Chapter 10 Energy and Energy changes Flashcards
thermochemistry
the study of energy and energy transfer in chemical processes
energy
the capacity to do work
work
a force acting through a distance
pontential energy
stored energy, gravitational potential energy, chemical bonds.
kinetic energy
energy of motion
system
the portion of universe of interest
surrounding
the rest of the universe not part of the system
temperature
a measure of the average kinetic energy
thermal energy
the internal energy present in a system due to its molecular motion
heat
the transfer of thermal energy due to differences in temperature
heat moves from a region of higher temperature to a region of lower temperature
the process of heat flow is defined with repeat to the system
endothermic
heat is absorbed by the system
exothermic
heat is released by the system
Si Units
Joule, J
Calorie
the energy required to raise exactly 1 g H2O by 1ºC
1cal=1000cal
1cal=1kcal
1cal=4.184 j
types of system
characterized by transfer between system and surrounding
state function
a value or variable that only depends on the state of the system not how it achieved that state
internal energy, E
the sum of all the kinetic and potential energy within a system
E is a state function
E is based on parameters such as temperature, pressure, concentration, phase, etc.
Change internal energy, ∆E
its is difficult to measure E so ∆E is measured
first law of thermodynamics
in chemical processes, energy is neither created nor destroyed. the total energy of the universe is constant in all processes, the change in energy of the universe is zero.
1 last applied
if energy cannot be crated nor destroyed…t=the energy lost by a system must be the same amount gained by the surrounding. The energy gained by a system must be the same amount lost by the surroundings
∆Esystem=-∆Esurrounding
∆Esystem+∆Esurrounding=0
changes in internal energy
the change in internal energy occurs via two forms:
1.heat(q) 2.work(w)
∆E=q+w
heat
+q=system gains/absorbs thermal energy
-q= system loses/releases thermal energy
work
+w=work is done on the system
-w=work is done by the system
pressure-volume
P-V work
work as a result of a volume change agains an external pressure
w=-P∆V
P:exnternal Pressure, atm
∆V: change in volume
∆V: Vfinal-Vintial
for expansions Vf>Vi work is done by the system w<0
for contraction Vf0
constants volume work
if the reaction is carried in a sealed container the volume is fixed
∆V=0 w=-P∆V=0J
therefore, the total change in internal energy at constant volume is:
∆E=q+w. ∆E=q+0 ∆E=qw
enthalpy
often reactions are carried out under constant pressure; not constant volume
∆E=q+w ∆E=qp+(-P∆V)
rearranging, qp=∆E+P∆V, this is defined as enthalpy H: qp=∆H
enthalpy H
for a system, enthalpy is the sum of the internal energy plus the pressure times volume
∆H>0 endothermic: heat flows into the system
∆H<0 exothermic:heat flows out of the system
Stoichiometry of ∆H
when ∆H is given for a reaction, it is related to the moles shown in the balanced reaction.
Calorimetry
the study of heat change in chemical reactions. when a system absorbs heat, there is a change in temperature
experimental it is observed: the change in temperature depends on the amount of heat absorbed
qα∆T
heat capacity, c
the constant of proportionality that relates the temperature change to the heat absorbed
heat capacity (of a system)
the quantity of heat required to change temperature of a system by 1ºC
q=C•∆T
the heat capacity depend on the amount of matter present in the system and is report for the system
specific heat capacity, Cs
the amount of energy required to raise the temperature of 1 gram of substance by 1ºC
q=m•Cs•∆T
molar heat capacity, Cm
the amount go energy required to raise the temperature of 1 mole of substance by 1ºC
Cm=J/mol ºC
constant pressure calorimetry
often reaction are conducted at constant pressure, not constant volume
∆Hrxn=qrxn/nreactant
thermochemical equation
chemical equation that shoes the enthalpy change in addition to the balanced species
Manipulating ∆H
- if a reaction is multiplied by n, the ∆His also multiplied by n
- if a reaction is reversed, the sign on ∆H is changed
- if a chemical equation can be expressed as the sum of a series of steps, then the ∆Hrxn for the overall reaction is the sum of the individual ∆Hs
Standard states for enthalpy
a. for a gas the pure gat as a pressure of exactly 1 atm
b. for liquid or solid the pure substance in its most stables from at a pressure of 1 atm and temperature of interest
c. for substance in solution the substance in solution at a concentration of exactly 1 M
Standard enthalpy change, ∆H
the change in enthalpy for a process when all reactants and products are in their standard states
standard enthalpy of formation,∆Hfº
for a pure compound the change in enthalpy when 1 mole of the compound is formed from its constituent elements in their standard states
standard enthalpy of reaction, ∆Hºrxn
the ∆Hfº provides the energy of each substance, therefore for a reaction the ∆Hrxn is given by:
∆Hºrxn= ∑nproduct∆Hfºproduct-∑nreactant∆Hfºreactant
n=number of moles
bond enthalpies, ∆HBE
the energy required to break 1 mole of the bond for the gas phase species. Also known as bond dissociation energy
