1. Chemical Thermodynamics

2. Plan 1. Main concepts 2. Classification of chemical reactions
3. The first law of thermodynamics
4. Thermochemistry
5. The second law of thermodynamics .

3. Main Concepts

4. Thermodynamics is a branch of science that deal with the conversion of energy from one form to another.
The energy transformations determine most features of chemical reactions

5. Thermodynamics answers the questions:
Whether or not a reaction spontaneous under a given set of conditions.
To what extent will the reaction proceed? Will it be completed or partial?
How much energy will be produced or absorbed in the reaction?
What conditions are to be fulfilled to enable the reaction to proceed, or inversely, to stop

6. Two very important conclusion:
Energy is transferred in the form of either work or heat
Heat arises at each step of energy transformation whereas work does not always occur
Consequently, energy of any kind can be completely transferred into heat, whereas the transformation of energy into work is never complete

7. There are kinds of energy released in a chemical system
The part that could be converted into Work is referred to as Free Energy
The other part which may only be converted into Heat is referred to as Bound Energy (or Reversible Heat)

8. Each state of the system is characterized by its thermodynamic probability
Without external influence any change in the system results in a more probable state whose thermodynamic probability is greater than that of the state before the change

9. Formal Definitions The example: the flask with reactive mixture
A System is a substance or group of interacting substances that we consider apart from its surroundings
The example: the flask with reactive mixture

10. Kinds of Systems An Open System can exchange either energy or matter
A Closed System can exchange only energy
An Isolated System can exchange neither energy nor matter with its surroundings and has a constant volume
A Phase is a part of a system which is homogeneous throughout and separated from other parts of the system by a boundary surface

11. Fe(solid)+2H2O(g) = H2(g)+Fe2O3(solid)
A Homogeneous System (only one phase) is uniform throughout
H2(g) + Cl2(g) = 2HCl(g)
A Heterogeneous System contains the substances in different aggregative states
Fe(solid)+2H2O(g) = H2(g)+Fe2O3(solid)

12. The State of the Chemical System is described by a set of the parameters (T, P, V, n, C):
Temperature
Pressure
Volume
Amount of substances
Concentrations

13. The state functions of a chemical system:
internal energy (U)
enthalpy (Н)
entropy (S)
free energy (G)
These are thermodynamic values which characterize energetic changes of a chemical system

14. Classification of the Chemical Process
1) A sign of the process
Endothermic reaction – the system absorbs heat (+)
Exothermic reaction – system evolves heat (–)

16. 2) The conditions of a reaction
Parameters Processes
Т - const Isothermal
Р - const Isobaric
V - const Isochoric

17. 3) The Principle of Spontaneity
Free energy(G) is the criterion of the process direction:
G < 0 – a spontaneous process
G > 0 – a no spontaneous process
G = 0 – an equilibrium state

18. The First Law of Thermodynamics
The energy of the universe is constant
A change in internal energy may be caused by a change either in heat or in work, or in both
U = Q - W
U = U2 - U1 – a change of the internal energy of a system
Q - heat
W - work

19. Heat and Work
The heat effect (Q) of a reaction can be measured under constant volume (QV) or constant pressure (Qp) and it is usually measured in isochoric conditions.
In chemical reactions a work can be obtained as a result of the change of volume:
W = pV, где V= V2 - V1

20. Total Energy of a System
Kinetic energy of movement of a system as a whole
Potential energy caused by a situation of a system in an external field
Internal energy

21. For chemical reactions the change of total energy in chemical systems are determined only by the change of its internal energy
The internal energy includes forward, rotary, oscillatory energy of atoms, molecules, and also the energy of movement of electrons in atoms, internuclear energy

22. Quantity of internal energy (U) of a substance is determined by the amount of a substance, its composition and state
The stability of system is defined by the quantity of internal energy: the more internal energy is, the less steady system is

23. The Change of Internal Energy
In isochoric process (V = 0):
U = Qv
the change of internal energy occurs as the change of a heat effect because:
A = pV = 0
In isobaric process (P-const)
U = Qp - pV
the change of internal energy is the heat minus the work (pV) of expansion or compression

24. Qp=U+pV= (U2-U1) + p(V2-V1)
Qp = (U2 + pV2) - (U1 + pV1)
U + pV = Н
H - enthalpy:
Qp = H2 - H1 = H
H – matches the heat of a chemical reaction including the work of the system kJ
Measuring unit
mole

25. The absolute value of energy formation (U, H) of a substance can not be measured

26. Enthalpy of Formation ( ) of a Simple Substance
is always zero under standard conditions
(N2,gas) = 0; (Сgraphite) = 0
Standard state:
Р = 101,3kPa
n = 1 mole
concentrations - 1 mol/L
Т- any temperature when a substance can exist

27. For example: K(solid)+1/2Cl2+3/2O2=KClO3(solid)
Enthalpy of Formation ( ) is defined as the heat of reaction in the formation of one mole of a compound from its elements taken in their standard states
For example: K(solid)+1/2Cl2+3/2O2=KClO3(solid)
= - 39,1 kJ/mol

28. Р = 101,3 kPa (760 mm Hg, 1 atm) T = 298,15К (25оC) n = 1 mole⁰ f
As a rule the measuring of H is done in standard conditions:
Р = 101,3 kPa (760 mm Hg, 1 atm)
T = 298,15К (25оC)
n = 1 mole
concentration - 1 mol/L

29. Thermochemistry
is a part of Thermodynamics which deals with reactions involving heat changes

30. C(solid) +O2(g) = CO2(g), Hо = -396 kJ
A thermochemical equation of a reaction is the equation in which heat effect, the conditions of the reaction and aggregative state are shown
C(solid) +O2(g) = CO2(g), Hо = -396 kJ

31. Hess’s Law
The heat evolved or absorbed in a given chemical process is always the same, whether the process takes place in one or several steps
This means that a given chemical reaction takes place in several steps, the overall heat change in the total process is equal to the algebraic sum of the reaction heats of the steps

32. This means that the thermochenical equations can be added to or substracted from one another, multiplied or divided by constant factors
(as can be done with ordinary algeraic equations)

33. For example:
the formation of CO2 from C and O2 can be demonstrated as following:
1.C(s)+O2(g)= CO2(g);Н1= -396 kJ
2.C(s)+1/2O2(g) = CO(g); Н2 = Х kJ
3.CO(g)+1/2O2(g)=CO2(g);Н3= -285,5 kJ
С Н CO2
Н Н3 СО

34. reaction(2)+ reaction(3) = reaction(1)
From Hess’s law:
reaction(2)+ reaction(3) = reaction(1)
H + H = H
Consequently:
H - H = H
(-285,5) =-110,5 kJ/mol О 2 О 3 О 1 О 1 О 3 О 2

35. Consequences from the Hess’s Law

36. The first consequence For isobaric process:
Under standard conditions the change of enthalpy (heat effect) of a chemical process is equal to the difference in enthalpy between products and reactants.

37. Н =nprod •H - ninit • Нr
prod
Н =nprod •H ninit • Н
aA + bB = cC + dD
For isobaric process:
H =[cHC+dHD]-[aHA+bHB]
init r О f О f О f О f

38. Enthalpy of decomposition of a chemical compound is equal, but is opposite on a mark of enthalpy of formation (under identical conditions)
Нform. = –Нdecomp.

39. The secod consequence (for organic substances)
The heat effect of an organic reaction is equal to the difference between combustion heats of reagents and combustion heats of products
Нr = ninit. • H - nprod. • Н
burn
init
burn
prod

40. The heating value of fuel is the heat, which precipitates out at combustion 1kg of dry or liquid fuel or 1 m3 of gas
The heating value of different kinds of fuel:
anthracite < wood charcoal < oil < benzine < natural gas < hydrogen

41. Entropy Bound energy = ST
The energy of the system which is not available as work, it is the bound energy
It is proportional to the temperature of the system with a factor known under the name of entropy (S)
Bound energy = ST

42. The second Law of Thermodynamics
The entropy of the universe is constantly increasing

43. What is entropy?
In statistical physics the notion of entropy is defined through the Thermodynamic Probability of the state (ώ). This is the number of different microscopic states corresponding to the macroscopic state which exists
The macroscopic state of the system is described by its macroscopic parameters - P, V, T etc. Each molecule has its own parameters, but all such molecular parameters taken at once for all the molecules of he system represent one of the numerous possible microscopic states of the system
There can be different distributions of molecular parameters between molecules but corresponding to the same values of T, P, and V.
The total number of such distributions corresponding to the same T, P, and V is the thermodynamic probability of the macroscopic state

44. Entropy is the measure of the disorder
The entropy of any system is given by the Boltzmann equation as
S = R lnώ
ώ is the number of complexions of the system.
As the number of complexions increases, the entropy increases as well
Entropy is the measure of the disorder
When the system passes from one aggregative state to another one entropy increases
Solid Liquid Gas J
mol•К

45. This is the Third Law of Thermodynamics
When temperature decreases Entropy decreases too and converts to zero under Т = 0 K (the state of ideal gas)
This is the Third Law of Thermodynamics

46. Entropy of Formation of a Substance
Standard entropy of formation of a substance (S ) is the entropy under
P = 101,3 kPa
T = 298 К
concentration – 1mol/L
Entropy can be determined as an absolute value O f

47. Calculation of Reactionary Entropy
Cgraphite + CO2(g) = 2CO(g) S
Sr = nprod. • S ninitial • S
Sr = 2· = 176 J
mol•К O f o
Prod. o
Init. J
mol•К

48. Regularities of the Entropy Changing
Sgas > Sliquid > Ssolid
S increases at the solving of a solid or a liquid substance and decreases at the solving of a gas
S increases with the increase of mass

49. The more stable and hardness a chemical bond the less entropy
The more complex of the composition of a substance the more entropy
Entropy of the elements and compounds is the periodic property

50. Spontaneous processes running into macroobjects go with losing of the part of energy on unhelpful heating of a system i.e. on orderless moving of microparticles
S > H T

51. Thermodynamic Equilibrium
In the reversible process the least amount of energy pass to unhelpful heat
Н = ТS
 Н – enthalpy factor (breaking and forming of chemical bonds)
ТS – entropy factor (the losing of energy connected with chaotic moving in equilibrium conditions
Н -ТS = 0

52. The changing of free energy is a criterion of a process directionkJ
mole
The changing of free energy is a criterion of a process direction
Н - ТS = G or G = H - TS
G is max. helpful energy, which can be done by a system in a simultaneous process and characterizes the deviation of the system from equilibrium

53. Criterion of the Process Direction
Equilibrium
G = 0; Н = ТS
Simultaneous process
G < 0; Н - ТS < 0
Nonsimultaneous process
G > 0; Н - ТS > 0

54. Change of Gibb’s Energy
 Н  S  G Direction
<0 > <0 Possible at any to
> 0 < > Impossible at any to
< < <0 & >0 Possible at low to
> > >0 & <0 Possible at high to

55. Interconnections Among Thermodinamic FunctionsН u pV G TS