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: triple point
the point at which the three phase boundaries meet
the temperature and pressure at which the three phases exist in equilibrium
what extends indefinitely from the triple point?
the phase boundary that separates the solid and the liquid phases
defn: critical point
the temperature and pressure above which there is no distinction between the phases
what terminates at the critical point?
the phase boundary between the liquid and gas phases
what type of substance exists beyond the critical point?
supercritical fluids
why do supercritical fluids form? (2)
- as a liquid is heated in a closed system, its density decreases and the density of the vapor sitting above it increases
- the critical point is the point at which the two densities become equal and there is no distinction between the phases
what is the heat of vaporization at the critical point and beyond?
beyond = for all temperatures and pressures above the critical point
0!
defn: temperature
related to the average kinetic energy of the particles of a substance
the way we scale how hot or cold something is
what is the relationship between the thermal energy (enthalpy) of a substance and the temperature of a substance?
when a substances thermal energy increases, its temperature also increases
why can there be no temperature below 0K (absolute zero)?
by definition, the system is said to be unable to lose any more heat energy
defn: heat (Q)
the transfer of energy from one substance to another as a result of their differences in temperature
defn: zeroth law of thermodynamics
objects are in thermal equilibrium only when their temperatures are equal
defn: first law of thermodynamics
the change in the total internal energy of a system is equal to the amount of heat transferred to the system minus the amount of work done by the system
defn: endothermic vs. exothermic
ENDOthermic = deltaQ > 0 = processes in which the system absorbs heat
EXOthermic = delta Q < 0 = processes in which the system releases heat
2units of heat + conversion
1 cal = 4.184 J
defn: enthalpy (deltaH)
is equivalent to heat (Q) under constant pressure
does energy move from warmer to cooler or cooler to warmer substance when substances of different temperatures are in thermal contact?
warmer to cooler
what happens when a substance undergoes and endothermic or exothermic reaction?
heat energy will be exchanged between the system and the environment
defn: calorimetry
the process of measuring transferred heat
what are the two main types of calorimetry? + an example of each
- constant-pressure calorimetry –> coffee-cup calorimeter
- constant-volume calorimetry –> bomb calorimeter
defn: specific heat
the amount of energy required to raise the temperature of one gram of a substance by one degree Celsius or kelvin
defn: heat capacities
the product mc (mass times specific heat)
process + explanation (3): constant-pressure calorimeter (think of the coffee-cup calorimeter)
- an insulated container covered with a lid, filled with a solution in which a reaction or some physical process (such as dissolution) is occurring
- the incident (atmospheric) pressure remains constant throughout the process and the temperature can be measured as the reaction progresses
- there should be sufficient thermal insulation (i.e. Styrofoam) to ensure that the heat being measured is an accurate representation of the reaction, without gain or loss of heat to the environment
what are 3 common commercial applications of these same principles?
- home insulation
- padded clothing
- certain food containers such as thermoses
what is another term for a bomb calorimeter (application of constant-volume calorimetry)?
decomposition vessel
what is happening in constant-volume calorimetry (a bomb calorimeter)? (5)
- a sample of matter (typically a hydrocarbon) is placed in the steel decomposition vessel, which is then filled with almost pure oxygen gas
- the decomposition vessel is then placed in an insulated container holding a known mass of water
- the contents of the decomposition vessel are ignited by an electric ignition mechanism
- the material combusts (burns) in the presence of the oxygen, and the heat that evolves is the heat of the combustion reaction
- no work is done in an isovolumetric process, so Wcalorimeter = 0
why can a bomb calorimeter be considered isolated from the rest of the universe? what does this imply about what counts as the SYSTEM and what counts as the SURROUNDINGS?
because of the insulation
system: sample + oxygen + steel vessel
surroundings: water
Why is Qcalorimeter of a bomb calorimeter 0?
because no heat is exchanged between the calorimeter and the rest of the universe
does the temperature of a compound change when it is undergoing a phase change? what type of graph demonstrates this?
no!
The temperature only changes within a phase
this is demonstrated by heating curves
we know intuitively that heat must continue to be added in order for the whole solid to melt, even though the temperature is constant during a phase change, so where does this heat go?
the solid absorbs energy, which allows particles to overcome the attractive forces that hold them in a rigid, 3-D arrangement
the opposite is also true: removing heat will cause a formation of the rigid, 3-D arrangement
what values do we focus on during phase changes?
enthalpy
defn: enthalpy (heat) of fusion
will this be positive or negative for each direction of phase change? why?
used to determine the heat transferred during the phase change when transitioning at the solid-liquid boundary
from solid to liquid: this will be positive because heat must be added
from liquid to solid: this will be negative because this must be removed
defn: enthalpy (heat) of vaporization
used at the liquid-gas boundary
sign convention follows a similar pattern to enthalpy of fusion
defn + unit: latent heat
a general term for the enthalpy of an isothermal process
units: cal/g
how do you find the total amount of heat needed to cross multiple phase boundaries?
simply a summation of the heats for changing the temperature of each of the respective phases and the heats associated with phase changes
defn: change in enthalpy
equal to the heat transferred into or out of the system at constant pressure
can enthalpy be measured directly?
no only change in enthalpy
defn + symbol: standard enthalpy of formation
the enthalpy required to produce one mole of a compound from its elements in their standard states
defn: standard state
the most stable physical state of an element or compound at 298 K and 1 atm
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
defn: Hess’s law
enthalpy changes of reactions are additive
Hess’s law applies to ANY state function including entropy and Gibbs free energy
what is the reaction between the enthalpy change for the forward of any reaction and the reverse of the same reaction?
the enthalpy change for the reverse of any reaction has the same magnitude, but opposite sign, as the enthalpy change for the forward reaction
defn + aka + unit: bond dissociation energies
aka: bond enthalpies
the average energy that is required to break a particular type of bond between atoms in the gas phase (average of the bond energies for the same bond in many different compounds)
unit: kJ/mol of bonds broken
is bond dissociation an endothermic or exothermic process? what about bond formation?
dissociation = endothermic
formation = exothermic
what is the relationship between bond formation and bond breaking?
bond formation has the same magnitude of energy but is negative rather than positive (energy is released when bonds are formed)
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
why are combustion reactions the ideal processes for measuring enthalpy change?
because measurements of enthalpy change require a reaction to be spontaneous and fast
what are 3 common oxidants in combustion reactions?
- atmospheric oxygen
- diatomic flourine
- hydrogen gas + chlorine gas will combust
process + defn: entropy
process: energy of some form goes from being localized or concentrated to being spread out or dispersed
defn: the measure of the spontaneous dispersal of energy at a specific temperature; how MUCH energy is spread out or how WIDELY spread out energy becomes, in a process
defn: second law of thermodynamics
energy spontaneously disperses from being localized to become spread out if it is not hindered from doing so
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
how does entropy increase or decrease as energy is distributed into or out of a system?
energy distributed INTO a system at a given temp: entropy increases
energy distributed OUT of a system at a given temp: entropy decreases
does the fact that energy will spontaneously disperse mean that energy can never be localized or concentrated?
no but the concentration of energy will rarely happen spontaneously in a closed system because work must usually be done to concentrate energy
what is an example of something that works against the direction of spontaneous heat flow by concentrating energy? how?
refrigerator!
they counteract the flow of heat from the “warm” exterior of the fridge to the “cool” interior thereby “concentrating” energy outside of the system in the surroundings
fridges consume a lot of energy to accomplish this movement of energy against a temp gradient
what does it mean that the second law of thermodynamics has been described as time’s arrow?
there is a unidirectional limitation on the movement of energy by which we recognize before and after or new and old
energy in a closed system will spontaneously spread out, and entropy will increase if it is not hindered from doing so
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
defn: change in Gibbs free energy, deltaG
a measure of the change in enthalpy and the change in entropy as a system undergoes a process
indicates whether a reaction is spontaneous or nonspontaneous
the change in the free energy is the maximum amount of energy released by a process occurring at constant temperature and pressure that is available to perform useful work
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
how can we visually think of Gibbs free energy?
it is like a valley between two hills
any system (including chem reactions) will move in whichever direction results in a reduction of the free energy of the system
the bottom of the valley represents equilibrium and the sides of the hill represent the various points in the pathway toward or away from equilibrium
defn + movement in relation to equilibrium position + gibbs free energy change: exergonic vs. endergonic
EXERGONIC = when a system releases energy = movement TOWARD equilibrium = decrease in Gibbs free energy = spontaneous
ENDERGONIC = movement AWAY form the equilibrium position = increase in Gibbs free energy = nonspontaneous
once the system is at equilibrium, what is the change in free energy?
0!
summarize the meanings of delta G being negative, positive, and zero
NEGATIVE = reaction is spontaneous
POSITIVE = reaction is nonspontaneous
ZERO = the system is in a state of equilibrium: deltaH = TdeltaS
why is the temperature variable in Gibbs free energy always positive?
it is in Kelvin!
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
what is the standard free energy of formation for any element under standard state conditions?
zero
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 does this alternative equation allow us to get qualitative assessments of the spontaneity of the reaction?
the greater the value of Keq = the more positive the value of its natural log = the more negative the standard free energy change = the more spontaneous the reaction
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
