Thermochemistry Flashcards
The law of conservation of energy
Energy can neither be created nor destroyed, but that all thermal, chemical, potential, and kinetic energies are interconvertible
Isolated Systems
no exchange of energy/matter with the environments (bomb calorimeters)
Closed Systems
can exchange energy but not matter with the environment
Open Systems
can exchange both energy and matter with the environment
Isothermal
temperature remains constant
Adiabatic
no heat exchange occurs
Isobaric
pressure of a system remains constant
Isovolumetric
aka isochoric
volume remains constant
Heat
the transfer of thermal energy from one object to another
Endothermic
Reactions that absorb thermal energy
Exothermic
Reactions that release thermal energy
Constant Volume and Constant Pressure Calorimetry
Used to indicate conditions under which the heat flow is measured
Equation to calculate heat absorbed or released
q=mcΔT
m= mass c= specific heat ΔT= change in temperature
State Functions
described by the macroscopic properties of the system, and the magnitudes of these properties depend on only the initial and final states of the system, not the path.
Include: pressure, density temperature, volume, enthalpy, internal energy, free energy, and entropy
Enthalpy
(H) used to express the heat changes at constant pressure
Standard heat of formation
the enthalpy change that would occur if one mole of a compound was formed directly from its elements in their standard states
Standard heat of reaction
the hypothetical enthalpy change that would occur if one mole of a compound was formed directly from its elements
Hess’s Law
states that enthalpies of reactions are additive
the reverse of any reaction has an enthalpy of the same magnitude as that of the forward reaction (sign is opposite)
Bond Dissociation Energy
an average of the energy required to beak a particular type of bond in one mole of gaseous molecules
Bond Enthalpy
the standard heat of reaction can be calculated using the values of bond dissociation energies of particular bonds
Entropy
(S) the measure of the distribution of energy throughout a system
Gibbs free energy
combines the two factors that affect the spontaneity of a reaction: changes in enthalpy (ΔH) and changes in entropy (ΔS)
ΔG= ΔH - TΔS
If ΔG is negative
the reaction is spontaneous
If ΔG is positive
the reaction is nonspontaneous
If ΔG is zero
the reaction is in equilibrium (ΔH=TΔS)
If ΔH is negative and ΔS is positive
reaction is spontaneous at all temperatures
If ΔH is positive and ΔS is negative
reaction is nonspontaneous at all temperatures
If ΔH is positive and ΔS is positive
spontaneous at very high temperatures
If ΔH is negative and ΔS is negative
spontaneous at very low temperatures
Reaction Quotient
(Q) once a reaction has started the standard state conditions no longer hold so we use Q to determine where we are in the reaction (before eqlbm, after eqlbm or at eqlbm)
Standard Conditions
298K
1 atm
1 M
Standard Temperature and Pressure
273K
1 atm
** used for ideal gas problems
Specific Heat
the specific heat required to raise the temperature of 1 gram of substance by 1 degree
Heat Absorbed or Released in a Process (Equation)
q=mcΔt
m- mass
c- specific heat
t- temp
During Phase Changes we use
q=mL
L- latent heat
Entropy Equation
ΔS = Qrev/T
Equation for Gibbs Free Energy
ΔG= ΔH - TΔS
Standard Free Energy for the Reaction Equation
ΔGrxn = -RT lnKeq
