Plan for Wed, 15 Oct 08. Lecture The nature of energy (6.1) Enthalpy and calorimetry (6.2). Thermodynamics. The study of energy and its interconversion . Energy is the capacity to do work or to produce heat
Potential Energy (PE) – the energy an object has by virtue of its placement in a field of force, like gravity.
PE = mgh, where m = mass, g = acceleration due to gravity, h = height
Ball A has higher PE than ball B.
Kinetic Energy (KE) – the energy an object has by virtue of its motion
KE = ½ mv2, where m = mass, v = velocity
ball A’s PE is converted to KE when it rolls down the hill...
ball A transfers its KE to ball B by doing work
it also loses some of its PE to frictional heating of the hill.Energy: Kinetic vs. Potential
xExample: Mass on a Spring
Joule = kg.m2/s2
x = 0
Molecules are subject to more forces than just gravity.
(and Chemical Bonds)
Petrucci, Fig. 7.9
Translation: motion through space
Rotation: motion about the center of mass
Vibration: motion directed through chemical bonds
Petrucci, Fig. 7.9
Molecular Kinetic Energy
Molecular Potential Energy
(and chemical bonds)
Chemical reactions and phase changes involve changes in both the kinetic and potential energy of molecules.
Petrucci, Fig. 7.9
First Law of Thermodynamics: Total energy of the universe is constant.
To help keep track of energy flow let’s define...
System: that one part of the universe you are interested in (i.e., you define it).
Surroundings: everything else in the universe.
Energy gained by the system must be lost by the surroundings.
Energy exchange can be in the form of heat (q), work (w), or both.Conservation of Energy
What do these have to do with molecular KE and PE?
By measuring the heat absorbed or evolved, and the work performed on or by the system, we can obtain information about how the internal energy of the system changes during a process.
q = 0
w < 0
Work: energy transferred as a force applied over a distance.
= Pext x A x Dh
P = Psurr
DV of sys!!!
systemHeat and Work
w = F x d
q = heat > 0
w = work < 0
System gains energy as heat(q), causing the gas to expand.
The expanding gas exerts force on the piston, causing it to move, doing work (w)on the surroundings. System loses energy as work done on surr.
System loses energy as heat.
Production of H2(g) causes piston to move... system loses energy as work done on surroundings.
q < 0 (q is negative)
If system gets heat
q > 0 (q is positive)Conservation of Energy Revisited
DE = q + w
= (1.3 x 108 J) + (-PDV)
= (1.3 x 108 J) + (-1 atm (Vfinal - Vinit))
= (1.3 x 108 J) + (-0.5 x 106 L.atm)
Conversion: 101.3 J per L.atm
(-0.5 x 106 L.atm) x (101.3 J/L.atm) = -5.1 x 107 J
DE = (1.3 x 108 J) + (-5.1 x 107 J)
= 8 x 107 J
In English: the system gained more energy through heat than it lost doing work. Therefore, the overall energy of the system has increased.
Efinal < Einitial
Efinal-Einitial = DE < 0.
Efinal > Einitial
Efinal-Einitial = DE > 0.
Suppose that you have two identical 500-mL bottles of water sitting on your desk at equilibrium.
Bottle 1 has been kept at 25°C since bottling.
Bottle 2 came from the same spring, but has been frozen, thawed, transported by air in an unpressurized compartment, and allowed to fluctuate wildly in temperature before being at equilibrium on your desk.
Which bottle has the higher internal energy?
sitting on your desk at equilibrium.U =q+wU as a state fxn
Combustion of gas in a bomb calorimeter:
DV = 0 w = -PDV = 0
All the energy produced in the combustion rxn is evolved as heat.
Combustion of gas in your car:
DV > 0 w = -PDV < 0
Some of the energy produced is evolved as heat, and some is evolved as work done on the surroundings.
BUT IN BOTH CASES DU IS THE SAME.
we could determine how the internal energy of a system changes during a process.
Enthalpy is defined as:
Consider a process carried out at constant P (w = -PDV):
Note similarity to our definition of enthalpy:
What is this telling us?
We can track the heat flow in a process occurring at constant P and that will give us direct info about DE.
Heat Capacity (C): the energy required to raise the temp. of a sample of a substance by 1oC; the ability of a substance to absorb heat.
Specific heat (s): the energy required to raise the temperature of 1 gram of a substance by 1oC.
If we know the specific heat of a substance, the mass, and the temperature change, we can determine the heat flow.
NaOH + HCl H2O + NaCl; DH = -58 kJ
The heat evolved in this reaction is trapped in the water…this heat increases the average Ek of the water molecules, leading to an increase in the temperature of the water.
qsystem = -qsurroundings
Formally, q is the amount of heat that must be exchanged with the surroundings to return the system to its original temperature.
In calorimetry we don’t let it escape…q goes into raising the temperature of the water.