L.12 Gibbs Free Energy Flashcards
How is a system classified?
It is the matter being observed and it is based on
what is or is not being exchanged with the surroundings
Isolated System
NO EXCHANGE in HEAT WORK or MATTER
Closed System
Exchange in HEAT & WORK
but
NOT MATTER
Open System
Exchange of Work Heat and Matter
Definition of undergoing a process
When a system undergoes a change in one or more of its properties
Concentration, Temperature or Pressure
also known as,
CHANGE OF STATE of SYSTEM
First Law of Thermodynamics
The total energy of an isolated system is constant;
energy can be transformed from one form to another but can be neither created nor destroyed.
∆U = Q - W
The internal energy of system = heat added to the system - work done BY the system
W = work = changes in pressure and volume
Four processes characterized by a single constant property.
- Isothermal
- Adiabatic
- Isobaric
- Isovolumetric (isochoric)
Isothermal Process
∆U = Q - W
Temperature is Constant
The total energy of the system is constant
∆U = 0, then Q = W
The heat added to the system is equal to the work done by the system
Adiabatic Process
∆U = Q - W
Constant heat/ no heat exchange with the environment
Q = 0, then ∆U = - W
The change in internal energy is equal to the work done ON the system
Isobaric Process
∆U = Q - W
No change in pressure/pressure is constant
The slope of Line = 0 in the pressure-volume graph
Isovolumetric Process
∆U = Q - W
Volume is constant, gas neither expands or compresses
W = 0, the ∆U = Q
The change in internal energy is equal to the heat added to the system
What is a State Function?
Name all 8 State Functions & 2 Process Functions
State Function
- Density
- Pressure
- Temperature
- Internal Energy U
- Gibbs Free Energy
- Entropy
- Enthalpy
- Volume
Process Function
- Work W
- Heat Q
Standard Conditions
vs
STP
Standard Conditions
1 atm, 25C/ 298K/ 77F
Gibbs free energy, Enthalpy & Entropy
(Kinetics, Thermodynamics & Equilibrium)
STP for IDEAL GAS Problems
0C/273/32F
The Standard State of an Element
is its most prevalent form under standard conditions; standard enthalpy, entropy and free energy are all calculated under standard conditions.
H2 gas
H2O liquid
NaCl solid
O2 gas
C solid, graphite
Temperature
Temperature: Scaled measure of the Average Kinetic Energy of molecules of a substance
Fusion & Freezing
Fusion/Melting &
Freezing/Crystalization/ solidification
Occur at the interface of solid and gas phase
Vaporization & Condensation
Vaporization
- Occurs only above boiling Point
- The ambient pressure is equal to the pressure of the liquid
Condensation
- Forcing gas back into the liquid phase
Sublimation & Deposition
Sublimation = Solid into Gas
Deposition = Gas into Solid
Critical Point & Triple point
At temperatures above the critical point, all three phases are indistinguishable
At the triple point, all three phases of matter exist in equilibrium
Amorphous solid
In condensed matter physics and materials science, an amorphous or non-crystalline solid is a solid that lacks the long-range order that is characteristic of a crystal.
Melts at large range of Temperatures
Phase Diagram
Heat
Heat is the transfer of energy that results from differences of temperature between two substances
Heat Q
Endothermic = Absorbs Heat
Exothermic = Releases Heat
∆H = Q at Constant Pressure
Heat Is transferred From Warmer Substances —-> To Cooler Ones
Enthalpy
Is a measure of the potential energy of a system found in intermolecular attractions and chemical bonds
Heat changes at Constant Pressure (state function)
∆H rxn = ∆H products - ∆H reactants
Positive Change = Endothermic
Negative Change = Exothermic
The heat content of a system undergoing heating, cooling, or phase changes is the sum of all the respective energy changes.
Hess’s Law
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.
Enthalpy changes of a reaction are additive
(switch signs when adding reactions/flipping them)
Enthalpy can also be calculated using heats of formation, heats of combustion, or bond dissociation energies.
Entropy
Is a measure of the degree to which energy has been spread throughout a system or between a system and its surroundings.
- is a ratio of heat transferred per mole unit Kelvin
- is maximized at equilibrium
∆S = Qrev / T
Second Law of Thermodynamics
The Second Law of Thermodynamics is commonly known as the Law of Increased Entropy.
Energy is spontaneously dispersed if it’s not hindered - at a specific Temperature.
∆Suniverse = ∆Ssystem - ∆S surroundings > 0
While quantity remains the same (First Law), the quality of matter/energy deteriorates gradually over time.
How so? Usable energy is inevitably used for productivity, growth, and repair. In the process, usable energy is converted into unusable energy. Thus, usable energy is irretrievably lost in the form of unusable energy.
Third Law of Thermodynamics
The Third Law of Thermodynamics refers to a state known as “absolute zero.” This is the bottom point on the Kelvin temperature scale. The Kelvin scale is absolute, meaning 0° Kelvin is mathematically the lowest possible temperature in the universe. This corresponds to about -273.15° Celsius, or -459.7 Fahrenheit.
Nothing can exist after this point
Standard Heat of Formation
Standard Heat of Reaction
- Enthalpy required to form 1 mole of compounds from its elements under standard conditions
- Enthalpy changes of a reaction carried under standard conditions
Bond Dissociation Energy
Breaking a particular bond between atoms in the gas phase
Bond Breaking = Endothermic
Bond Forming = Exothermic (becoming more stable)
KJ/mol of bonds broken
Gibbs Free Energy
The change is Gibbs free energy determines if a process will be spontaneous or non-spontaneous.
∆G = ∆H - T∆S
∆H ∆S OUTCOME
+ + Spontaneous at HIGH T
+ - Non-spontaneous at ALL T
- Spontaneous at ALL T
- Spontaneous at LOW T
Kinetics vs Thermodynamics
Sometimes the thermodynamically favored/stable product will have high activation energy, therefore will have slower kinetics.
Standard Gibbs Free energy from Equilibrium Constant
∆G* rxn = -RT ln Keq
Standard free energy change = - ideal gas constant ln (equilibrium constant)
When in Process
∆G rxn = ∆G* rxn + RT ln Q = RT ln Q/Keq
Q/Keq < 1 (spontaneously move forward)
Q/Keq > 1 (spontaneously move backward)
Q/Keq = 1 (equillibrium)
Equations for
Heat Transferred during phase change
Heat Transferred during no phase change
q = mc∆T
q =mL (latent heat)
