Thermodynamics Flashcards
Energy
Capacity to do work or produce heat
SI Unit: joule(J) = kg m ^2/ s ^2
Another Unit: 1 cal = 4.184J
System
Object or collection of objects being studied
Like Chemical Reactions
Surrondings
Everything outside the system that can exchange energy and/or matter with the system
Thermodynamics
The study of energy and the exchange of energy between the system and the surroundings.
State Functions
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
Path Functions
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
First Law Of Thermodynamics
Law of Conservation Of Energy
Energy can be converted, but NOT created nor destroyed
Internal Energy
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
Chemical Energy
Energy associated with bonds in molecules
Making bonds RELEASE energy
Breaking bonds COSTS energy
Reaction Energy
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)
Change in Energy
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
-deltaE vs +deltaE
-deltaE means energy is LEAVING
+deltaE means energy is being ADDED
Exothermic vs Endothermix
Exothermic- Heat is LEAVING/ EXITING
Exo -> sounds like “Exit”
Endothermic- Heat is ENTERING/ ABSORBED
Endo -> sounds like “In To”
Heat Capacity
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
Specific Heat Capacity
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)
Molar Heat Capacity
amount of heat required to raise the temperature of ONE MOLE of a substance by one degree Celsius
Coffee Cup Calorimeter
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)
Specific Heat Of Water
4.184
(Bomb) Calorimeter
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)]
Enthalpy
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
Extensive Property
When you double the reactants, you also double the product
Example: Enthalpy
Enthalpy Of Phase Changes
Getting Hotter:
solid—(Melting)—>liquid—(Vaporizing)—> gas
Getting Colder:
gas—(Condensation)—>liquid—(Freezing)—> solid
Skipping Phases:
solid —(Sublimation)—> gas
gas—(Deposition)—> solid
Heat of Fusion vs Heat of Vaporization
Heat of Fusion = Melting
Heat of Vaporization = Vaporize
Bond Enthalpy
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
Hess’s Law
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
Enthalpy of Formation
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)
Standard Enthalpy of Formation
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
Spontaneous
Process that occurs without any ongoing outside intervention; naturally
Entropy
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
Second Law of Thermodynamics
The entropy of the universe is always increasing
deltaS(surroundings)
-deltaH(reaction) / T
or
q(surroundings) / T = q(system) / T
Third Law of Thermodynamics
The entropy of a pure, perfect crystal at 0 K is zero
Standard Entropy of a Reaction
deltaS(rxn) = summation((n)(deltaS(products))) - summation((m)(deltaS(reactants)))
where n and m are stoichiometric coefficients for each product and reactant
Gibbs Free Energy
deltaG = deltaH - T(deltaS)
Denoted as: G
A state function
deltaG > 0 —> non-spontaneous
deltaG < 0 —> spontaneous
T = deltaH / deltaS
Temperature Dependance of Gibbs Free Energy
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
