7. Thermochemistry DONE

Question 1

What is thermochemistry?

Answer 1

Thermochemistry: the study of the energy changes that accompany chemical and physical processes

Question 2

What’s the difference between the system and the environment when thinking about a reaction? (thermochemistry)

Answer 2

System: is the matter that is being observed, the reactants and the products
Surroundings/Environment: everything apart from the reactants and the products.

Question 3

What’s the difference between?
Isolated system:
Closed Systems:
Open Systems:

Answer 3

this is ENERGY and MATTER

Isolated Systems: exchange neither matter nor energy with the environment (e.g. insulated bomb calorimeter)

Closed Systems: can exchange energy but not matter with the environment (steam radiator)

Open Systems: can exchange both energy and matter with the environment (pot of boiling water)

Question 4

What is the first law of thermodynamics and what is the related equation?

Answer 4

1st Law of Thermodynamics - Energy cannot be created or destroyed.

∆U= Q – W

∆U = change in the internal energy of the system (temperature)
Q= heat added to the system (Enthalpy or ∆H)
W= work done

Question 5

What is the difference between:
Isothermal:
Adiabatic:
Isobaric:
Isovolumetric/isochoric:

∆U= Q – W

∆U = change in the internal energy of the system
Q= heat added to the system
W= work done

Answer 5

Isothermal: processes occur at a constant temperature (work is done BY the system) ∆U= W (Q=0)
Adiabatic: processes exchange no heat with the environment (work is done ON the system) ∆U= -W
Isobaric: processes occur at a constant pressure
Isovolumetric (isochoric): processes occur at a constant volume ∆U= Q

===============================

∆U= Q – W

∆U = change in the internal energy of the system
Q= heat added to the system
W= work done

Question 6

What is the difference between enthalpy and temperature?

Answer 6

Temperature is the average kinetic energy within a system.

Enthalpy ΔH- is the exchange of temperature/heat with the system and the surroundings

Question 7

What is the first law of thermodynamics?

What is the equation?

Answer 7

∆U= Q – W

∆U = change in the internal energy of the system
Q= heat added to the system
W= work done

First law of thermodynamics is that energy cannot be created or destroyed.

Question 8

What is a coupling reaction?

Answer 8

This occurs when you put a spontaneous and non-spontaneous reaction together, the spontaneous reaction will supply the energy the that the non-spontaneous reaction needs.

Question 9

A person snaps an ice pack and places it on one leg. In terms of energy transfer, what would be considered the system, and what would be the surroundings in this scenario?

Answer 9

System: the ice inside the bag

Surroundings: things outside the bag (i.e. the leg that the coolness will flow to)

Question 10

What is unique about each of the following processes?
In other terms, what is NOT happening

Isothermal:
Adiabatic:
Isobaric:
Isovolumetric (isochoric):

Answer 10

∆U= Q – W

Isothermal: There is no change in temperature: ∆U=0, Q=W
Adiabatic: There is no heat exchange: ∆U= -W, Q=0
Isobaric: No change in pressure, line appears flat in a P-V (pressure to volume) graph
Isovolumetric: no change in volume W=0, ∆U=Q

∆U= Q – W

∆U = change in the internal energy of the system or temperature
Q= heat added to the system
W= work done by the system

Question 11

What is an adiabatic process?

Answer 11

There is no heat exchange:
∆U= -W,
because Q=0 (i.e. no heat exchange)

∆U= Q – W

Example: release of air from a tire

∆U = change in the internal energy of the system
Q= heat added to the system
W= work done

In thermodynamics, an adiabatic process (Greek: adiábatos, “impassable”) is a type of thermodynamic process that occurs **without transferring heat or mass ** between the thermodynamic system and its environment. Unlike an isothermal process, an adiabatic process transfers energy to the surroundings only as work.[1][2] As a key concept in thermodynamics, the adiabatic process supports the theory that explains the first law of thermodynamics.

Question 12

What is a state function and what are the different state functions?

Answer 12

State functions describe the physical properties of an equilibrium state; they are pathway independent

Pressure, density, temperature, volume, enthalpy (H), internal energy (U), Gibbs free energy (G), and entropy (S)

Mnemonic: State Pressure: When I’m under pressure and feeling dense, all I want to do is watch TV and get HUGS

Question 13

What are the standard conditions?

Answer 13

Temperature: 298 K
Mass: 1 ATM
Concentration: 1 M

Question 14

What’s the difference between standard pressure and standard conditions?

Answer 14

The main difference is temperature: 0 C (SP)vs 25C (SC)

STP: 0 C (273 K), 1 ATM pressure (standard conditions)

Standard Conditions= 25 C (298K), 1 atm pressure, 1 M concentration

Question 15

What is the standard state of the element?

Answer 15

it is it’s most prevalent form at standard conditions

Question 16

What is:

Sublimation/ deposition:
Fusion:
Vaporization:

Answer 16

Sublimation: from solid to gas (dry ice)
Fusion: melting (solid to liquid)
Vaporization: evaporation or boiling

Question 17

What is the triple point?

Answer 17

point where all three phases meet. Temperature and pressure that all three phases live at equilibrium. There is no distinction between the phases

Question 17

What are standard conditions? When are standard conditions used for calculations?

Answer 17

Kinetics, equilibrium, and thermodynamics calculations used standard conditions,

Standard conditions are: 25°C (298K), 1atm pressure, and 1 M concentrations

Question 18

What is a state function vs process function?

Answer 18

State Function: State functions are properties of a system at equilibrium and are independent of the path taken to achieve the equilibrium; they may be dependent on one and other

Process Function: Process functions define the path between equilibrium states and include Q(heat) and W (work)

Question 19

On a phase diagram, what is the definition of a critical point?

Answer 19

The critical point is the temperature and pressure above which the liquid and gas phases are indistinguishable and the heat of vaporization is zero.

Question 20

What’s the difference between temperature and heat?

Answer 20

Temperature is a scaled measure of the average kinetic energy of a substance

Heat (∆Q) is the transfer of energy that results from differences of temperature between two substances

Heat is the same as enthalpy (at constant pressure)

Heat and temperature are different. Heat is a specific form of energy that can enter or leave a system, while the temperature is a measure of the average kinetic energy of the particles in a system

Question 21

What is the zeroth Law of thermodynamics?

Answer 21

Zeroth Law of Thermodynamics – implies that objects are in thermal equilibrium only when their temperatures are equal.

Question 22

How are enthalpy and heat related?

Answer 22

Enthalpy (∆H) is equivalent to heat (Q) under constant pressure, which is an assumption the MCAT usually makes for thermodynamics problems.

Question 23

If a system is endothermic, what is the ∆Q?
If a system is exothermic, what is the ∆Q?

Answer 23

Endothermic: ∆Q> 0 Positive
Exothermic: ∆Q< 0 Negative

Question 24

What is the specific heat defined as?

Answer 24

defined as the amount of energy required to raise the temperature of one gram of a substance by one degree Celsius.

Question 25

What is the specific heat of water? (x2)

Answer 25

1 (cal)/(gram x Kelvin)

or

4.186 J (g K)

Question 26

What equation do you have to use when doing phase changes with one substance?

Why is temperature not included in the phase change equation?

Answer 26

Temperature is not included because with a phase change, there is no change in temperature. You just continue to add heat energy, it does not increase the temp of the system, just converts it to another phase.

You need to know the latent heat for phase change.

q=mL

m= mass
L= latent heat (cal/g)
q= heat

Question 27

What is enthalpy? ΔH

Answer 27

Enthalpy is a measure of the potential energy of a system found in intermolecular attractions and chemical bonds

Also, enthalpy is heat

Notation: ΔH

Question 28

What is Hess’s law?

Answer 28

States that the total change in potential energy of a system is equal to the changes of potential energies of the individual steps of the process.

They are always PATH INDEPENDENT.

Question 29

What is the equation for the enthalpy of bond dissociation?

Answer 29

ΔH°rxn = Σ ΔHbonds broken − Σ ΔHbonds formed

Question 30

How do you define entropy?

Answer 30

while often thought of as disorder, it is a measure of the degree to which energy has been spread throughout a system or between a system and its surroundings.

AKA, how much energy is spread out, or how widely spread-out energy becomes, in a process.

Question 31

What is the equation for entropy?

Answer 31

Entropy = ∆S

∆Change in Entropy, units= J/(mol*K)
When energy goes into a system, entropy INCREASES.
When energy goes out of a system, the energy decreases

Qrev= heat that is gained or lost in a reversible process

T= temperature in Kelvin

Question 32

When is entropy (∆S) at its maximum?

Answer 32

At equilibrium.

Entropy is NOT ∆G

∆G= Gibbs free energy
∆H=Enthalpy
∆S= Entropy

Question 33

What is the second law of thermodynamics?

Answer 33

Entropy

Energy spontaneously disperses from being localized to being spread out if it is not hindered from doing so.

Question 34

Rank the phases of matter from lowest to highest entropy?

Solids vs liquids vs gas

Answer 34

Gas has the highest, liquids in the middle, solids have the lowest entropy.

Question 35

Describe entropy in terms of energy dispersal and disorder.

Answer 35

Entropy increases as a system has more disorder or freedom of movement, and energy is dispersed in a spontaneous system.

Entropy of the universe can never be decreased spontaneously

Question 36

Which of the following situation result in an increase or decrease in entropy?

H20 going from liquid to solid
Dry ice sublimates into carbon dioxide
NaCl from solid to aqueous solution
N2+ H2–> 2NH3
An Icepack is placed on a wound.

Answer 36

H20 going from liquid to solid –> Decreasing (freezing)
Dry ice sublimates into carbon dioxide–>Increasing (sublimation)
NaCl from solid to aqueous solution –>Increasing (dissolution)
N2+ H2 –> 2NH3–> Decreasing (fewer moles)
An Icepack is placed on a wound. –>Increasing (heat is transferred)

Question 37

What is gibbs free energy?

Answer 37

It is a value that is derived from both enthalpy and entropy.

∆G is a measure of the change in the enthalpy and the change in entropy as a system undergoes a process, and it indicates whether a reaction is spontaneous.

Change in the free energy is the maximum amount of energy released by a process, occurring at constant temperature and pressure.

∆G=∆H-T∆S

∆G= Gibbs free energy
∆H= enthalpy
T= temperature
∆S= Entropy

Question 38

How will you know whether Gibbs free energy will be positive or negative based on enthalpy and entropy?

Answer 38

You will not, will will depend on temperature.

Question 39

Given the equation, will the reaction be spontaneous or not spontaneous given the following?

∆G=∆H-T∆S

Answer 39

Question 40

If you have the following equation…
How would you solve for:

1. Determining Gibbs Free energy
2. At what temperature would it be at equilibrium?
3. What it you suddenly put in a lot of ammonia in the beaker?

Answer 40

1. Convert -93 kg/mol to J/Mol.
   Then use equation ∆G=∆H-T∆S
   Answer @ 500 K = 7000 J/mol or 7kJ/mol
2. Set ∆G to 0 (equilibrium) and solve for T
   Answer=465 K
3. The value of Q would increase significantly. The reaction would shift to the left, forming more reactants.