7: Thermochemistry

Question 1

system

Answer 1

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

Question 2

surrounding/enivornment

Answer 2

everything outside of system… boundary can be moved

Question 3

isolate system

Answer 3

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

Question 4

closed system

Answer 4

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

Question 5

open system

Answer 5

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

Question 6

first law of thermodynamics

Answer 6

∆U = Q - W

• ∆U is change in internal energy of system
• Q is heat added to system
• W is work done by the system

Question 7

isothermal processes

Answer 7

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)

Question 8

adiabatic processes

Answer 8

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)

Question 9

isobaric processes

Answer 9

occur when the pressure of the system is constant

Question 10

isovolumetric (isochoric) processes

Answer 10

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)

Question 11

coupling

Answer 11

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

Question 12

state functions

Answer 12

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)

Question 13

standard conditions

Answer 13

25ºC, 1 atm, 1 M

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

Question 14

standard state

Answer 14

most stable form of a substance is called standard state of that substance

• H₂ (g)
• H₂O (l)
• NaCl (s)
• O₂ (g)
• C (s, graphite)

Question 15

evaporation/vaporation

Answer 15

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

Question 16

boiling

Answer 16

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

Question 17

gas-liquid equilibrium

Answer 17

occurs when rate of evolration and condensation are same

Question 18

boiling point

Answer 18

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

Question 19

liquid-solid equilibrium

Answer 19

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

Question 20

gas-solid equilibrium

Answer 20

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

Question 21

phase diagram

Answer 21

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

Question 22

triple point

Answer 22

point at which 3 phase boundaries meet

Question 23

critical point

Answer 23

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

Question 24

heat vs. temperature

Answer 24

heat is a form of energy, while temperature is a measure of average kinetic energy of the particles in a system (related to enthalpy)

Question 25

heat (Q)

Answer 25

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

Question 26

heat (q) absorbed or released in a given process

Answer 26

q = mc∆T

• m=mass
• c=specific heat

Question 27

constant-pressure calorimeter

Answer 27

coffee-cup calorimeter

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

Question 28

constant-volume calorimeter

Answer 28

bomb calorimeter

• heats of certain reactions (like combustion) can be measured indirectly by assessing temperature change in a water bath around the reaction vessel
• q_system = -q_surroundings
• abdiatic process because insulation prevents heat exchange

Question 29

heating curves

Answer 29

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∆H_{phase change}
• solid-liquid boundary… use ∆H_fus
• liquid-gas boundary… use ∆H_vap

Question 30

specific heat (c)

Answer 30

amount of energy required to raise temperature of one gram of a substance by 1ºC

C_H2O(l) = 1 cal/g•K or 4.18 J/g•K

Question 31

heat capacity

Answer 31

heat capacity = mc

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

Question 32

enthalpy (H)

Answer 32

heat changes at constant pressure… state function

∆H_rxn = H_products - H_reactants

positive ∆H_rxn corresponds to endothermic process

negative ∆H_rxn corresponds to exothermic process

Question 33

standard enthalpy of formation

Answer 33

∆H_fº 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)

• ∆H_fº of an element in its standard state is zero

Question 34

standard enthalpy of reaction

Answer 34

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

∆Hº_rxn = Σ∆Hº_{f, products} - Σ∆Hº_{f, reactants}

Question 35

Hess’s law

Answer 35

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

Question 36

bond dissociation energies

Answer 36

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)

Question 37

standard heat of combustion

Answer 37

∆Hº_com 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

Question 38

second law of thermodynamics

Answer 38

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
• ∆S_universe = ∆S_system + ∆S_surroundings > 0
• ∆Sº_rxn = Σ∆Sº_{f, products} - Σ∆Sº_{f, reactants}

Question 39

entropy

Answer 39

measure of the spontaneous dispersal of energy at a specific temperature

∆S= Q_rev/T

Q_rev is heat gained/lost in a reversible process

Question 40

Gibb’s Free energy

Answer 40

∆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 E_a

Question 41

standard free energy

Answer 41

∆Gº_rxn is free energy change of reactions under standard state conditions

∆Gº_f 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

∆Gº_rxn = Σ∆Gº_{f, products} -Σ∆Gº_{f, reactants}

Question 42

free energy, K_eq, and Q

Answer 42

∆Gº_rxn = -RTlnK_eq

• once a reaction begins, the standard state conditions no longer apple , so for a reaction in progress… ∆G_rxn = ∆Gº_rxn + RTlnQ