Ch. 7: Thermochemistry (Complete) Flashcards
defn: system
the matter that is being observed – the total amount of reactants and products in a chemical reaction
defn: surroundings (environment)
everything outside of that system
how do you determine where to place the boundary between system and surroundings?
it depends on what phenomenon one is studyign
defn + example: isolated system
the system cannot exchange energy (heat and work) or matter with the surroundings
example: insulated bomb calorimeter
defn + example: closed system
the system can exchange energy (heat and work) but not matter with the surroundings
example: steam radiator
defn + example: open system
the system can exchange both energy (heat and work) and matter with the surroundings
example: pot of boiling water
defn: process
when a system experiences a change in one or more of its properties (such as concentrations of reactants or products, temperature, or pressure)
defn + first law of thermodynamics: isothermal processes
how does this appear on a P-V graph?
occur when the system’s temperature is constant
so U is constant, so
delta U = 0 so,
first law of thermodynamics simplifies to Q = W (the heat added to the system equals the work done by the system)
on P-V graph: hyperbolic
if the temperature of the system is constant, what does this imply about the total internal energy of the system?
the total internal energy of the system (U) is also constant throughout the process
defn + first law of thermodynamics: adiabatic processes
how does this appear on a P-V graph?
no heat is exchanged between the system and the environment
so the thermal energy of the system is constant throughout the process
when Q = 0, first law is delta U = - W (the change in internal energy of the system is equal to work done on the system)
on P-V graph: hyperbolic
defn + first law of thermodynamics: isobaric processes
how does this appear on a P-V graph?
occur when the pressure of the system is constant
no effect on the first law
flat line on P-V graph
defn + first law of thermodynamics: isovolumetric/isochoric processes
how does this appear on a P-V graph?
no change in volume
first law: delta U = Q (the change in internal energy is equal to the heat added to the system)
vertical line on a P-V graph
why is there no work performed in an isochoric process?
what implication does this have on the P-V graph?
the gas neither expands nor compresses
the area under the curve is zero (which represents the work done by the gas)
defn: spontaneous process
one that can occur by itself without having to be driven by energy from an outside source
how do we predict whether the process will be spontaneous or not?
by calculating the change in the Gibbs free energy (delta G) for a process
defn: coupling
a common method for supplying energy for nonspontaneous reactions by coupling nonspontaneous reactions and spontaneous ones
defn: state functions
certain macroscopic properties that describe the system in an equilibrium state
cannot describe the process, only useful for comparing equilibriums
defn: process functions
a way to quantitatively describe the pathway taken from one equilibrium state to another
what are the 2 most important process functions?
work (W) and heat (Q)
what are the 8 state functions + mnemonic for remembering them?
When I’m under PRESSURE and feeling DENSE, all I want to do is watch TV (temperature, volume) and get HUGS (enthalpy (H), internal energy (U), gibbs free energy, and entropy (S))
char (2): state functions
- when the state of a system changes from one equilibrium to another, one or more of these state functions will change
- they are independent of the path/process taken, but not necessarily independent of one another
defn: standard conditions
25 deg C (298 K), 1 atm pressure, and 1 M concentrations
defn: standard temperature and pressure
0 deg C (273 K) and 1 atm pressure
when are standard conditions used vs. STP used?
STANDARD conditions –> for kinetics, equilibrium, and thermodynamics
STP –> for ideal gas
defn: standard state
the most stable form of a substance under standard conditions
defn + symbols: standard enthalpy, entropy, and free energy changes
the changes in enthalpy, entropy, and free energy that occur when a reaction takes place under standard conditions
deltaHknot, deltaSknot, deltaGknot
what does the degree sign in these variables represent?
0, as the standard state is used as the “zero point” for all thermodynamic calculations
defn: phase diagrams
graphs that show the standard and nonstandard states of matter for a given substance in an isolated system, as determined by temperatures and pressures
defn + char (2): phase changes
solid <–> liquid <–> gas
- reversible
- an equilibrium of phases will eventually be reached at any given combo of temp and pressure
what are phase equilibria analogous to? why?
analogous to the dynamic equilibria of reversible chemical reactions
why? the concentrations of reactants and products are constant because the rates of the forward and reverse reactions are equal
defn: evaporation/vaporization
when molecules in the liquid phase near the surface of the liquid have enough kinetic energy to leave the liquid phase and escape into the gaseous phase
what happens every time the liquid loses a high-energy particle to the gas phase? what does this mean about evaporation?
the temperature of the remaining liquid decreases
evaporation is an endothermic process for which the heat source is liquid water
defn: boiling
a specific type of vaporization that occurs only under certain conditions
the rapid bubbling of the entire liquid with rapid release of the liquid as gas particles
what is the difference between when evaporation and boiling can occur?
evaporation: happens in all liquids at all temperatures
boiling: can only occur above the boiling point of a liquid and involves vaporization throughout the entire volume of the liquid
defn: condensation
in a covered or closed container, the molecules that are escaping the liquid phase are trapped above the solution
these molecules exert a countering pressure, which forces some of the gas back into the liquid phase
what 2 physical conditions facilitate condensation?
- lower temperature
- higher pressure
defn: vapor pressure of the liquid
the pressure that the gas exerts over the liquid at equilibrium
why does vapor pressure increase as temperature increases?
because more molecules have sufficient kinetic energy to escape into the gas phase
defn: boiling point
the temperature at which the vapor pressure of the liquid equals the ambient pressure
3 aka: ambient
external
applied
incident
describe how entropy changes as the temperature of a solid increases
what does this mean in layman’s terms
the availability of energy microstates increases as the temperature of the solid increases
the molecules have greater freedom of movement and the energy disperses
when heat is applied, what happens to the vibrational motions of the atoms or molecules of a solid from their equilibrium position?
the vibrational motions increase when heat is applied
defn: fusion/melting
if atoms or molecules in the solid phase absorb enough energy, the 3-D structure of the solid will break down, and the atoms or molecules will escape into the liquid phase
defn + 2 akas: freezing
the reverse of melting, the transition from liquid to solid
aka: solidification, crystallization
defn: melting point or freezing point
the temperature at which melting or freezing occurs, respectively
what types of materials have distinct, precise melting points? a larger range of melting/freezing points (+ why?)?
distinct, precise: pure crystalline solids
large range: amorphous solids: glass, plastic, chocolate, candle wax –> because of their less ordered molecular structure
defn: sublimation
when a solid goes directly into the gas phase
defn: deposition
the transition from gas to solid
defn + process: cold finger
a device that may be used to purify a product that is heated under reduced pressure, causing it to sublime
the desired product is more volatile than the impurities, and so impurities are left in the solid state and the gas is purer than the original prodcut
the gas then deposits onto the cold finger, which has cold water flowing through it, yielding a purified solid product that can be collected
defn + aka + func: lines of equilibrium
aka: phase boundaries
the line on a phase diagram
func: indicate the temperature and pressure values for the equilibria between phases
defn + symbol: standard enthalpy of formation
the enthalpy required to produce one mole of a compound from its elements in their standard states
what is the standard enthalpy of formation of an element in its standard state?
by definition, zero
defn + symbol: standard enthalpy of a reaction
the enthalpy change accompanying a reaction being carried out under standard conditions
how do you calculate the standard enthalpy of a reaction?
take the difference between the sum of the standard heats of formation for the products and the sum of the standard heats of formation of the reactants
how do we describe the standard enthalpy of a reaction in terms of bonds being broken and formed?
which also equals total energy absorbed - total energy released
defn + symbol: standard heat of combusion
the enthalpy change associated with the combustion of a fuel
eqn: change in entropy
where delta S = change in entropy
qrev = the heat that is gained or lost in a reversible process
T = temperature in Kelvin
what is the equation that represents that the entropy of the universe is increasing
this equation is > 0
eqn: standard entropy change for a reaction
the standard entropy change for a reaction can be calculated using the standard entropy of the reactants and products
eqn: change in Gibbs free energy
where T is the temperature in kelvin
TdeltaS = the total amount of energy that is absorbed by a system when its entropy increases reversibly
defn + symbol: change in standard free energy reaction
the free energy change of reactions measured under standard state conditions
defn + symbol: change in standard free energy compound
the free energy change that occurs when 1 mole of a compound in its standard state is produced from its respective elements in their standard states under standard state conditions
equation: change in standard free energy of a reaction
what is an alternative equation for finding the change in standard free energy for a reaction?
where R is the ideal gas constant
Te is the temperature in kelvin
K = Keq is the equilibrium constant
how and why does this equation change once a reaction begins?
we relate deltaGrxn to the reaction quotient, Q because the standard state conditions no longer apply once a reaction begins
this equation can also be written as the following. what does the ratio of Q/Keq tell us?
Q/Keq < 1 = Q<Keq = ln negative = free energy change negative = reaction proceeds spontaneously forward until equilibrium
Q/Keq > 1 = Q>Keq = ln positive = free energy change positive = reaction proceeds spontaneously reverse until equilibrium
Q/Keq = 1; Q = Keq; reaction is at equilibrium, free energy change is zero
