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

4. Specific Heat:

ρwater =

5. Heat Capacity:

Calculating thermal energy transfer using C or ρLet q =

Conventions for q:

q ( )

q (-)

Note:

1st Law of Thermodynamics

Energy Changes Associated with Phase TransitionsA. Melting/freezing:

B. Vaporization/condensation:

Enthalpies of vaporization and fusion for some selected substances

Silberberg Fig. 12.1

Heating/cooling Curve for Water at 1 atm

-40-20

020406080

100120140

Thermal Energy Added ->

Temp

erat

ure,

C

How are Thermal Energy Changes Measured?1. Constant Pressure Calorimetry

Silberberg Fig 6.10

2. Thermal Energy Change at Constant Volume:

Silberberg Fig 6.11

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)

Some components of internal energy

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 =

Determining Work Done by (or on) a Reaction System at Constant Pressure

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

Rules for Working with Thermochemical Equations

1.

2.

3.

4.

5.

6.

7.

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

Standard Enthalpy of Formation ∆Hof

∆Hof =

C(s) + O2(g) → CO2(g) ∆Ho = ∆Hof of CO2 (g)

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

Using ∆Hof to Calculate ∆Ho

rctn

Entropy = Disorder = S

Third Law of Thermodynamics:

Second Law of Thermodynamics:

For a system at constant pressure…

Under standard state conditions,

If ∆G < 0 (-) ⇒

If ∆G > 0 (+) ⇒

If ∆G = 0 ⇒∆H ∆ S ∆G- ++ -- -+ -

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

Two Ways to Calculate ∆Go