Transcript Thermodynamics

Thermodynamics
from Greek thermo dy’namis (heat and power)
Studies energy changes and the direction of flow of energy usually in a well-defined
part of the universe (the system)
Definitions:
System: part of the universe in which we are interested
Surroundings: where we make our observations (the universe)
Boundary: separates above two
Heat and Work
Heat: transfer of energy that
changes motions of atoms
in the surroundings in a
chaotic manner
Work: transfer of energy
that changes motions of
atoms in the surroundings
in a uniform manner
=Fxd
Energy
• Definition: the capacity to do WORK
• Units are Joules (J) = kg.m2/s2
(from KE=1/2mv2)
Work done on a system - system gains energy (w +ve)
Work done by the system - system loses energy (w -ve)
Heat absorbed by the system (endothermic) - system gains energy (q +ve)
Heat released by the system (exothermic) - system loses energy (q +ve)
SYSTEM TOTAL ENERGY (kinetic plus potential) is the
INTERNAL ENERGY (U sometimes E)
Usually measure CHANGE in internal energy ( U )
U=Ufinal – Uinitial
U is a STATE FUNCTION (independent of path)
4
st
1
LAW of Thermodynamics
Internal energy of an isolated system is constant
(energy can neither be created nor destroyed)
U = q+w
Pressure-Volume work
Against constant external pressure
w = -F.dz but Pex=F/A therefore w= -Pex.dV
Free expansion
w=0
Calorimetry
Can measure internal energy
changes in a “bomb” calorimeter
U=q-P V, but in a constant
volume “bomb”, V=0
Thus U=q
Heat Capacity
Amount of energy required to raise the temperature of a substance by
1C (extensive property)
For 1 mol of substance: molar heat capacity (intensive property)
For 1g of substance: specific heat capacity (intensive property)
 U 
CV  

 T V
U  CV T  qV 
If heat capacity is independent of
Temperature over the range of interest
Most reactions we investigate occur under
conditions of constant PRESSURE (not Volume)
Enthalpy
Heat of reaction at constant pressure!
H  U  PV
H  U  P V
but w  - PV
H  q P
Heat capacity
Use a “coffee-cup” calorimeter
to measure it
 H 
CP  

 T  P
H  CP T  qP 
Excercise: When 50mL of 1M HCl is mixed with 50mL of 1M
NaOH in a coffee-cup calorimeter, the temperature increases
from 21oC to 27.5oC. What is the enthalpy change, if the
density is 1g/mL and specific heat 4.18 J/g.K?
Problem: Heat Capacities & Temperature Changes
How much heat is required to raise the temperature of 10 g
of water and 10g of lead from 0 to 50oC?
specific heat of H2O = 4.18 J/g-oC
specific heat of Pb = 0.128 J/g-oC
q = m×c×∆T
q(H2O) = 10g×4.18 J/g-oC×50oC
= 2090 J
q(Pb) = 10g×0.128 J/g-oC×50oC
= 64 J
Problem: Heats of Chemical Reaction
100 ml solutions of 1.00 M NaCl and 1.00 M AgNO3 at 22.4 oC are mixed in coffee
cup calorimeter and the resulting temperature rises to 30.2 oC.
What is the heat per mole of product? Assume the solution density and specific heat
are the same as pure water.
Write balanced chemical reaction:
Net ionic: Ag+(aq) + Cl-(aq) → AgCl(s)
Determine heat of reaction:
qrxn= -qcal = -m×c×∆T
m = 200 ml × 1.0g/ml = 200g
c = cH2O = 4.18 J/g-oC
= -200g × 4.18 J/g-oC × (30.2-22.4)
= -6,520 J
Determine heat per mole of product:
stoichiometric reactants, 0.1 mol in 100 ml
qrxn/mol = -6.52 kJ/0.1 mol
See student activities
= -65.2 kJ/mol
U and H (see chapter 10)
H  U  PV
Only differ significantly
when gases are involved
 nRT 
 RT 
V  
  ng 

 P 
 P 
 H  U  ng RT
Standard Enthalpy Changes, Ho
H for a process in which the initial and final species are in their
standard states.
Can be reported for any T. Use 298K unless otherwise indicated
Hvapo:1 mole pure liquid vapourises to a gas at 1bar
(+40.66 kJmol-1 at 373K for water) endothermic
Hfuso:1mole pure solid melts to a pure liquid at 1bar
(+6.01 kJmol-1 at 273K for ice) endothermic
Standard Reaction Enthalpy Changes
CaO(s) + CO2(g)  CaCO3(s)
rxnHo = -178.3kJmol-1
Thermochemical equations:
standard heats of reaction,
rxnHo
Represent by an
Enthalpy Diagram
Hess’s Law
If a reaction is the sum of two
separate reactions then the enthalpy
change during that reaction is also
the sum of the enthalpy changes in
the component reactions.
Hess’s Law
rxnHo = Hoproducts - Horeactants
Standard Heats of Formation
If one mole of the compound is formed under standard conditions
from its elements in their standard state then the resulting enthalpy
change is said to be the standard molar enthalpy (Heat) of
formation, fHo where the subscript indicates this.
By definition the enthalpies of formation of
the elements in their standard states are zero.
H2 (g) + 1/2O2 (g)  H2O (l) fHo = -285.8kJmol-1
2C (s) + 3H2(g) + 1/2O2(g)  C2H5OH (l) fHo = -277.7kJmol-1
Hess’s Law
Hess‘s Law is particularly useful for calculating fHo which would not
be easy to measure experimentally. fHo for CO cannot be measured
as CO2 is also formed when graphite is burned
C(s) + 1/2O2  CO fHo = x
CO + 1/2O2  CO2 rxnHo = -283 kJmol-1
_______________________________________
C(s) + O2  CO2 fHo = -393.5 kJmol-1
From looking at these equations it is fairly obvious that the sum of the
first two enthalpies is equal to the third by Hess‘s Law.
i.e. x - 283 = -393.5 or x = -110.5 kJmol-1.
1/2H2(g)
Enthalpy Changes and Bond Energies
Energy is absorbed when bonds break. The energy required to
break the bonds is absorbed from the surroundings.
If there was some way to figure out how much energy a single bond absorbed when broken,
the enthalpy of reaction could be estimated by subtracting the bond energies for bonds
formed from the total bond energies for bonds broken.
O2(g) 2O(g) H°=490.4 kJ
H2(g) 2H(g) H° =431.2 kJ
H2O(g)2H(g) + O(g) H°=915.6 kJ
We can estimate the bond enthalpies of O=O, H-H, and O-H as 490.4 kJ/mol, 431.2 kJ/mol,
and 457.7 kJ/mol, respectively.
2H2(g) + O2(g)  2H2O(g)
H°= ?
2H2(g) + O2(g)
moles of
bonds broken
Energy absorbed
 2H2O(g)
moles of bonds
formed
Energy released
2 H-H @ 431.2 kJ each
862.4kJ
1 O=O @ 490.4 kJ each
490.4kJ
_____________________________________________
1352.7kJ
4 O-H @ 457.7 kJ each
1830.9kJ
1830.9kJ
H°= 1352.7 - 1830.9 kJ = -478.2 kJ.
(Remember that the minus sign means "energy released", so you add the bond energies for
broken bonds and subtract energies for bonds formed to get the total energy.)
A calculation based on enthalpies of formation gave H° = -483.7 kJ
Bonds in a molecule influence each other, which means that bond energies aren't really additive.
An O-H bond in a water molecule has a slightly different energy than an O-H bond in H2O2,
because it's in a slightly different environment.
Reaction enthalpies calculated from bond energies are very rough approximations!
Foods and Fuels
Enthalpies (heats) of combustion: complete reaction of compounds
with oxygen. Measure using a bomb calorimeter.
Most chemical reactions used for the production of heat are
combustion reactions. The energy released when 1g of material is
combusted is its Fuel Value. Since all heats of combustion are exothermic, fuel
values are reported as positive.
Most of the energy our body needs comes from fats and carbohydrates.
Carbohydrates are broken down in the intestines to glucose. Glucose
is transported in the blood to cells where it is oxidized to produce CO2,
H2O and energy:
C6H12O6(s) + 6O2(g)  6CO2(g) + 6H2O(l) H°rxn=-2816 kJ
The breakdown of fats also produces CO2 and H2O
Any excess energy in the body is stored as fats
About 100 kJ per kilogram of body weight per
day is required to keep the body functioning at a
minimum level
Compound
Fuel Value
(kJ/gram)
Fats
38
Carbohydr
ates
17
Proteins
17
Fuels
Energy comes primarily from the combustion of fossil fuels
Coal represents 90% of the fossil fuels on earth. However, it typically contains
sulfur, which when combusted can lead to environmental pollution (acid rain)
Solar energy: on a clear day the sun's energy which strikes the earth equals 1kJ
per square meter per second.
Hydrogen: clean burning (produces only water) and high fuel value. Hydrogen
can be made from coal as well as methane
C(coal) + H2O(g)  CO(g)+H2(g)
CH4(g) + H2O(g)  CO(g) + 3H2(g)
C
(%)
Fuel
The greater the percentage of
carbon and hydrogen in the
fuel the higher the fuel value
H
(%)
O
(%)
Fuel
Value
(kJ/g)
Wood
50
6
44
18
Coal
77
5
7
32
Petrol
85
15
0
48
Hydrogen
0
100
0
142
Fuel cells
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•
•
Biofuel cell research in NUIG
Biomednano website
Combustion chemistry
Global Energy Consumption