Thermochemistry (7) Flashcards
Isolated
The system cannot exchange energy (heat and work) or matter with the surrounding
ex: insulated bomb calorimeter
Closed
The system can exchange energy (heat and work) but not matter with the surrounding;
ex: a steam radiator
Open
The system can exchange both energies (heat and work) and matter with the surrounding
ex: a pot of boiling water
First law of Thermodynamics
ΔU = Q - W
ΔU = change in internal energy of the system Q = heat added TO THE system W = work done BY THE system
Isothermal Process
Occurs when the temperature is constant
Constant Temperate = Constant ΔU = ΔU = 0
Q = W
Hyperbolic curve of Pressure vs. Volume graph
Adiabatic Process
Occurs when no heat is exchanged between the system and the environment, so Q = 0
ΔU = -W
Temperature is not constant
Isobaric Process
occurs when the pressure of the system is constant.
Pressure vs. Volume graph line appears flat line b/c Pressure is constant
Isovolumetric (Isochoric) proccess
No change in volume, no work is performed in such a process.
ΔU = Q (change in internal heat is equal to the heat added to the system)
Pressure vs. Volume graph line appears verticle line
State functions
When I’m under PRESSURE, and feeling DENSE, all I want to do is watch TV and get HUGS.
- Pressure (P)
- density (p)
- Temperature (T)
- Volume (V)
- Enthalpy (H)
- Internal energy (U)
- Gibbs free energy (G)
- Entropy (S)
Standard conditions vs. STP
Standard conditions: 25 °C (298K) , 1atm pressure, 1M concentration
STP: 0 °C (273K) , 1atm pressure
Heat
the specific form of energy that can enter or leave a system
Temperature
the measure of the average kinetic energy of the particles in a system.
Exothermic and Endothermic
Exothermic = -ΔH Enthalpy
Endothermic = + ΔH Enthalpy
Calorimetry
process of measuring transferred heat
Heat/energy
J vs Cal Conversion
1cal = 4.184J
Heat transfer equation
q = mCΔT
Standard enthalpy of formation
enthalpy required to produce one mole of a compound from its element in their standard states.
Standard enthalpy of reaction
enthalpy change accompanying a reaction being carried out under the standard condition.
ΔH°rxn = ΣH°f, products - ΣH°f, reactants
Hess’s Law
total change in potential energy of a system is equal to the change of potential energies of the individual step of the process.
Enthalpy changes in reactions are additive.
ΔH = ΔH1 + ΔH2 + ΔH3
Bond Energy
ΔH°rxn = ΔH bond broken - ΔH bonds formed
bond breakage is endothermic
the bond formation is exothermic
\+ΔH = total energy absorbed -ΔH = total energy released
Second law of Thermodynamics
energy spontaneously disperses from being localized to becoming spread out. disorder and entropy.
ΔS universe = ΔS system + ΔS surrounding > 0
ΔS = ΣS°f, products - ΣS°f, reactants
Entropy
the measure of the spontaneous dispersal of energy at a specific temperature.
S = Qrev / T
Gibbs free energy
ΔG = ΔH - TΔS Goldfish are (equal) Horrible without (minus) Tartar Sauce
If ΔG = 0 , the system is in a state of equilibrium,
ΔH = TΔS
Exergonic
The movement towards the equilibrium. When energy released.
ΔG < 0
spontaneous
Endergonic
The movement away from the equilibrium
ΔG > 0
non-spontaneous
Signs of ΔH and ΔS corresponding to Temperature
ΔH (+) ΔS (+) Spontaneous at High T
ΔH (+) ΔS (-) Nonspontaneous at all T
ΔH (-) ΔS (+) Spontaneous at all T
ΔH (-) ΔS (+) Spontaneous at low T
Standard Gibbs free energy formula
ΔG = ΣG°f, products - ΣG°f, reactants
Free energy, Keq, and Q
ΔG°rxn = -RT lnKeq
greater the Keq value, more negative the Gibbs free energy and leads to spontaneous rxn.
ΔGrxn = ΔG°rxn + RTlnQ = RTln Q/Keq
If the Q/Keq ratio is less than one, then the Gibbs free energy is negative, spontaneous rxn.
Q < Keq –> spontanous
Q > Keq –> non-spontanous
