ch 7 - Thermochemistry Flashcards
system
the matter that is being observed - the total amount of reactants and products in a chemical reaction
surroundings or environment
everything outside of the system
characterizations of systems
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)
process
when a system experiences a change in one or more of its properties (such as concentrations of reactants or products, temperature, or pressure
first law of thermodynamics
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
isothermal processes
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
Adiabatic processes
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
isobaric processes
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
isovolumetric (isochoric) processes
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
spontaneous process
one that can occur by itself without having to be driven by energy from an outside source
state functions
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)
process functions
pathway taken from one equilibrium state to another, quantitatively. The most important of these are work and heat
standard conditions
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
standard temp and pressure (STP)
used for ideal gas calculations: temp is 0 degrees C (273 K) and pressure is 1 atm.
standard state
most stable state of a substance under standard conditions
standard enthalpy (delta H degree sign), standard entropy (delta S degree sign), standard free energy changes (delta G degree sign)
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
Phase diagrams
graphs that show the standard and nonstandard states of matter for a given substance in an isolated system, as determined by temps and pressures
evaporation
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
condensation
gas to liquid; facilitated by lower temp or higher pressure
melting or fusion
transition from solid to liquid
solidification, crystallization, freezing
liquid to solid
sublimation
solid directly to gas phase
deposition
from gas to solid
temperature
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
Heat (Q)
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
zeroth law of thermodynamics
implies that objects are in thermal equilibrium only when their temps are equal
calorimetry
process of measuring transferred heat; two basic types are constant-pressure calorimetry and constant-volume calorimetry
equation for heat absorbed or released in a given process
q = mc deltaT (q = mcAt); m = mass; c = specific heat of the substance; delta T = change in temp (C or K)
specific heat (c)
the amount of energy required to raise the temp of one gram of a substance by one degree Celsius (or one Kelvin)
specific heat of H2O
c sub H2O = 1 cal/g x K
heat capacity
product of mc (mass times specific heat)
constant-pressure calorimeter
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
constant-volume calorimeter
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
enthalpy (or heat) of fusion (delta H sub fus)
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
enthalpy (or heat) of vaporization (delta H sub vap)
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
enthalpy (H)
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
enthalpy change of a reaction
delta H sub rxn = H sub products - H sub reactants; positive delta H sub rxn corresponds to an endothermic process; negative corresponds to exothermic
standard enthalpy of formation (delta H, degree sign, sub f)
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
standard heat of reaction (delta H degree sign, sub rxn)
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
Hess’s Law
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
bond enthalpies
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
equation for enthalpy change associated with a reaction
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
standard heat of combustion (delta H, degree sign, sub comb)
the enthalpy change associated with the combustion of a fuel
second law of thermodynamics
states that energy spontaneously disperses from being localized to becoming spread out if it is not hindered from doing so.
entropy
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
equation for change in entropy
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
entropy of the universe
second law ultimately claims entropy of universe is increasing: delta S sub universe = delta S sub system + delta S sub surroundings > 0
equation for standard entropy change for a reaction
delta S, degree sign, sub rxn = sum of delta S, degree sign, sub f, products - sum of delta S, degree sign, sub f, reactants
Gibbs Free energy (G)
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
equation for change in Gibbs free energy
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
delta G < 0
decrease in Gibbs free energy meaning movement is spontaneous; exergonic
delta G > 0
increase in Gibbs free energy meaning movement away from equilibrium which is not spontaneous; endergonic
delta G = 0
system is in equilibrium; delta H = T delta S
equilibrium between gas and solid
delta G = G (g) - G (s) = 0, therefore G (g) = G (s)
temp and delta G
delta G is temp dependent when delta H and delta S have the same sign.
spontaneity based on delta H and delta S
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
standard free energy
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
standard free energy of formation of a compound
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
equation for free energy of the reaction
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
equation for deriving free energy change for a reaction from K sub eq
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
free energy change of a reaction in progress
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
Euler’s number
about 2.7; natural log of some number means e to the power of whatever the natural log of that number equals
difference between endergonic and endothermic vs exothermic and exergonic
endergonic and exergonic refer to change in free energy (usually Gibbs); endothermic and exothermic refer to transfer of heat or change in enthalpy
