7: Thermochemistry Flashcards Preview

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Flashcards in 7: Thermochemistry Deck (42):


matter that is being observed... the total amount of reactants and products in a chemical reaction



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)



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)



some of molecules near surface of liquid have enough kinetic energy to leave liquid phase and escape into gaseous phase



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

A image thumb

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

A image thumb

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




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