MCAT General Chemistry - Thermochemistry

system

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

surroundings/environment

everything outside of the system

Isolated system

cannot exchange energy (heat and work) or matter with the surroundings

ex. insulated bomb calorimeter

Closed system

can exchange energy (heat and work) but not matter with the surroundings

ex. steam radiator

Open system

can exchange both energy (heat and work) and matter with the surroundings

ex. pot of boiling water

process

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

conservation of energy

ΔU = Q – W

where ΔU is the change in internal energy of the system, Q is the heat added to the system, and W is the work done by the system

isothermal processes

system’s temperature is constant; ΔU = 0, Q=W, hyperbolic curve on a pressure–volume graph

pressure–volume graph

displays changes in volume compared with changes in pressure; Work is represented by the area under such a curve

Adiabatic processes

no heat is exchanged between the system and the environment; Q = 0; ΔU = –W; appears hyperbolic on a P–V graph

Isobaric processes

pressure of the system is constant; flat line on a P–V graph

isovolumetric (isochoric) processes

no change in volume; W=0; ΔU = Q; vertical line on a P–V graph

spontaneous process

one that can occur by itself without having to be driven by energy from an outside source; negative ΔG; will not necessarily happen quickly and may not go to completion

enzymes/biological catalysts

selectively enhance the rate of certain spontaneous (but slow) chemical reactions so that the biologically necessary products can be formed at a rate sufficient for sustaining life

coupling

A common method for supplying energy for nonspontaneous reactions is by pairing nonspontaneous reactions to spontaneous ones that create the necessary energy

state functions

describe the system in an equilibrium state, but not how it got there

ex. pressure (P), density (ρ), temperature (T), volume (V), enthalpy (H), internal energy (U), Gibbs free energy (G), and entropy (S)

process functions

describes pathway taken from one equilibrium state to another

ex. work (W) and heat (Q)

standard conditions

defined for measuring the enthalpy, entropy, and Gibbs free energy changes of a reaction

25°C (298 K), 1 atm pressure, and 1 M concentration

kinetics, equillibrium, thermodynamics

standard temperature and pressure (STP)

temperature is 0°C (273 K) and pressure is 1 atm

ideal gas

standard state

the most stable form of a substance; “zero point” for all thermodynamic calculations

standard enthalpy - ΔH°

standard entropy - ΔS°

standard free energy changes - ΔG°

Phase diagrams

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

Phase changes

change between states of matter; reversible, and an equilibrium of phases will eventually be reached at any given combination of temperature and pressure

evaporation/vaporization

liquid → gas; endothermic process for which the heat source is the liquid water

Boiling

specific type of vaporization; rapid bubbling of the entire liquid with rapid release of the liquid as gas particles

condensation

gas → liquid; facilitated by lower temperature or higher pressure (vapor pressure)

boiling point

liquid-gas equilibrium temperature; the temperature at which the vapor pressure of the liquid equals the ambient (external, applied, or incident) pressure

microstates

freedom of movement; allows energy dispersion; involved in entropy

fusion/melting

solid → liquid

solidification/crystallization/freezing

liquid → solid

melting/freezing point

solid-liquid equilibrium temperature

sublimation

solid → gas

deposition

gas → solid

cold finger

device used to purify a product that is heated under reduced pressure, causing it to sublime, then deposits onto the instrument

lines of equilibrium/phase boundaries

indicate the temperature and pressure values for the equilibria between phases; interfaces

triple point

point at which the three phase boundaries meet; temperature and pressure at which the three phases exist in equilibrium

gas phase

found at high temperatures and low pressures

solid phase

low temperatures and high pressures

liquid phase

moderate temperatures and moderate pressures

critical point

where phase boundary between the liquid and gas phases terminates; temperature and pressure above which there is no distinction between the phases; densities of ‘liquid’ and ‘vapour‘ become equal; heat of vaporization at this point and above is zero

supercritical fluids

fluids existing above the critical point

Temperature (T)

related to the average kinetic energy of the particles of a substance; how hot or cold something is

scales: Fahrenheit, Celsius, and Kelvin

thermal energy (enthalpy)

realated average kinetic energy of the particles in a substance and how much substance is present

absolute temperature scale (Kelvin)

determined via the third law of thermodynamics, which elucidated that there is a finite limit to temperature below which nothing can exist

Heat (Q)

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

q = mcΔT

unit of energy: joule (J) or calorie (cal) (1 cal = 4.184 J)

zeroth law of thermodynamics

objects are in thermal equilibrium only when their temperatures are equal

endothermic

Processes in which the system absorbs heat; ΔQ > 0

exothermic

processes in which the system releases heat; ΔQ < 0

Enthalpy (ΔH)

equivalent to heat under constant pressure
ΔH_rxn = H_products – H_reactants

calorimetry

process of measuring transferred heat; constant pressure and sonstant volume

Specific heat

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

heat capacities

mass times specific heat

constant-pressure calorimeter

insulated container covered with a lid and filled with a solution in which a reaction or some physical process, such as dissolution, is occurring

bomb calorimeter/decomposition vessel/constant pressure calorimeter

a sample of matter, typically a hydrocarbon, is placed in the steel decomposition vessel, which is filled with almost pure oxygen gas, then ignited by an electric ignition mechanism; heat that evolves is the heat of the combustion reaction; no work

Heating curves

show that phase change reactions do not undergo changes in temperature

enthalpy/heat of fusion/vaporization (ΔH_fus/vap)

used to determine the heat transferred during the phase change

q = mL

where m is the mass and L is the latent heat, a general term for the enthalpy of an isothermal process, given in the units cal/g

standard enthalpy of formation (ΔH°_f)

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

standard enthalpy of a reaction (ΔH°_rxn)

enthalpy change accompanying a reaction being carried out under standard conditions

ΔH°_rxn = Σ ΔH°_f,products − Σ ΔH°_f,reactants

Hess’s law

enthalpy changes of reactions are additive

ΔH_{reactants → elements} = –ΔH_{elements → reactants}

bond enthalpies/dissociation energies

average energy that is required to break a particular type of bond between atoms

kJ/mol

standard heat of combustion (ΔH°_comb)

enthalpy change associated with the combustion of a fuel

Entropy

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

where ΔS is the change in entropy, Qrev is the heat that is gained or lost in a reversible process, and T is the temperature in kelvin.

units: J/mol*K

second law of thermodynamics

energy spontaneously disperses from being localized to becoming spread out if it is not hindered from doing so; time’s arrow: unidirectional limitation on the movement of energy by which we recognize before and after or new and old

ΔS_universe = ΔS_system + ΔS_surroundings > 0

standard entropy change for a reaction (ΔS°_rxn)

ΔS°_rxn = Σ ΔS°_f,products − Σ ΔS°_f,reactants

Gibbs free energy, G

measure of the change in the enthalpy and the change in entropy as a system undergoes a process; indicates whether a reaction is spontaneous or nonspontaneous; maximum amount of energy released by a process—occurring at constant temperature and pressure—that is available to perform useful work

ΔG = ΔH – TΔS

where T is the temperature in kelvin and TΔS represents the total amount of energy that is absorbed by a system when its entropy increases reversibly

exergonic

system releases energy; spontaneous

endergonic

system absorbs energy; nonspontaneous

standard free energy (ΔG°_rxn)

free energy change of reactions can be measured under standard state conditions

ΔG°_rxn = Σ ΔG°_f,products − Σ ΔG°_f,reactants

ΔG°_rxn = –RT ln K_eq

where R is the ideal gas constant, T is the temperature in kelvin, and K_eq is the equilibrium constant

ΔG_rxn = ΔG°_rxn + RT ln Q = RT ln Q/K_eq