Thermochemistry Flashcards
the energy of a system can be changed either by heat transfer, or by doing work
only two ways in which energy can enter/leave a system
conservation of energy of a system
first law of thermodynamics
energy can never be created or destroyed
conservation of energy
Equation: Energy of a system: ∆E
∆E = q - W
SI Unit= Joule?
Note: q = heat absorbed by the system from its surroundings; W is the work done BY the system on its surroundings
***If W is work done ON the system by the surroundings, then ∆E = q + W
energy transfer that occurs as a result of a temperature difference between the system and its surroundings
transfer will occur spontaneously from a warmer system to a cooler system
measured in the same units as Energy (Joules), calories (cal), kcal (kilocalorie), or kJ (kilojoule)
Heat
process in which heat is transferred from one particle to another through collisions
metals are good ______ of heat
conduction
conductors
transfer of energy by electromagnet waves
type of heat transfer does not require a material medium
can occur in a vacuum
how the sun warms the earth
radiation
transfer of heat by the bulk movement of fluids
Ex: warm air rises and cool air falls b/c of the difference in density
convection
functions whose value depends only on the position of the system and not on how the system got there
Ex: elevation is an example but distance traveled it not
state function
measure of the disorder, or randomness, of a system
units = Energy / temperature = J/K or cal/K (often J*K/mol)
entropy (S)
states that all spontaneous processes proceeding in an isolated system lead to an increase in entropy
second law of thermodynamics
states that the absolute entropy of a pure crystalline substance at absolute zero is zero
corresponds to a state of “perfect order” bc all of the atoms in the hypothetical state possess no kinetic energy and do not vibrate at all, so there is absolutely no randomness and no disorder in the spatial arrangement of the atoms
third law of thermodynamics
phases of matter w/ increasing entropy
solid < liquid < gas
Dissolution of particles in solution ________ the entropy
increases
Entropy _______ with temperature
increases
measure of the “heat content” of a system
Enthalpy (H)
equal to the heat absorbed or evolved by the system at constant pressure
∆H; change in enthalpy
change in enthalpy that corresponds to an endothermic process (one that absorbs heat)
*products at higher PE than reactants
+∆H
change in enthalpy that corresponds to an exothermic process (releases heat)
*products at lower PE than reactants
-∆H
the enthalpy change that would occur if one mole of that compound were formed directly from its elements in their standard states
Ex: formation of water = enthalpy change of the rxn between 1 mole of diatomic hydrogen and 1/2 mole of diatomic oxygen
standard enthalpy (or heat) of formation of a compound (∆Hºf)
∆Hºf of an element in its standards state
= zero
hypothetical enthalpy change that would occur if he reaction were carried out under standard conditions
i.e. when reactants in their standard states are converted to products under standard conditions (1 atm and usually 298 K)
standard heat of reaction (∆Hºrxn)
Equation: Standard heat of rxn (∆Hºrxn)
∆Hºrxn = (sum of ∆Hºf of products) - (sum of ∆Hºf of reactants)
*Note: multiply ∆Hºf by the stoichiometric coefficient; switch sign if compound is broken down instead of formed
states that if a reaction can be broken down into a series of steps, the enthalpy change for the overall net reaction is just the sum of the enthalpies of each step
consequence of enthalpy being a state function
Hess’s Law
describes the overall spontaneity of a reaction
Gibbs Free Energy (∆G)
Equation: Gibbs Free Energy
∆G = ∆H - T∆S
indicates a spontaneous reaction
- ∆G
indicates a nonspontaneous reaction
+ ∆G
indicates that system is not in a state of equilibrium
∆G = 0
-∆H and +∆S: rxn is _______ at all temperatures
spontaneous (-∆G)
+∆H and -∆S: rxn is _______ at all temperatures
not spontaneous (+∆G)
+∆H and +∆S: rxn is _______ only at high temperatures
spontaneous (-∆G)
-∆H and -∆S: rxn is _______ only at low temperatures
spontaneous (-∆G)
the amound of heat needed to raise the temperature of one mass unit of a substance by 1 degree Celsius
Specific Heat (c)
Equation: Amount of heat (q) gained or given off by a substance that changes in temp
q = mc∆T
amount of heat required to change the phase of one mass unit of a substance
heat of transformation (L)
Equation: amount of heat (q) gained or given off by a substance that cahnges phase
q = m * L
During a phase change, the _______ of a system stays constant
temperature
The thermal energy supplied to a system during a phase change is used to increase the ______ of the substance instead of the KE, which corresponds to temperature
Potential Energy (PE)
Breaking bonds ______ energy.
Requires
Melting, vaporization, and sublimation are all ________ phase changes
endothermic (require energy)
Forming bonds ________ energy
releases
Freezing, condensation, and deposition are all ________ phase changes
exothermic
corresponds to the energy needed to effect the phase change from ice to water (solid to liquid)
heat of fusion: ∆H(fusion)
corresponds to the energy needed to effect the phase change from water to vapor
heat of vaporization: ∆H(vap)
phase change from solid to liquid
fusion
melting
phase change from liquid to solid
freezing
phase change from liquid to gas
vaporization
phase change from gas to liquid
condensation
phase change from solid to gas
sublimation
phase change from gas to solid
deposition
point (on phase diagram) at which all three phases are in equilibrium (solid, liquid, and gas) at a certain temp and pressure unique to that substance
triple point
point (on a phase diagram) at which the differences between the properties of the liquid and gas phases disappear
at extremely high temperature and pressure
beyond this point, the substance exists as a supercritical fluid
critical point
