the distribution of energy throughout a system

22
Q

Gibbs free energy

A

combines the two factors that affect the spontaneity of a reaction: changes in enthalpy (ΔH) and changes in entropy (ΔS)

ΔG= ΔH - TΔS

23
Q

If ΔG is negative

A

the reaction is spontaneous

24
Q

If ΔG is positive

A

the reaction is nonspontaneous

25
If ΔG is zero
the reaction is in equilibrium (ΔH=TΔS)
26
If ΔH is negative and ΔS is positive
reaction is spontaneous at all temperatures
27
If ΔH is positive and ΔS is negative
reaction is nonspontaneous at all temperatures
28
If ΔH is positive and ΔS is positive
spontaneous at very high temperatures
29
If ΔH is negative and ΔS is negative
spontaneous at very low temperatures
30
Reaction Quotient
(Q) once a reaction has started the standard state conditions no longer hold so we use Q to determine where we are in the reaction (before eqlbm, after eqlbm or at eqlbm)
31
Standard Conditions
298K 1 atm 1 M
32
Standard Temperature and Pressure
273K 1 atm ** used for ideal gas problems
33
Specific Heat
the specific heat required to raise the temperature of 1 gram of substance by 1 degree
34
Heat Absorbed or Released in a Process (Equation)
q=mcΔt m- mass c- specific heat t- temp
35
During Phase Changes we use
q=mL | L- latent heat
36
Entropy Equation
ΔS = Qrev/T
37
Equation for Gibbs Free Energy
ΔG= ΔH - TΔS
38
Standard Free Energy for the Reaction Equation
ΔGrxn = -RT lnKeq