ch 7 - Thermochemistry Flashcards Preview

General Chemistry > ch 7 - Thermochemistry > Flashcards

Flashcards in ch 7 - Thermochemistry Deck (63)
Loading flashcards...


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)



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



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



gas to liquid; facilitated by lower temp or higher pressure


melting or fusion

transition from solid to liquid


solidification, crystallization, freezing

liquid to solid



solid directly to gas phase



from gas to solid



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



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