Mcat - Gem Chem. Ch 7

Front 1

system

Back 1

- matter being observed

- total amount of products and reactants in the reaction

- characterized by whether or not they can exchange heat or matter with surroundings

Front 2

surroundings

Back 2

- "environment"

- everything outside of the system

Front 3

isolated system

Back 3

- system cannot exchange energy or matter with the surroundings

Front 4

Closed system

Back 4

- system can exchange energy but not matter with the surroundings

Front 5

Open system

Back 5

- system can exchange both energy and matter with the surroundings

Front 6

Process

Back 6

- when a system experiences a change in one or more of its properties

Front 7

first law of thermodynamics

Back 7

- U = Q - W

U = change in internal energy of the sys.

Q = heat added to the system

W = work done by the sys.

Front 8

isothermal processes

Back 8

occur when the systems temp is constant

- implies that total energy of the sys. is constant

- U = 0, Q=W

- hyperbolic curve on PV graph

- work = area under graph

Front 9

Adiabatic processes

Back 9

- no heat exchange between sys. and environment, temp is not constant

- Q = 0 , U = -W

- appears hyperbolic

Front 10

isobaric processes

Back 10

- occurs when the pressure of a system is constant

- appears as a flat line on a P-V graph (constant pressure = zero slope)

Front 11

isovolumetric (isochoric) processes

Back 11

- no change in volume

- no work is performed

- W=0, U=Q

- vertical line on P-V graph (V=0)

Front 12

spontaneous process

Back 12

- can occur by itself without having to be driven by energy from an outside source

- may not happen quickly

- may not go to completion

Front 13

coupling

Back 13

- method for supplying energy for nonspontaneous reactions

Front 14

state functions

Back 14

- describe system in equilibrium state

- pressure, density, temperature, volume, enthalpy, internal energy, gibbs free energy, entropy

-"When I'm under pressure and feeling dense, all I want to do is watch TV and get HUGS"

Front 15

process functions

Back 15

- describes pathway taken from one equilibrium state to another

- Ex: Work (W) and Heat(Q)

Front 16

Standard Conditions

Back 16

- for measuring enthalpy, entropy, and gibbs free energy

25 deg C, 298 K, 1 atm pressure and 1 M concentrations

Front 17

Phase diagrams

Back 17

- graphs that show the temp and pressure at which a substance will be thermodynamically stable in a particular phase

- show temp and pressures of the phase at equilibrium (lines of equilibrium/ phase boundaries)

Front 18

Phase changes

Back 18

- solid to liquid to gas

- reversible

- equilibrium of phases will eventually be reached at any given combination of temp and pressure

Front 19

evaporation / vaporization

Back 19

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

- endothermic process

- heat source = liquid water

Front 20

Boiling

Back 20

- rapid bubbling of the entire liquid with rapid release of the liquid as gas particles

Front 21

condensation

Back 21

-escaping gas molecules are trapped above solution which forces the gas to go back into the liquid phase

- @ lower temp, or @ higher pressure

Front 22

boiling point

Back 22

- temp where the vapor pressure of the liquid equals the external pressure

Front 23

energy of microstates

Back 23

increases as temp increases

Front 24

fusion

Back 24

- "melting"

- solid to liquid

- occurs at melting point

Front 25

solidification

Back 25

- "crystallization"

- "freezing"

- occurs at freezing point

Front 26

sublimation

Back 26

- solid to a gas

Front 27

deposition

Back 27

- gas to solid

Front 28

triple point

Back 28

- point where all 3 phase boundaries meet in the phase diagram

Front 29

critical point

Back 29

- phase boundary btwn liquid and gas ends

Front 30

Temperature (T)

Back 30

- way to scale how hot or cold something is

Front 31

thermal energy

Back 31

- average kinetic energy of the particles in a substance

Front 32

Heat (Q)

Back 32

- transfer of energy from one substance to another as a result of their differences in temperature

- process function

Front 33

Zeroth law of thermodynamics

Back 33

- objects are in thermal equilibrium when their temperatures are equal

Front 34

endothermic

Back 34

- process in which the system absorbs heat

Front 35

exothermic

Back 35

- system releases heat

Front 36

Calorimetry

Back 36

- process of measuring transferred heat

Front 37

specific heat

Back 37

- amount of energy required to raise the temperature of one gram of a substance by one degree celcius

Front 38

heat capacities

Back 38

- mass * specific heat

Front 39

heat absorbed / released eqn

Back 39

= q=mcΔT

Front 40

Heating curve

Front 41

Heating curves show...

Back 41

- phase change rxns do not undergo changes in temp.

Front 42

heat of fusion

Back 42

- used to determine the heat transferred during the phase change at solid-liquid boundary

- Ex: solid to liquid = + b/c heat must be added

- q=mL

m= moles

L = latent of heat

Front 43

heat of vaporization

Back 43

- at liquid-gas boundary

- + indicates the addition of heat

- q= mL

m = moles

L = latent of heat

Front 44

enthalpy (H)

Back 44

- equivalent to heat under constant pressure

- state function

- change in enthalpy is equal to the heat transferred into or out of the system at constant pressure

- property of the equilibrium state

ΔH = H (products) - H (reactants)

- -H = exothermic

- + H = endothermic

Front 45

Standard enthalpy of formation

Back 45

- enthalpy required to produce one moles of a compound from its elements in their standard states

Front 46

standard heat of reaction

Back 46

- enthalpy change accompanying a reaction being carried out under standard conditions

Hreaction = SHproducts - SHreactants

Front 47

Hess's law

Back 47

- enthalpy changes of reactions are additive

Front 48

bond dissociation energies

Back 48

- average energy that is required to break a particular type of bond between atoms in the gas phase

- endothermic process

- units: kJ/mol of bonds broken

ΔH = ∑ ΔH(bonds broken) - ∑ ΔH(bonds formed)

Front 49

standard heat of combustion

Back 49

- enthalpy change associated with he combustion off a fuel

Front 50

entropy

Back 50

- measure of the spontaneous dispersal of energy at a specific temperature

ΔS = Q/T

S(universe) = S (sys) + S(surroundings)

Front 51

second law of thermodynamics

Back 51

- energy spontaneously disperses from being localized to becoming spread out if it is not hindered from doing so

Front 52

Gibbs Free Energy

Back 52

- state function

- measure of the change in enthalpy and the change in entropy as a system undergoes a process

- indicates if a rxn is spontaneous or nonspontaneous

- change in free energy is the max amount of energy released by a process that is available to perform useful work

- +G = spontaneous rxn

- -G = nonspontaneous rxn

- G = 0 = in equilibrium

- ΔG = temp dependent when H and S have the same sign

Front 53

Standard free energy change

Back 53

- free energy change that occurs when 1 mol of a compound in its standard state is produced from its respective elements in their standard states under standard conditions

- ΔG (rxn) = -RTlnK(eq)

- the more negative, the more spontaneous

Front 54

Free energy of the reaction

Back 54

ΔG (rxn) = sum ΔG (prod) - sum ΔG (reactants)

Front 55

Free energy change for a reaction in progress

Back 55

ΔGrxn = ΔG(0)rxn + RT ln(Q) = RT ln (Q/Keq)

Q/Keq < 1= negative enrgy change = rxn will spontaneously move forward until equilibrium

Q/Keq > 1 = positive energy change = rxn will spontaneously move in the reverse direction until equilibrium is reached