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Flashcards in Thermodynamics Deck (70)
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1
Q

What is a thermodynamic system?

A

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

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

2
Q

Define:

an open thermodynamic system

A

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

For example, a bottle of gas with no lid is an open thermodynamic system.

3
Q

Define:

a closed thermodynamic system

A

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

For example, a bottle of gas with the lid securely on is a closed thermodynamic system.

4
Q

Define:

an isolated thermodynamic system

A

An isolated thermodynamic system cannot exchange energy or mass with the environment.

For example, a closed, insulated container with a temperature that is independent of its environment is an isolated system.

5
Q

With regard to a thermodynamic system, what are the surroundings?

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 the characteristics of the system at that 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 when a system moves from a value of X1 to X2?

A

ΔX = X2 - X1

Since X is a state function, the path by which it gets from state 1 to state 2 is irrelevant; the change in X depends only on the initial and final 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, ΔH, is calculated 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. In such reactions, ΔH < 0, and heat is lost from the system to the environment.

Remember that “exo” is like “exit”; heat exits the system in an exothermic reaction.

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. In such reactions, ΔH > 0, and heat is absorbed by the system from the environment.

Remember that “endo” is like “into”; heat flows into the system in an endothermic reaction.

11
Q

Define:

standard enthalpy of formation (ΔHof)

A

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

For example, 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

0

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

13
Q

What are the requirements for a reaction to take place under standard conditions?

A

Standard conditions for a reaction require that:

  • Pressure = 1 atm
  • Temperature = 25º C = 298 K
  • Concentration = 1 M for all products and reactants
14
Q

If the chemical reaction

(1) 2A ⇒ C ΔH1

can be broken down into two steps:

(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 component steps 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?

2H2(g) ⇒ 4H(g)
ΔH1= -870 kJ/mol

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

A

-152 kJ

1) 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)

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

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

3) Finally, multiply by the number of moles produced (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)

To calculate the enthalpy change of a reaction, sum the enthalpies of formation of the products. Then subtract the sum of the enthalpies of formation of the reactants.

17
Q

Is this reaction endothermic or exothermic?

CH4 + 2O2 ⇒ CO2 + 2H2O

  • 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)]

-891 kJ/mol

Remember, on the MCAT, you won’t have a calculator. Approximate to get close to the final answer.

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 therefore 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

Remember, you won’t have a calculator on the MCAT, so approximate to get close to the correct answer.

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 1 ºC.

The higher the specific heat, the more energy input required to raise the substance’s temperature.

22
Q

What formula can be used to calculate the necessary quantity of heat to raise the temperature of a material?

A

q = mcΔT

Where:

q = heat required (J)
m = mass of substance present (g)
c = substance’s specific heat (J/g*ºC)
ΔT = temperature change (ºC)

23
Q

What is the specific heat of water?

A

4.18 J/g*K

This is a value that you should have memorized. According to this measurement, it requires 4.18 J to raise 1 g of water by 1 K (or ºC).

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 25 ºC. Which material changes its 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 temperature of the iron changes 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 = 45 ºC), while the water will only increase by 2 degrees (final T = 27 ºC).

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 the MCAT, you should focus on the fact that they all measure heat generated when a system’s temperature changes, 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 change in phase by breaking intermolecular forces, 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

The curve below represents a sample’s temperature versus the 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

28
Q

The curve below represents a sample’s temperature versus the heat added. Which heat (q) formula must be applied to calculate heat added to the system at each labeled point on the plot?

A

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

29
Q

Define:

entropy

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.

Colloquially, entropy is said to represent the possibility for “disorder” of a system. This definition is effective for most entropy questions on the MCAT.

30
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.

31
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.

32
Q

What is the sign of the entropy change for this chemical reaction?

2H2(g) + O2(g) ⇒ 2H2O(l)

A

ΔS < 0; in other words, the change in entropy is negative.

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 a lower final entropy value.

33
Q

What are the signs of 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.

34
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.

35
Q

Define:

Gibbs free energy (ΔG)

A

ΔG is a measure of the work which can be extracted from a thermodynamic system. ΔG can be calculated as such:

ΔG = ΔH - TΔS

A handy mnemonic for remembering the formula for ΔG is “Get Higher Test Scores.”

36
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.

37
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 nonspontaneous and the reverse reaction is spontaneous. This favors the creation of more reactants.

38
Q

Define:

an exergonic reaction

A

An exergonic reaction is any reaction for which ΔG < 0; because of this, 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.

39
Q

Define:

an endergonic reaction

A

An endergonic reaction is any reaction for which ΔG > 0; because of this, 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.

40
Q

Describe the spontaneity of a reaction where:

ΔHrxn > 0
ΔSrxn < 0

A

This reaction will always be nonspontaneous.

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

41
Q

Describe the spontaneity of a reaction where:

ΔHrxn < 0
ΔSrxn < 0

A

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

Remember that ΔG = ΔH - TΔS. When T is sufficiently low, the second term can be ignored, and ΔG will be negative due to ΔH. But when T becomes large enough, the positive sign of the -TΔS term dominates, and ΔG becomes positive.

42
Q

Describe the spontaneity of a reaction where:

ΔHrxn > 0
ΔSrxn > 0

A

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

Remember that ΔG = ΔH - TΔS. When T is sufficiently low, the second term can be ignored, and ΔG will be positive (corresponding to a nonspontaneous reaction) because ΔH > 0. However, when T becomes large enough, the negative sign of the -TΔS term dominates and ΔG becomes negative (corresponding to a spontaneous reaction).

43
Q

Describe the spontaneity of a reaction where:

ΔHrxn < 0
ΔSrxn > 0

A

This reaction will always be spontaneous.

Remember that 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.

44
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 it allows for a common point of reference between them.

45
Q

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

A

Temperature and internal energy are equivalent concepts.

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. In other words, they are at the same temperature.

46
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

Where:

k = Boltzmann’s constant (1.38 x 10-23 J/K)
T = absolute temperature (K)
n = number of moles of fluid molecules

47
Q

Define:

work

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 and measured in joules.

48
Q

Define:

heat

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 and measured in joules.

49
Q

Explain the First Law of Thermodynamics.

A

The First Law of Thermodynamics states that heat and work are the only possible mechanisms of energy flow into and out of a system. It also states that energy is always conserved; for this reason, energy change can be calculated as:

ΔE = Δq + Δw

50
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, on the system.

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

51
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, on the environment.

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

52
Q

What is the direction of heat flow for a thermodynamic system that is surrounded by an environment at a higher temperature? 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 as well.

53
Q

What is the direction of heat flow for a thermodynamic system that is surrounded by an environment at a lower temperature? 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.

54
Q

What are the characteristics of an adiabatic thermodynamic process?

A

An adiabatic process is one in which heat cannot flow. In such a process, Δq = 0 and ΔE = Δw; all energy change is due to work being done on or by the system.

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

55
Q

What happens to the temperature of a system 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’s temperature will increase as well.

56
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’s temperature will decrease as well.

57
Q

What are the characteristics of an isothermal thermodynamic process?

A

An isothermal process is one in which temperature is held constant. In such processes, 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 usually happen slowly enough that the system’s temperature can constantly equilibrate.

58
Q

In what direction does 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.

59
Q

In what direction does 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 and losing energy. Since any isothermal process must have an overall ΔE = 0, Δq must be positive to compensate.

60
Q

Explain the Second Law of Thermodynamics.

A

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

ΔSuniverse > 0
ΔSsystem+ ΔSsurroundings > 0

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

61
Q

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

A

ΔSX > 0

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

62
Q

What formula can be used to convert between temperatures in Celsius and Kelvin?

A

TK = TC + 273

Where:

TK = Kelvin temperature
TC = Celsius temperature

Some standard temperatures to memorize:

  • 0º C = 273 K
  • 25º C = 298 K
  • 100º C = 373 K
63
Q

What formula can be used to convert between temperatures in Celsius and Fahrenheit?

A

TF = (9/5)*TC + 32

Where:

TF = Fahrenheit temperature
TC = Celsius temperature

Some standard conversions to memorize:

  • 32º F = 0º C
  • 77º F = 25º C
  • 212º F = 100º C
64
Q

Define:

conduction

A

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

65
Q

Define:

convection

A

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

66
Q

Define:

radiation

A

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

Radiation does not require a medium, so it can occur through a vacuum.

67
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 will dominate.

68
Q

When the upstairs rooms of a house are 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.

69
Q

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

A

Radiation

For the heat to move from the sun to the earth, it must travel across millions of miles of near-empty space. Only electromagnetic radiation can carry energy across this distance, and electromagnetic waves transfer heat through radiation.

70
Q

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

A

Work = PΔV

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