7: Thermochemistry Flashcards
system
matter that is being observed… the total amount of reactants and products in a chemical reaction
surrounding/enivornment
everything outside of system… boundary can be moved
isolate system
system cannot exchange energy (heat/work) or matter with surroundings… ie. insulated bomb calorimeter
closed system
system can exchange energy but not matter with the surroundings… ie. a steam radiator
open system
system can exchange both energy and matter with the surroundings… ie. a pot of boiling water
first law of thermodynamics
∆U = Q - W
- ∆U is change in internal energy of system
- Q is heat added to system
- W is work done by the system
isothermal processes
occur when the system’s temperature is constant… total energy of system is constant so ∆U=0
- Q=W (heat added to the system equals work done by the system)
adiabatic processes
occur when no heat is exhcange between the system and environment… Q=0
∆U=-W (change in internal energy of the system is equal to work done on the system)
isobaric processes
occur when the pressure of the system is constant
isovolumetric (isochoric) processes
experience no change in volume because gas neither expands nor compresses… no work is performed
∆U=Q (change in internal energy is equal to the heat added to the system)
coupling
common method for supplying energy for nonspontaneous reactions is by coupling nonspontaneous reactions to spontaneous ones
state functions
properties of system in an equilibrium state… independent of pathway… pressure, desnsity, TV HUGS
- presure (P)
- density
- temperature (T)
- volume (V)
- enthalpy (H)
- internal energy (U)
- Gibbs free energy (G)
- entropy (S)
standard conditions
25ºC, 1 atm, 1 M
- used for kinetics, equilibrium, and thermodynamics probelms
- different than STP (0ºC, 1 atm)
standard state
most stable form of a substance is called standard state of that substance
- H2 (g)
- H2O (l)
- NaCl (s)
- O2 (g)
- C (s, graphite)
evaporation/vaporation
some of molecules near surface of liquid have enough kinetic energy to leave liquid phase and escape into gaseous phase
boiling
specific type of vaporization that occurs above the boiling point of a liquid and involves vaporization through the entire volume of the liquid
gas-liquid equilibrium
occurs when rate of evolration and condensation are same
boiling point
temperature at which the vapor pressure of the liquid equals the ambient pressure
liquid-solid equilibrium
occurs when rates of fusion/melting equal rates of freezing/solidification/crrystallization
gas-solid equilibrium
- sublimation: when a solid –> gas (ie. Dry ice)
- deposition: gas –> solid
phase diagram
- lines on diagram are called lines of equilibrium or phase boundaries
triple point
point at which 3 phase boundaries meet
critical point
point where phase boundary between liquid and gas phases terminates… temperature and pressure at which densities of liquid and gas become equal and there is no distinction between phases
- heat of vaporization is 0
heat vs. temperature
heat is a form of energy, while temperature is a measure of average kinetic energy of the particles in a system (related to enthalpy)
heat (Q)
transfer of energy as a result of their differences in temperature in J or cal
- process function
- endothermic when ∆Q>0
- exothermic when ∆Q<0
- enthalpy (∆H) is equivalent to heat under constant pressure
heat (q) absorbed or released in a given process
q = mc∆T
- m=mass
- c=specific heat
constant-pressure calorimeter
coffee-cup calorimeter
- temperature of the contents is measured to determine heat of reaction
constant-volume calorimeter
bomb calorimeter
- heats of certain reactions (like combustion) can be measured indirectly by assessing temperature change in a water bath around the reaction vessel
- qsystem = -qsurroundings
- abdiatic process because insulation prevents heat exchange
heating curves
show that phase change reactions do not undergo changes in temperature so you cannot use q=mc∆T during the interval because ∆T=0
- use q=mc∆T in phase
- during phase changes, use q=m∆Hphase change
- solid-liquid boundary… use ∆Hfus
- liquid-gas boundary… use ∆Hvap
specific heat (c)
amount of energy required to raise temperature of one gram of a substance by 1ºC
CH2O(l) = 1 cal/g•K or 4.18 J/g•K
heat capacity
heat capacity = mc
- energy required to raise any given amount of a substance by 1ºC
enthalpy (H)
heat changes at constant pressure… state function
∆Hrxn = Hproducts - Hreactants
positive ∆Hrxn corresponds to endothermic process
negative ∆Hrxn corresponds to exothermic process
standard enthalpy of formation
∆Hfº is enthalpy required to produce 1 mole of a compound from its elements in their standard states (most stable state of an element at 298 K and 1 atm)
- ∆Hfº of an element in its standard state is zero
standard enthalpy of reaction
∆Hºrxn is enthalpy change accompanying a reaction being carried out under standard conditions
∆Hºrxn = Σ∆Hºf, products - Σ∆Hºf, reactants
Hess’s law
enthalpy changes of reactions are additive because enthalpy is a state funcion
∆H = ∆H1 + ∆H2 + ∆H3 + …
- applies to ANY state function… like entropy and Gibbs free energy
bond dissociation energies
average energy required to break a particular type of bond between atoms in the gas phase (endothermic)
- bond formation has same magnitude of energy but different sign (exothermic)
standard heat of combustion
∆Hºcom is enthalpy change associated with combustion of a fuel
- usually hydrocarbon + O2 –> CO2 + H2O
- the larger the alkane reactant, the more numerous the combustion products
second law of thermodynamics
energy spontaneously dispreses from being localized to becoming spread out if it is not hindered from doing so
- concentration of energy will rarely happen spontaneously in a closed system… work must be done to concentrate energy
- ∆Suniverse = ∆Ssystem + ∆Ssurroundings > 0
- ∆Sºrxn = Σ∆Sºf, products - Σ∆Sºf, reactants
entropy
measure of the spontaneous dispersal of energy at a specific temperature
∆S= Qrev/T
Qrev is heat gained/lost in a reversible process
Gibb’s Free energy
∆G = ∆H - T∆S
- system moves in whichever direction results in a reduction of the free energy of the system
- ∆G < 0 is spontaneous… system is exergonic
- ∆G > 0 is nonspontaneous… system is endergonic
- ∆G = 0, system is in a state of equilibrium
- ∆H = T∆S
*thermodynamic spontaneity has no bearing on kintetics of reaction… rate of reaction depends on Ea
standard free energy
∆Gºrxn is free energy change of reactions under standard state conditions
∆Gºf is free energy change that occurs when 1 mole of a compound in its stardard state is produced from its respective elements in their standard states
∆Gºrxn = Σ∆Gºf, products -Σ∆Gºf, reactants
free energy, Keq, and Q
∆Gºrxn = -RTlnKeq
- once a reaction begins, the standard state conditions no longer apple , so for a reaction in progress… ∆Grxn = ∆Gºrxn + RTlnQ
