7: Thermochemistry Flashcards Preview

MCAT Chem > 7: Thermochemistry > Flashcards

Flashcards in 7: Thermochemistry Deck (42):
1

system

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

2

surrounding/enivornment

everything outside of system... boundary can be moved

3

isolate system

system cannot exchange energy (heat/work) or matter with surroundings... ie. insulated bomb calorimeter

4

closed system

system can exchange energy but not matter with the surroundings... ie. a steam radiator

5

open system

system can exchange both energy and matter with the surroundings... ie. a pot of boiling water

6

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

7

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)

8

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)

9

isobaric processes

occur when the pressure of the system is constant

 

10

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)

11

coupling 

common method for supplying energy for nonspontaneous reactions is by coupling nonspontaneous reactions to spontaneous ones

12

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)

13

standard conditions

25ºC, 1 atm, 1 M

  • used for kinetics, equilibrium, and thermodynamics probelms 
  • different than STP (0ºC, 1 atm)

14

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)

15

evaporation/vaporation

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

16

boiling

specific type of vaporization that occurs above the boiling point of a liquid and involves vaporization through the entire volume of the liquid

17

gas-liquid equilibrium

occurs when rate of evolration and condensation are same

18

boiling point

temperature at which the vapor pressure of the liquid equals the ambient pressure

19

liquid-solid equilibrium

occurs when rates of fusion/melting equal rates of freezing/solidification/crrystallization

 

20

gas-solid equilibrium

  • sublimation: when a solid --> gas (ie. Dry ice)
  • deposition: gas --> solid 

21

phase diagram

  • lines on diagram are called lines of equilibrium or phase boundaries

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22

triple point

point at which 3 phase boundaries meet

23

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

24

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)

25

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 

26

heat (q) absorbed or released in a given process

q = mc∆T

  • m=mass
  • c=specific heat

27

constant-pressure calorimeter

coffee-cup calorimeter

  • temperature of the contents is measured to determine heat of reaction

28

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

29

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

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30

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

31

heat capacity

heat capacity = mc

  • energy required to raise any given amount of a substance by 1ºC

32

enthalpy (H)

heat changes at constant pressure... state function

∆Hrxn = Hproducts - Hreactants

positive ∆Hrxn corresponds to endothermic process

negative ∆Hrxn corresponds to exothermic process

33

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

34

standard enthalpy of reaction

∆Hºrxn is enthalpy change accompanying a reaction being carried out under standard conditions

∆Hºrxn = Σ∆Hºf, products - Σ∆Hºf, reactants

35

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

36

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) 

37

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

38

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

 

39

entropy

measure of the spontaneous dispersal of energy at a specific temperature 

∆S= Qrev/T

Qrev is heat gained/lost in a reversible process

 

40

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

41

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

42

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