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General Chemistry > ch 7 - Thermochemistry > Flashcards

ch 7 - Thermochemistry Flashcards

1
Q

system

A

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

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2
Q

surroundings or environment

A

everything outside of the system

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3
Q

characterizations of systems

A

isolated: system cannot exchange energy (heat and work) or matter with the surroundings; closed: system can exchange energy (heat and work) but not matter with the surroundings, ex is a steam radiator; open: system can exchange both energy (heat and work) and matter with the surroundings (pot of boiling water)

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4
Q

process

A

when a system experiences a change in one or more of its properties (such as concentrations of reactants or products, temperature, or pressure

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5
Q

first law of thermodynamics

A

delta U = Q - W; delta U = change in internal energy of the system; Q = heat added to the system; W = work done by the system

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6
Q

isothermal processes

A

system’s temp is constant which implies internal energy of the system (U) is also constant; in this case delta U = 0 and Q = W; P-V graph shows as hyperbolic and work is the area under the graph

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7
Q

Adiabatic processes

A

occur when no heat is exchanged between the system and the environment; thermal energy of the system is constant throughout process. When Q = 0, delta U = -W (change in internal energy of the system is equal to work done on the system); appears hyperbolic on P-V (pressure-volume) graph

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8
Q

isobaric processes

A

occur when the pressure of the system is constant; do not alter the first law; appears as a flat, horizontal line on the P-V (pressure-volume) graph

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9
Q

isovolumetric (isochoric) processes

A

experience no change in volume; no work is performed. W = 0, delta U = Q (change in internal energy is equal to the heat added to the system); vertical line on P-V graph

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10
Q

spontaneous process

A

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

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11
Q

state functions

A

certain macroscopic properties that describe a system in equilibrium state; pressure (P), density, temp (T), volume (V), enthalpy (H), internal energy (U), Gibbs free energy (G), entropy (S)

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12
Q

process functions

A

pathway taken from one equilibrium state to another, quantitatively. The most important of these are work and heat

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13
Q

standard conditions

A

used for measuring the enthalpy, entropy, and Gibbs free energy changes of a reaction: 25 degrees C (298 K), 1 atm pressure, and 1 M concentrations; used for kinetics, equilibrium and thermodynamics problems

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14
Q

standard temp and pressure (STP)

A

used for ideal gas calculations: temp is 0 degrees C (273 K) and pressure is 1 atm.

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15
Q

standard state

A

most stable state of a substance under standard conditions

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16
Q

standard enthalpy (delta H degree sign), standard entropy (delta S degree sign), standard free energy changes (delta G degree sign)

A

change in enthalpy, entropy, and free energy that occur when a reaction takes place under standard conditions; degree sign represents zero, as the standard state is used as the “zero point” for all thermodynamic calculations

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17
Q

Phase diagrams

A

graphs that show the standard and nonstandard states of matter for a given substance in an isolated system, as determined by temps and pressures

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18
Q

evaporation

A

also vaporization: liquid to gas; every time liquid loses a high energy particle, temp of remaining liquid decreases; endothermic process for which the heat source is the liquid water

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19
Q

condensation

A

gas to liquid; facilitated by lower temp or higher pressure

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20
Q

melting or fusion

A

transition from solid to liquid

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21
Q

solidification, crystallization, freezing

A

liquid to solid

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22
Q

sublimation

A

solid directly to gas phase

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23
Q

deposition

A

from gas to solid

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24
Q

temperature

A

T - related to average kinetic energy of particles of a substance; way we scale how hot or cold something is; average kinetic energy is related to thermal energy (enthalpy); what is hot does not necessarily have a greater thermal energy but when thermal energy increases in a substance so does temp

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25
Q

Heat (Q)

A

the transfer of energy from one substance to another as a result of their differences in temperature; process function; processes that absorb heat are endothermic (delta Q>0) those that release heat are exothermic (delta Q<0); unit of heat is joule (J) or calorie (cal) - one cal = 4.184 J

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26
Q

zeroth law of thermodynamics

A

implies that objects are in thermal equilibrium only when their temps are equal

27
Q

calorimetry

A

process of measuring transferred heat; two basic types are constant-pressure calorimetry and constant-volume calorimetry

28
Q

equation for heat absorbed or released in a given process

A

q = mc deltaT (q = mcAt); m = mass; c = specific heat of the substance; delta T = change in temp (C or K)

29
Q

specific heat (c)

A

the amount of energy required to raise the temp of one gram of a substance by one degree Celsius (or one Kelvin)

30
Q

specific heat of H2O

A

c sub H2O = 1 cal/g x K

31
Q

heat capacity

A

product of mc (mass times specific heat)

32
Q

constant-pressure calorimeter

A

an insulated container covered with a lid and filled with a solution in which a reaction or some physical process, such as dissolution, is occurring; incident pressure (atmospheric pressure) remains constant throughout process

33
Q

constant-volume calorimeter

A

example is a bomb calorimeter, also called decomposition vessel: a sample of matter, like hydrocarbon, is placed in a steel decomposition vessel that is then filled with pure oxygen gas; vessel is placed in an insulated container holding a known mass of water, material is ignited and heat that evolves is heat of combustion reaction. no work is done in isovolumetric process

34
Q

enthalpy (or heat) of fusion (delta H sub fus)

A

used at the solid-liquid boundary to determined the heat transferred during phase change; when transitioning from solid to liquid, the change in enthalpy will be positive because heat must be added; when transitioning to a liquid from a solid, the change in enthalpy will be negative because heat must be removed. q = mL; m = mass and L = latent heat

35
Q

enthalpy (or heat) of vaporization (delta H sub vap)

A

used at the liquid-gas boundary to determine heat transferred during phase change. q = mL; m = mass and L = latent heat, a general term for enthalpy of an isothermal process, given in units cal/g

36
Q

enthalpy (H)

A

potential energy; state function; used to express heat changes at a constant pressure; we can calculate the change in enthalpy of a system that has undergone a process by comparing final to initial regardless of path taken

37
Q

enthalpy change of a reaction

A

delta H sub rxn = H sub products - H sub reactants; positive delta H sub rxn corresponds to an endothermic process; negative corresponds to exothermic

38
Q

standard enthalpy of formation (delta H, degree sign, sub f)

A

enthalpy required to produce one mole of a compound from its elements in their standard states. (standard state refers to the most stable physical state of an element or compound at 298 K and 1 atm); delta H, degree sign, sub f of an element in its standard state, by definition, is zero

39
Q

standard heat of reaction (delta H degree sign, sub rxn)

A

the enthalpy change accompanying a reaction being carried out under standard conditions; calculated by taking the difference between the sum of the standard heats of formation for the products and the sum of the standard heats of formation of the reactants: delta H, degree sign, sub rxn = sum of delta H, degree sign, sub f, products - sum of delta H, degree sign, sub f, reactancts

40
Q

Hess’s Law

A

states that enthalpy changes of reactions are additive; when thermochemical changes (chemical equations for which changes are known) are added to give the net equation for a reaction, the corresponding heats of reaction are also added to give the net heat of reaction

41
Q

bond enthalpies

A

also called bond dissociation energies; average energy that is required to break a particular type of bond between atoms in the gas phase; given in units kJ/mol of bonds broken; Hess’s Law can be expressed in terms of these

42
Q

equation for enthalpy change associated with a reaction

A

delta H, degree sign, sub rxn = sum of delta H sub bonds broken - sum of delta H sub bonds formed = total energy absorbed - total energy released

43
Q

standard heat of combustion (delta H, degree sign, sub comb)

A

the enthalpy change associated with the combustion of a fuel

44
Q

second law of thermodynamics

A

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

45
Q

entropy

A

the measure of the spontaneous dispersal of energy at a specific temp; how much energy is spread out or how widely spread out energy becomes, in a process

46
Q

equation for change in entropy

A

delta S = Q sub rev/T; delta S is change in entropy; Q sub rev is heat that is gained or lost in reversible process; T = temp in kelvin; units are usually J/mol x K; entropy distributed into a system at a given temp makes system’s entropy increase; entropy distributed by a system to surroundings makes system’s entropy decrease

47
Q

entropy of the universe

A

second law ultimately claims entropy of universe is increasing: delta S sub universe = delta S sub system + delta S sub surroundings > 0

48
Q

equation for standard entropy change for a reaction

A

delta S, degree sign, sub rxn = sum of delta S, degree sign, sub f, products - sum of delta S, degree sign, sub f, reactants

49
Q

Gibbs Free energy (G)

A

state function that is a combination of temp, enthalpy and entropy; delta G (change in Gibbs free energy) is a measure of the change in enthalpy and change in entropy as a system undergoes a process; indicates whether a reaction is spontaneous or nonspontaneous. change in free energy is the max amount of energy released by a process - occurring at constant temp and pressure - available to perform useful work

50
Q

equation for change in Gibbs free energy

A

delta G = delta H - T delta S; T = temp in kelvins; T delta S = total amount of energy that is absorbed by a system when its entropy increases reversibly

51
Q

delta G < 0

A

decrease in Gibbs free energy meaning movement is spontaneous; exergonic

52
Q

delta G > 0

A

increase in Gibbs free energy meaning movement away from equilibrium which is not spontaneous; endergonic

53
Q

delta G = 0

A

system is in equilibrium; delta H = T delta S

54
Q

equilibrium between gas and solid

A

delta G = G (g) - G (s) = 0, therefore G (g) = G (s)

55
Q

temp and delta G

A

delta G is temp dependent when delta H and delta S have the same sign.

56
Q

spontaneity based on delta H and delta S

A

When delta H and delta S are both positive, reaction will be spontaneous at high temps; when delta H is positive and delta S is negative, rxn will be nonspontaneous at all temps; when delta H is negative and delta S is positive, rxn will be spontaneous at all temps; when delta H and delta S are negative rxn will be spontaneous at low temps

57
Q

standard free energy

A

free energy change of reactions measured under standard state conditions; delta G, degree sign, sub rxn. for standard free energy determinations, concentrations of any solutions in the reaction are 1 M

58
Q

standard free energy of formation of a compound

A

delta G, degree sign, sub f = free energy change that occurs when 1 mole of a compound in its standard state is produced from its respective elements in their standard states under standard state conditions; this for any element under standard state conditions is 0

59
Q

equation for free energy of the reaction

A

calculated from free energies of formation of the reactants and products: delta G, degree sign, sub rxn = sum of delta G, degree sign, sub f, products - sum of delta G, degree sign, sub f, reactants

60
Q

equation for deriving free energy change for a reaction from K sub eq

A

delta G, degree sign, sub rxn = -RT x ln of Keq; where R = ideal gas constant; T = temp in kelvins; Keq = equilibrium constant; the more positive Keq the greater the natural log and the more negative (more spontaneous) standard free energy change

61
Q

free energy change of a reaction in progress

A

delta G sub rxn = delta G, degree sign, sub rxn + RT x ln of Q = RT x ln of Q/Keq; if ratio of Q/Keq is less than one (Q < Keq) then natural log will be negative and free energy change will be negative (meaning reacting will spontaneously proceed forward); if ratio is greater than one reaction will move spontaneously in reverse

62
Q

Euler’s number

A

about 2.7; natural log of some number means e to the power of whatever the natural log of that number equals

63
Q

difference between endergonic and endothermic vs exothermic and exergonic

A

endergonic and exergonic refer to change in free energy (usually Gibbs); endothermic and exothermic refer to transfer of heat or change in enthalpy

General Chemistry (11 decks)

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