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CHEM 107 > Thermodynamics > Flashcards

Thermodynamics Flashcards

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1
Q

Energy

A

Capacity to do work or produce heat

SI Unit: joule(J) = kg m ^2/ s ^2

Another Unit: 1 cal = 4.184J

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2
Q

System

A

Object or collection of objects being studied

Like Chemical Reactions

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3
Q

Surrondings

A

Everything outside the system that can exchange energy and/or matter with the system

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4
Q

Thermodynamics

A

The study of energy and the exchange of energy between the system and the surroundings.

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5
Q

State Functions

A

A PROPERTY of the system that DEPNDS ONLY on its PRESENT STATE

Represented with UPPERCASE VARIABLES

Analogy:
The floor you are on in a building

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6
Q

Path Functions

A

Are DEPENDENT on the path the system takes to get from initial to final state

Represented with LOWERCASE VARIABLES

Analogy:
How you get from the 1st floor to the 3rd floor

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7
Q

First Law Of Thermodynamics

A

Law of Conservation Of Energy

Energy can be converted, but NOT created nor destroyed

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8
Q

Internal Energy

A

energy contained within a system

Two Kinds:
1) Kinetic Energy (Energy associated with motion)
Example: Translation, Vibration, Rotation

2) Potential Energy (Energy associated with position)
Example: Chemical Bonds, Intermolecular Forces

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9
Q

Chemical Energy

A

Energy associated with bonds in molecules

Making bonds RELEASE energy

Breaking bonds COSTS energy

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10
Q

Reaction Energy

A

Energy takin in or given off during reaction

change in E = E(products) minus E(reactants)

or

change in E = E(final) minus E(initial)

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11
Q

Change in Energy

A

deltaE(reaction) = deltaE(products) minus deltaE(reactants)

deltaE(reaction) is NOT the same as deltaE(reactants)

Basically with this, chemical reactions can go both ways, it just depends on if energy is going in or out

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12
Q

-deltaE vs +deltaE

A

-deltaE means energy is LEAVING

+deltaE means energy is being ADDED

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13
Q

Exothermic vs Endothermix

A

Exothermic- Heat is LEAVING/ EXITING
Exo -> sounds like “Exit”

Endothermic- Heat is ENTERING/ ABSORBED
Endo -> sounds like “In To”

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14
Q

Heat Capacity

A

A transfer of thermal energy into or out of a system
Not the same as Temperature

Lower Heat Capacity = Easier to change temperature
Higher Heat Capacity = Harder to change temperature

T = measure of thermal energy
q = change in thermal energy
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15
Q

Specific Heat Capacity

A

amount of heat required to raise the temperature of one gram of a substance by one degree Celsius

q = (m)(C)(deltaT)

m = mass(g)
deltaT = change in temperature(K, C- as long as they match)
C = specific heat capacity (J/g K, J/g C- as long as they match)
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16
Q

Molar Heat Capacity

A

amount of heat required to raise the temperature of ONE MOLE of a substance by one degree Celsius

17
Q

Coffee Cup Calorimeter

A

q(reaction) = -q(water) + (-q(calorimeter))

where:
-q(water) = -[(m(H2O)))(C(H2)))(deltaT(H2O))] –> You know this
-q(calorimeter) = (-C(calorimeter))(deltaT(calorimeter))
This is taking into account that the “walls” of the calorimeter are
absorbing some heat

deltaT(H2O) = deltaT(calorimeter)

18
Q

Specific Heat Of Water

A

4.184

19
Q

(Bomb) Calorimeter

A

All energy is contained inside. Essentially, its a chamber inside of a bigger chamber like a steel cup.

q(reaction) = -(C(calorimeter))(deltaT(calorimeter))

-C(calorimeter) = -[(m(H2O))(C(H2O)) + C(calorimeter)]

20
Q

Enthalpy

A

The heat flow under constant pressure

Is a state function

Denoted as H

deltaH = H(products) - H(reactants)
or
deltaH = H(final) - H(initial)

deltaH>0 –> Endothermic
deltaH<0 –> Exothermic

21
Q

Extensive Property

A

When you double the reactants, you also double the product

Example: Enthalpy

22
Q

Enthalpy Of Phase Changes

A

Getting Hotter:
solid—(Melting)—>liquid—(Vaporizing)—> gas

Getting Colder:
gas—(Condensation)—>liquid—(Freezing)—> solid

Skipping Phases:
solid —(Sublimation)—> gas
gas—(Deposition)—> solid

23
Q

Heat of Fusion vs Heat of Vaporization

A

Heat of Fusion = Melting

Heat of Vaporization = Vaporize

24
Q

Bond Enthalpy

A

The energy required to break one mole of a particular bond in the gaseous state

deltaH(reaction) = summation(reactant bond enthalpies) - summation(product bond enthalpies) —> initial - final

Takes energy to break bonds
Energy is given off when a bond is formed

25
Q

Hess’s Law

A

Basically, if a reactions takes several steps, deltaH is equal to the sum of all the individual steps

Extensive Property

Can treat different chemical equations as a system of equations to add everything on the reactants and the products. If you add, the equations together, you also add the deltaH’s together

If you flip a chemical equation, you also change the sign of the deltaH

26
Q

Enthalpy of Formation

A

The enthalpy change for the reaction in which one mole of a compound is made from its elements in their elemental form

Denoted as: deltaH(f)

27
Q

Standard Enthalpy of Formation

A 
When deltaH(f) is measured under standard conditions:
Temperature = 25 C
atm = 1.00
Molarity(M) = 1

deltaH(f) of a pure element in its standard state = 0

deltaH(reaction) = summation((n)(deltaH(f(products)))) - summation((m)(deltaH(f(reactants))))
where n and m are the stoichiometric coefficients of each product or reactant

28
Q

Spontaneous

A

Process that occurs without any ongoing outside intervention; naturally

29
Q

Entropy

A

A measure of the molecular disorder in a system

Denoted as: S

Units: J/mol K

More Disorder = More Entropy

Increase in Temperature = More Entropy

30
Q

Second Law of Thermodynamics

A

The entropy of the universe is always increasing

31
Q

deltaS(surroundings)

A

-deltaH(reaction) / T

or

q(surroundings) / T = q(system) / T

32
Q

Third Law of Thermodynamics

A

The entropy of a pure, perfect crystal at 0 K is zero

33
Q

Standard Entropy of a Reaction

A

deltaS(rxn) = summation((n)(deltaS(products))) - summation((m)(deltaS(reactants)))

where n and m are stoichiometric coefficients for each product and reactant

34
Q

Gibbs Free Energy

A

deltaG = deltaH - T(deltaS)

Denoted as: G

A state function

deltaG > 0 —> non-spontaneous
deltaG < 0 —> spontaneous

T = deltaH / deltaS

35
Q

Temperature Dependance of Gibbs Free Energy

A

deltaH - T(deltaS) = deltaG

deltaH deltaS
negative positive = Always Spontaneous (deltaG = -)
positive negative = Never Spontaneous (deltaG = +)
positive positive = Spontaneous at HIGH Temperatures (deltaG =-/+)
negative negative = Spontaneous at LOW Temperatures (deltaG =-/+)

Temperature is always positive because it is in Kelvin

CHEM 107 (3 decks)

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