GE-CHEM 1103 Module 4.1 Flashcards
the ability to do
work or transfer heat
Energy
the study of energy and its
transformations
Thermodynamics
the study
of chemical reactions and the energy
changes that involve heat
Thermochemistry
most important form
of potential energy in
molecules
electrostatic
potential energy (Eel = kQ1Q2/d)
the unit of
energy commonly used
Joule ( 1J = 1 kg m^2/s^2)
it is
seen between
oppositely charged ions
Electrostatic attraction
what happens to energy when chemical bonds are formed?
Energy is released
what happens to energy when chemical bonds are broken?
Energy is
consumed
State the First Law of Thermodynamics
Energy can be converted from one form to
another, but it is neither created nor destroyed.
The portion of the
universe that we single
out to study is called
system
are
everything else that is not being studied
surroundings
a region of the
universe being studied that
can exchange heat AND
mass with its surroundings.
Open System
a region of
the universe being studied
that can ONLY exchange
heat with its surroundings
Closed System
a region of
the universe that can NOT
exchange heat or mass with
its surroundings
Isolated System:
the sum of all
kinetic and potential energies of all components
of the system
Internal energy; E
But we dont know E only how it changes so ΔE.
final energy of the system minus the initial
energy of the system
change in internal energy, ΔE
ΔE = Efinal − Einitial
the system absorbed energy from the
surroundings
ΔE > 0, Efinal > Einitial
the system released energy to the
surroundings
ΔE < 0, Efinal < Einitial
What is the value of ΔE if Efinal equals Einitial
ΔE = 0
Thermodynamic Quantities
Have Three Parts
1) A number
2) A unit
3) A sign
results when the system
gains energy from the surroundings.
positive ΔE
results when the system
loses energy to the surroundings.
negative ΔE
When energy is
exchanged between
the system and the
surroundings, it is
exchanged as either
heat (q) or work (w).
ΔE = q + w
Sign conventions for q
+q = system gains heat
-q = system loses heat
Sign conventions for w
+w = work done on system
-w = work done by system
Sign conventions for ΔE
+ΔE = net gain of energy by system
-ΔE = net loss of energy by system
When heat is absorbed by the system from the
surroundings
endothermic
When heat is released by the system into the
surroundings
exothermic
internal energy of a system
Is a state function (ΔE depends only on Einitial and Efinal)
q(heat) and w(work)
Is not a state function
the mechanical
work associated
with a change in
volume of gas
the only
work done by
chemical or
physical change
measure the work done by the gas if the
reaction is done in a vessel that has been fitted with a piston. work is NEGATIVE because it is work done
BY the system
w = −PΔV
the internal energy plus the
product of pressure and volume
Enthalpy
When the system changes at constant
pressure, the change in enthalpy, ΔH, is
ΔH = ΔE + PΔV to
Since ΔE = q + w and w = −PΔV, we
can substitute these into the enthalpy
expression:
ΔH = ΔE + PΔV
ΔH = (q + w) − w
ΔH = q
at constant pressure, the change in
enthalpy is
the heat gained or lost
ΔH is
positive
process is
endothermic
ΔH is negative
process is
exothermic
is the
enthalpy of the
products minus the
enthalpy of the
reactants
The change in
enthalpy, ΔH
ΔHrxn
enthalpy of
reaction, or the heat of reaction
The Truth about Enthalpy of reaction
- Enthalpy of reaction is an extensive property.
- The enthalpy change for a reaction is
equal in magnitude, but opposite in
sign, to ΔH for the reverse reaction. - The enthalpy change for a reaction
depends on the states of the reactants
and the products.
is
defined as the enthalpy change for the
reaction in which a compound is made
from its constituent elements in their
elemental forms
enthalpy of formation, ΔHf
are
measured under standard conditions (25 °C
and 1.00 atm pressure).
Standard enthalpies of formation, ΔHf
°
enthalpy of formation for an element
in its elemental state
0 because it takes no energy to form a
naturally-occurring compound.
Calculation of ΔH (see the answer right way)
Do some exercises on it
Calculation of ΔH using Hess law
ΔH = ΣnΔHf,products – ΣmΔHf
°,reactants
The enthalpy associated with breaking one
mole of a particular bond in a gaseous
substance
Bond Enthalpy
always positive because energy is
required to break chemical bonds
bond enthalpy
always released when a bond forms between
gaseous fragments
Energy
greater the bond enthalpy
stronger the bond.
ALL bonds made
ADD bond energy
ALL bonds
broken
SUBTRACT bond energy
Σ (bond enthalpies – Σ (bond enthalpies
of bonds broken) of bonds formed)
ΔHrxn
If a reaction is carried out in a series of steps,
ΔH for the overall reaction equals the sum of the enthalpy
changes for the individual steps.
Hess’s law
measure ΔH through measurement of heat flow
calorimetry
instrument used to
measure heat flow
calorimeter
The amount of energy required to raise the temperature of a
substance by 1 K (1 °C)
heat capacity
the amount of
the substance heated is one gram
specific heat
If
the amount is one mole
molar heat capacity
The specific heat for water
4.184 J/g∙K (can be used for dilute solutions)
calculate ΔH for the reaction
qsoln = Cs × msoln × ΔT = –qrxn
- Reactions can be carried
out in a sealed “bomb” - Because the volume in the
bomb calorimeter is
constant, what is measured
is really the change in
internal energy, ΔE, not ΔH.
Bomb Calorimetry
The heat absorbed (or
released) by the water is
a very good approximation
of the enthalpy change for
the reaction.
qrxn = – Ccal × ΔT
The energy released when one gram of food is
combusted
fuel value
Most of the energy in foods comes from carbohydrates,
fats, and proteins
Carbohydrates (17 kJ/g)
Fats (38 kJ/g)
Proteins (17 kJ/g)
The vast majority of the
energy consumed in
this country
fossil fuels
8.6% of the
U.S. energy needs
Nuclear fission
produce 9.9% of the
U.S. energy needs
Renewable energy
sources
