# Thermodynamics Flashcards

## From the definition of the various state functions to the properties of complex thermodynamic systems and processes, use these cards to master the topic of thermodynamics as tested in most introductory undergrad chemistry courses and on the AP Chemistry exam.

1
Q

Define and give an example of:

a thermodynamic system

A

A thermodynamic system is a macroscopic body which is engaged in mass and/or energy exchange with its surroundings.

Ex: The classic example of a thermodynamic system is a piston filled with gas.

2
Q

Define and give an example of:

an open thermodynamic system

A

An open thermodynamic system can exchange both mass and energy with the environment.

Ex: A bottle of gas with no lid is an open thermodynamic system.

3
Q

Define and give an example of:

a closed thermodynamic system

A

A closed thermodynamic system cannot exchange mass with the environment, but can exchange energy.

Ex: A bottle of gas with the lid securely on is a closed thermodynamic system.

4
Q

Define and give an example of:

an isolated thermodynamic system

A

An isolated thermodynamic system can not exchange either energy or mass with the environment.

Ex: A closed, double-walled (insulated) container which is temperature-independent of its environment is an isolated system.

5
Q

Define “the surroundings” in a thermodynamic problem

A

The surroundings (also known as the environment) are everything capable of exchanging mass and/or energy with the system.

6
Q

Define:

a state function

A

A state function is any property of a thermodynamic system that depends only on comparing the characteristics of the system at that moment compared to a prior moment.

Since state functions are calculated based on current vs past properties only, their values do not depend on the path by which the current state was achieved; they are path-independent.

7
Q

If X is a state function, what is the change in X between a system’s values X1 and X2?

A

ΔX = X2 - X1

Since X is a state function, the path by which it gets from state 1 to 2 is irrelevant, the change in X depends only on the starting and finishing states.

8
Q

Define:

enthalpy, H

A

Enthalpy is a measure of the heat contained in a system.

Enthalpy’s absolute value cannot be directly measured, so the change in enthalpy’s value, ΔH, is measured instead.

9
Q

What are the properties of an exothermic reaction?

A

An exothermic reaction is any reaction whose products have a lower enthalpy than the reactants. Therefore, ΔH < 0, and heat is lost from the system to the environment.

(Exo = exit)

10
Q

What are the properties of an endothermic reaction?

A

An endothermic reaction is any reaction whose products have a higher enthalpy than the reactants. Therefore, ΔH > 0, and heat is absorbed by the system from the environment.

(Endo = into)

11
Q

Define:

a material’s standard enthalpy of formation, ΔHof

A

The standard enthalpy of formation (ΔHof) is the enthalpy change of the formation reaction for the material from its fundamental elements under standard conditions.

Ex: The enthalpy of formation for NaCl is -411.12 kJ mol−1 and follows from the general equation:

12
Q

What is the standard enthalpy of formation, ΔHof, of oxygen gas, O2?

A

Zero.

By definition, the enthalpy of formation for any material in its standard state is zero.

13
Q

What are standard conditions?

A

Standard conditions for a reaction require that:

• Pressure = 1 atm
• Temperature = 25º C = 298 K
14
Q

If the chemical reaction

(1) 2A ⇒ C ΔH1

can be broken down into the sub-reactions

(2) 2A ⇒ B ΔH2
(3) B ⇒ C ΔH3

what does Hess’s Law tell you about the overall enthalpy change ΔH1?

A

ΔH1 = ΔH2 + ΔH3

Hess’s Law simply states that the enthalpy of a reaction can be calculated by adding together the enthalpies of a chain of sub-reactions which add up to the overall reaction.

Although most commonly applied to enthalpy, Hess’s Law applies to all state functions.

15
Q

What is the enthalpy change when 2 moles of CH4 are formed, according to the following reactions?

rxn1: 2H2(g)⇒4H(g)
ΔH1= -870 kJ/mol

rxn2: C(s) + 4H(g)⇒CH4(g)
ΔH2= +794 kJ/mol

A

-152 kJ

Adding the reactions together yields the formation reaction of CH4:

C(s) + 4H(g) + 2H2(g) ⇒
CH4(g) + 4H(g)
Canceling common terms leaves:
C(s) + 2H2(g) ⇒CH4(g)

To complete the calculation, combine the reactions’ enthalpies in the same way the reactions were combined.

ΔHrxn = ΔH1 + ΔH2
-870 + 794 = -76kJ/mol

Finally, multiply by the number of moles (2) to get the final answer.

16
Q

How can the enthalpy change of a reaction be calculated from the enthalpies of formation of the reactants and products?

A

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

Sum the enthalpies of formation of the products, and subtract the sum of the enthalpies of formation of the reactants.

17
Q

Is this reaction endothermic or exothermic?

CH4+ 2O2 ⇒ CO2+ 2H20

• ΔHof (CH4) = -75 kJ/mol
• ΔHof (CO2) = -394 kJ/mol
• ΔHof (H2O) = -286 kJ/mol
A

The reaction is exothermic.

∆H°rxn= Σ∆Hof(products)-Σ∆Hof(reactants)
= [CO2 + 2*H2O] - [CH4 + O2]

= [-394 kJ/mol + 2(-286 kJ/mol)]
- [-75 kJ/mol + 2(0 kJ)]

= -891 kJ/mol

18
Q

Define:

bond enthalpy

A

Bond enthalpy is the energy absorbed or released when a particular chemical bond is broken.

Most chemical bonds are stabilizing, so most bond-breaking reactions are endothermic, and most bond enthalpies are positive.

19
Q

How can the enthalpy of a reaction be calculated from the bond enthalpies of the reactants and products?

A

∆Hrxn = Σ∆H(bonds broken) - Σ∆H(bonds formed)

The reaction’s enthalpy change is identical to the energy needed to break all the bonds in the reactants, minus the energy released when the bonds in the products form.

20
Q

What is the overall enthalpy change of this reaction?

CH4+ 2O2 ⇒ CO2+ 2H20

• ΔH (C-H) = 411 kJ/mol
• ΔH (O=O) = 494 kJ/mol
• ΔH (C=O) = 799 kJ/mol
• ΔH (O-H) = 463 kJ/mol
A

-818 kJ/mol

∆Hrxn = Σ∆H(bonds broken) - Σ∆H(bonds formed)
= [4*(C-H) + 2*(O=O)] - [2*(C=O) + 2*2*(H-O)]

= [(4 * 411) + (2 * 494)]
- [(2 * 799) + (4 * 463)] kJ/mol

= -818 kJ/mol

21
Q

Define:

specific heat, c

A

Specific heat is a characteristic property of a material, and is the amount of heat which must be added to raise 1 g of the substance by 1ºC.

The higher the value of c, the more heat it takes to raise the substance’s temperature.

22
Q

What is the formula for necessary quantity of heat in a specific heat problem?

A

q = mcΔT

Where q is the necessary heat, m is the mass of substance present, c is the substance’s specific heat, and ΔT is the desired temperature change.

23
Q

What is the specific heat of water?

A

4.184 J/g*K

This is a value that you should have memorized. It means that it takes 4.184 J to raise 1 g of water by 1ºK.

This value is equal to 1 cal/g*K.

24
Q

8 J of heat is applied to 1 g of both iron and water at 25ºC. Which changes temperature more?

The specific heat of water is 4.184 J/g-K.

The specific heat of iron is 0.46 J/g-K.

A

The iron changes its temperature more.

Applying the equation q = mcΔT to both cases and solving for ΔT reveals that the iron will change temperature by about 20 degrees (final T = 45ºC), while the water will only increase by 2 degrees (final T = 27ºC).

The higher a material’s specific heat, the less responsive its temperature is to heat flow.

25
Q

What does a calorimeter measure?

A

A calorimeter measures the amount of heat given off by a particular chemical reaction or process.

There are many different styles of calorimeter, but for most chemistry courses, including the AP Chemistry exam, you should focus on the fact that they all measure heat generated via a system’s temperature change, using the equation

q = mcΔT

26
Q

Why doesn’t the specific heat equation q = mcΔT apply during a phase change?

A

During a phase change, temperature stays constant as heat is added. The added heat causes the material to go through the phase change, rather than increasing its temperature.

The amount of heat needed to make a material change its phase is known as the latent heat of that phase change.

27
Q

What is the equation for calculating the heat of a phase change?

A

q = mΔHL

where q = necessary heat, m = mass of the substance present, and ΔHL is the latent heat of the phase change.

The higher the ΔHL, the more heat it takes to force the substance to go through the phase change.

28
Q

The curve below represents a sample’s temperature vs. heat added. What phases (solid, liquid, and/or gas) are present at each labeled point on the plot?

A

a. solid
b. both solid and liquid
c. liquid
d. both liquid and gas
e. gas

29
Q

The curve below represents a sample’s temperature vs. heat added. What heat (q) formula would need to be applied, in order to calculate heat added to the system at each labeled point on the plot?

A

a. q=mcΔT
b. q=mΔHfus
c. q=mcΔT
d. q=mΔHvap
e. q=mcΔT

30
Q

Define:

entropy of a system

A

Entropy is a macroscopic property of a system, representing the number of possible ways the atoms or molecules of the system can arrange themselves.

Note: Colloquially, entropy is said to represent the ‘possibility for disorder’ of a system, and this definition is effective enough for most entropy questions on the AP Chem exam.

31
Q

What does increasing entropy say about a system’s properties?

A

As a system’s entropy increases, it becomes more disordered.

Systems spontaneously tend towards arrangements with higher entropy.

32
Q

Arrange the relative entropy levels of the 3 phases of matter:

Ssolid, Sgas, Sliquid

A

Ssolid < Sliquid < Sgas

Since entropy is a measure of disorder, the more ordered a system, the lower its entropy. Solid matter, with its regularly-repeating units and fairly well-defined locations for the atoms, is therefore low in entropy, while gases, with their continuous, random motion, have the highest entropy values.

33
Q

What is the entropy change for this chemical reaction?

2H2(g) + O2(g) ⇒ 2H2O(l)

A

ΔS < 0

Gases have more entropy than liquids; hence, in any chemical reaction, the side with more moles of gas will have higher entropy.

In general, reactions with fewer moles of products than reactants will have lower final entropy.

34
Q

What are the enthalpy and entropy changes for this reaction?

H2O (l) ⇒ H2O (g)

A

ΔHrxn > 0
ΔSrxn > 0

The reaction is endothermic, since heat must be added to vaporize the water.
Since the reaction creates gas, it represents an increase in entropy.

35
Q

What are the enthalpy and entropy changes for this reaction?

CO2 (g) ⇒ CO2 (s)

A

ΔHrxn < 0
ΔSrxn < 0

The reaction is exothermic, since heat must be removed to deposit the CO2.
Since the reaction results in a net decrease of gas, it represents an decrease in entropy.

36
Q

Define and give the formula for calculating:

Gibbs’ Free Energy, ΔG

A

ΔG is a measure of the work which can be extracted from a thermodynamic system.

ΔG = ΔH - TΔS

It also is used to measure the spontaneity of a system; chemical systems will always tend to move in a direction of decreasing ΔG.

A handy mnemonic for remembering the formula for ΔG is: ‘Get Higher Test Scores’.

37
Q

What does a negative value for ΔG imply about a chemical reaction?

A

If ΔG < 0 for a chemical reaction, then the forward reaction is spontaneous, favoring the creation of more products.

38
Q

What does a positive value for ΔG imply about a chemical reaction?

A

If ΔG > 0 for a chemical reaction, then the forward reaction is non-spontaneous, and the reverse reaction is spontaneous, favoring the creation of more reactants.

39
Q

Define:

an exergonic reaction

A

An exergonic reaction is any reaction for which ΔG < 0; hence all exergonic reactions are spontaneous.

This is very similar to the definition of exothermic, for which ΔH < 0. “Thermic” refers to Enthalpy, “gonic” to Gibbs’ Free Energy.

(Exerg = exit G)

40
Q

Define:

an endergonic reaction

A

An endergonic reaction is any reaction for which ΔG > 0; hence all endergonic reactions are non-spontaneous.

This is very similar to the definition of endothermic, for which ΔH > 0. “Thermic” refers to Enthalpy, “gonic” to Gibbs’ Free Energy.

(Enderg = enter G)

41
Q

Describe the spontaneity of a reaction where:

ΔHrxn > 0
ΔSrxn < 0

A

This reaction will always be non-spontaneous.

Remember, ΔG = ΔH - TΔS. Since T is in Kelvin and will always be positive, then both ΔH and -TΔS are positive. The quantity ΔG will always be positive then, so the reaction is endergonic at all temperatures.

42
Q

Describe the spontaneity of a reaction where:

ΔHrxn < 0
ΔSrxn < 0

A

This reaction will be spontaneous at low temperatures, and non-spontaneous at sufficiently high temperatures.

Remember, ΔG = ΔH - TΔS. When T is low (near zero Kelvin), the second term can be ignored, and ΔG will be negative due to ΔH<0. But when T becomes large enough, the positive sign of the -TΔS term dominates, and ΔG becomes positive.

43
Q

Describe the spontaneity of a reaction when:

ΔHrxn > 0
ΔSrxn > 0

A

This reaction will be spontaneous at high temperatures and non-spontaneous at sufficiently low temperatures.

Remember, ΔG = ΔH - TΔS. When T is low (near zero Kelvin), the second term can be ignored, and ΔG will be positive due to ΔH>0. But when T becomes large enough, the negative sign of the -TΔS term dominates, and ΔG becomes negative.

44
Q

Describe the spontaneity of a reaction where:

ΔHrxn < 0
ΔSrxn > 0

A

This reaction will always be spontaneous.

Remember, ΔG = ΔH - TΔS. Since T is in Kelvin and will always be positive, both ΔH and -TΔS are negative. The quantity ΔG will always be negative, so the reaction is exergonic at all temperatures.

45
Q

What is the difference between ΔG and ΔGº?

A

ΔG describes the change in Gibbs’ Free Energy for a chemical system at a particular pressure and temperature, which must be given.

ΔGº describes the change in Gibbs’ Free Energy for a chemical system at standard conditions (1 atm, 298 K, 1 M in all concentrations).

Typically, different reactions’ ΔGº values will be compared, since this gives a common point of reference between them.

46
Q

What is the relationship between a thermodynamic system’s temperature and its internal energy?

A

Temperature and internal energy are equivalent concepts.

Ex: If a system’s absolute temperature doubles, so does the amount of energy that it contains.

This is the foundation of the zeroth law of thermodynamics, which states that if systems A and C are both in thermal equilibrium with system B, they are also in equilibrium with each other; i.e. they are at the same temperature.

47
Q

Give the relationship between a fluid’s temperature and the internal energy (average kinetic energy) of the molecules in the fluid.

A

U = KEavg = (3/2)nkT

k is Boltzmann’s constant(1.38x10-23 J/K);
T is the absolute temperature (in Kelvin);
n is the number of moles of fluid molecules (in mol)

48
Q

Define:

work (in thermodynamics)

A

Work is the flow of energy between a system and its surroundings in the form of changing pressure and volume of the system.

In thermodynamics, work is signified by the symbol w.

49
Q

Define:

heat (in thermodynamics)

A

Heat is the flow of energy between a system and its surroundings in any form other than work.

In thermodynamics, heat is signified by the symbol q.

50
Q

What is the First Law of Thermodynamics?

A

The First Law of Thermodynamics states that heat and work are the only ways in which energy can flow, and energy is always conserved, hence energy change can be calculated:

ΔE = Δq + Δw

51
Q

Explain the direction of work being done and the sign of ΔE during the compression of a thermodynamic system.

A

During a compression, work is being done by the environment, and on the system.

Since energy is flowing into the system due to the work being done, ΔE > 0.

52
Q

Explain the direction of work being done and the sign of ΔE during the expansion of a thermodynamic system.

A

During an expansion, work is being done by the system, and on the environment.

Since energy is flowing out of the system due to the work being done, ΔE < 0.

53
Q

If a thermodynamic system is surrounded by an environment which is at a higher temperature, what is the direction of heat flow? What does this mean for the sign of ΔE?

A

When the environment is at a higher temperature than the system, heat flows from the environment into the system, and Δq > 0.

This direction of heat flow leads to energy being added to the system, and ΔE > 0.

54
Q

If a thermodynamic system is surrounded by an environment which is at a lower temperature, what is the direction of heat flow? What does this mean for the sign of ΔE?

A

When the environment is at a lower temperature than the system, heat flows from the system to the environment, and Δq < 0.

This direction of heat flow leads to energy being removed from the system, and ΔE < 0.

55
Q

What are the characteristics of an adiabatic thermodynamic process?

A

An adiabatic process is any process where heat cannot flow. Hence, Δq = 0 and ΔE = Δw; all energy change is due to work being done on or by the system.

Adiabatic processes are processes that happen in either heat-insulated systems, or that happen so quickly that heat cannot flow between system and environment.

56
Q

What happens to the system’s temperature during an adiabatic compression?

A

The system’s temperature must increase during an adiabatic compression.

Remember, Δq = 0 for all adiabatic processes. Furthermore, during any compression, work is being done on the system, raising ΔE. Since T and ΔE are proportional, the system must gain temperature during an adiabatic compression.

57
Q

What happens to the system’s temperature during an adiabatic expansion?

A

The system’s temperature must decrease during an adiabatic expansion.

Remember, Δq = 0 for all adiabatic processes. Furthermore, during any expansion, the system is doing work, causing it to lose energy. Since T and ΔE are proportional, the system must lose temperature during an adiabatic expansion.

58
Q

What are the characteristics of an isothermal thermodynamic process?

A

An isothermal process is any process where temperature is held constant. Hence, ΔE = 0, and Δq = -Δw. Any heat flow into or out of the system is compensated by work done by or on the system, respectively.

Isothermal processes are usually processes that happen slowly enough that the system’s temperature can constantly equilibrate.

59
Q

What is the direction of heat flow during an isothermal compression?

A

Heat flows out of the system during an isothermal compression.

Remember, during any compression, work is being done on the system, raising ΔE. Since an isothermal compression must have an overall ΔE = 0, Δq must be negative to compensate.

60
Q

What is the direction of heat flow during an isothermal expansion?

A

Heat flows into the system during an isothermal expansion.

Remember, during any expansion, the system is doing work, losing energy. Since any isothermal process must have an overall ΔE = 0, Δq must be positive to compensate.

61
Q

Define:

the Second Law of Thermodynamics

A

The Second Law of Thermodynamics states that for any thermodynamic process, the total entropy of the universe increases.

ΔSuniverse > 0
ΔSsystem+ ΔSsurroundings > 0

since the universe can be defined as a system and its surroundings. If the entropy of a system decreases, therefore, the entropy of its surroundings must increase by a greater amount in order for the universal law to still hold true.

62
Q

If System X is completely isolated from its surroundings, what must be true of ΔSX for any process that occurs inside System X?

A

ΔSX > 0

Since System X is isolated it cannot exchange energy or entropy with its surroundings. For the Second Law to hold, the system’s entropy must increase.

63
Q

What is the formula for converting between temperature in Celsius and Kelvin?

A

TK = TC + 273

TK = Kelvin Temperature
TC = Celsius Temperature

Some standard temperatures to memorize:
0º C = 273º K
25º C = 298º K
100º C = 373º K

64
Q

What is the formula for converting between temperature in Centigrade and Fahrenheit?

A

TF = (9/5)*TC + 32

TF = Fahrenheit Temperature
TC = Celsius Temperature

Some standard conversions to memorize:
32º F = 0º C
77º F = 25º C
212º F = 100º C

65
Q

Define:

conduction

A

Conduction is the transfer of thermal energy via molecular collisions. It requires physical contact between the systems exchanging energy.

When the molecules collide, molecules of the higher-energy system transfer some of their energy to the lower energy molecules of the other system, cooling the first and heating the second.

66
Q

Define:

convection

A

Convection is the thermal energy transfer via the movement of fluid in currents. It requires at least one medium which is capable of motion.

Differences in pressure or density drive warm fluid in the direction of cooler fluid, transferring heat away from a warmer object or towards a cooler one.

67
Q

Define:

A

Radiation is the thermal energy transfer via emission or absorption of electromagnetic waves.

Radiation does not require a medium and can occur through a vacuum.

68
Q

When a hot piece of metal is placed on a wooden table, causing the table to start to smoke, what is the primary form of heat transfer being exhibited?

A

Conduction

Conduction is the most efficient form of heat transfer; since the metal and table are in direct physical contact, conduction transfer will dominate.

69
Q

When the upstairs of a house is warmer than the basement, what form of heat transfer is being exhibited?

A

Convection

The decreased density of the warmer air causes it to rise to the top of the house, a classic example of a convection current.

70
Q

When the sun rises above the horizon, warming the ground, what form of heat transfer is being exhibited?

A

For the heat to get from the sun to the earth, it must travel across millions of miles of nearly empty space. Only electromagnetic radiation can carry energy across this distance.

71
Q

Define a material’s heat of vaporization, ΔHvap.

A

ΔHvap is the energy needed to vaporize one mole of a substance from its liquid phase to the gas phase at constant pressure.

ΔHvap may also be reported as heat per gram, in which case it is the specific heat of vaporization.
ΔHvap is always positive, since vaporization is an endothermic process.

72
Q

Define a material’s heat of fusion, ΔHfus.

A

ΔHfus is the energy needed to melt one mole of a substance from its solid phase to its liquid phase at constant pressure.

ΔHfus may also be reported as heat per gram, in which case it is the specific heat of fusion.
ΔHfus is always positive, since melting is an endothermic process.

73
Q

What is the work done on a thermodynamic system as a function of its pressure and volume?

A

Work = Δ(PV)

At constant pressure, W = PΔV
More generally, the work done during a thermodynamic process is the area under the curve of a P vs V diagram.