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Chemical Thermodynamics:Energy Changes in Chemical Systems
Conversion of energy from one form to another
Transfer of energy from one place to another
Why do we care about Thermodynamics?
Practical applications:
energy production; fuels & foods
energy conversion, storage & transfer
Fundamental to understanding most areas of chemistry (equilibrium, kinetics, electrochemistry…)
Allows us to predict reaction spontaneity
Definitions and Basic Concepts1. Energy: the capacity to do work
Units: Joule
calorie
BTU
1 cal = J
a. Kinetic Energy: energy of motion
KE = ½ m v2
b. Potential Energy: “stored energy”
-energy available due to an object’s position
-chemical energy: energy stored in chemical bonds
c. thermal energy : energy associated with the random motion of particles (atoms/ions/molecules)
heat = transfer of thermal energy from
warmer to cooler object
symbol = q2.a. System: specific part of universe under study
2.b. Surroundings: everything else in the universe
Types of Systems:- open:- closed:- isolated:
3. State Functions: properties that are determined only by the current physical state of the system.
Examples of State Functions include:
Example: ∆T =Fictitious Temperature - Time Plot
-10
-5
0
5
10
15
20
5:00am
6:00am
7:00am
8:00am
9:00am
10:00am
11:00am
noon 1:00pm
2:00pm
Tem
pera
ture
, F
Calculating thermal energy transfer using C or ρLet q =
Conventions for q:
q ( )
q (-)
Note:
1st Law of Thermodynamics
Heating/cooling Curve for Water at 1 atm
-40-20
020406080
100120140
Thermal Energy Added ->
Temp
erat
ure,
C
In addition to gaining or losing thermal energy (q) in a process,
Sign conventions
q (+)q (-)w (+)w (-)
E = internal energy: many factors contribute to the value of E for a system.
Components of Internal EnergyContributions to the kinetic energy:
• The molecule moving through space, Ek(translation)
• The molecule rotating, Ek(rotation)
• The bound atoms vibrating, Ek(vibration)
• The electrons moving within each atom, Ek(electron)
Contributions to the potential energy:
• Forces between the bound atoms vibrating, Ep(vibration)
• Forces between nucleus and electrons and between electrons in each atom, Ep(atom)
• Forces between the protons and neutrons in each nucleus, Ep(nuclei)
• Forces between nuclei and shared electron pair in each bond, Ep(bond)
Note: it is not possible to determine absolute values for the internal energy of a system (E), but it is possible to determine changes in internal energy, ∆E.
C8H18(l) + 12.5O2 → 8CO2(g) + 9H2O(l)
C8H18(l) + 12.5O2
8CO2(g) + 9H2O(l)
How does a system (like a chemical reaction) do work?
∆V = Vf - Vi
Work done by a chemical reaction:
Zn(s) + 2HCl(aq) → ZnCl2(aq) + H2(g)
Note: opposing pressure P is constantSilberberg Fig 6.4Silberberg Fig 6.7
A chemical reaction can do work if
Zn(s) + 2HCl(aq) → ZnCl2(aq) + H2(g)
∆n =
3C(s) + 4H2(g) → C3H8(g)
∆n =
Enthalpy of Reaction = ∆H or ∆Hrctn
Some Important Types of Enthalpy Change
heat of combustion (∆Hcomb)
C4H10(l) + 13/2O2(g) 4CO2(g) + 5H2O(g)
heat of formation (∆Hf)
K(s) + 1/2Br2(l) KBr(s)
heat of fusion (∆Hfus)
NaCl(s) NaCl(l)
heat of vaporization (∆Hvap)
C6H6(l) C6H6(g)
Thermochemical EquationsInformation about both reaction stoichiometry
andenthalpy change for the reaction as written
H2O(l) H2O(g) ∆H = +44.0 kJ @ 25oC (= ∆Hvap)
CH4(g) + 2O2(g) CO2(g) + 2H2O(l) ∆H =-890.4 kJ @ 25oC
Standard Enthalpy of Reaction ∆Ho
When a reaction is carried out under thermo-dynamic standard conditions, the enthalpy change is ∆Ho
CH4(g) + 2O2(g) CO2(g) + 2H2O(l) ∆Ho =-890.4 kJ @ 25oC
CH4(g) + 2O2(g) CO2(g) + 2H2O(g) ∆Ho =-802.4 kJ @ 25oC
Table 6.5 Selected Standard Heats of Formation at 250C(298K)
Formula ∆H0f(kJ/mol)
calciumCa(s)CaO(s)CaCO3(s)
carbonC(graphite)C(diamond)CO(g)CO2(g)CH4(g)CH3OH(l)HCN(g)CSs(l)
chlorineCl(g)
0-635.1
-1206.9
01.9
-110.5-393.5-74.9
-238.6135
87.9
121.0
hydrogen
nitrogen
oxygen
Formula ∆H0f(kJ/mol)
H(g)H2(g)
N2(g)NH3(g)NO(g)
O2(g)O3(g)H2O(g)
H2O(l)
Cl2(g)
HCl(g)
0
0
0
-92.30
218
-45.990.3
143-241.8
-285.8
107.8
Formula ∆H0f(kJ/mol)
silverAg(s)AgCl(s)
sodium
Na(s)Na(g)NaCl(s)
sulfurS8(rhombic)S8(monoclinic)SO2(g)
SO3(g)
0
0
0
-127.0
-411.1
2-296.8
-396.0
When ∆H is (-) and ∆S is (-):
∆G vs Temp: ∆G = ∆H - T∆S
-200
0
200
-100 0 100 200 300 400 500 600 700 800 900 1000
Temp, C
ΠG,
kJ
∆H = -150 kJ; ∆S = -250 J/K Temp, C ∆H T∆S ∆G
-50 -150 -55.7875 -94.2125-25 -150 -62.0375 -87.9625
0 -150 -68.2875 -81.712525 -150 -74.5375 -75.462550 -150 -80.7875 -69.2125
100 -150 -93.2875 -56.7125200 -150 -118.2875 -31.7125300 -150 -143.2875 -6.7125400 -150 -168.2875 18.2875500 -150 -193.2875 43.2875600 -150 -218.2875 68.2875700 -150 -243.2875 93.2875800 -150 -268.2875 118.2875900 -150 -293.2875 143.2875
When ∆H is (+) and ∆S is (+):
Free Energy Change vs Temp
-200
0
200
-100 0 100 200 300 400 500 600 700 800 900Temp, C
delt
a G
, kJ
∆H= +95kJ; ∆S = +225 J/K