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