1. 1 The basic concepts of thermodynamics
Theoretic bases of bioenergetics. Chemical kinetics and biological processes. Electrochemistry.
Plan
1 The basic concepts of thermodynamics
2. First law of thermodynamics. Heat (Q) and Work ( W)
3. Heat capacity
ass. prof. Dmukhalska Ye. B. prepared

2. The branch of science which deals with energy changes in physical and chemical processes is called thermodynamics
Some common terms which are frequently used in the discussion of thermodynamics are:

3. common terms of thermodynamics
System
characterized
Parameter
Condition (state)
characterizes
Process

4. Classification of the thermodynamics systems according to a structure
homogeneous
heterogeneous
KNO3
KNO3
PbI2↓

5. System is a specified part of the universe which is under observation
The remaining portion of the universe which is not a part of the system is called the surroundings
The system is separated by real or imaginary boundaries.

6. Types of Systems ISOLATED Open Close
(a system which can not exchange mass and energy with the surroundings )
Open
Close
CLOSE
A system which can exchange energy but no mass with its surroundings

7. (p, T, C, viscosity, surface tension, vapour pressure)
Parameters
Extensive
(m, V, U, H, G, S, c)
The properties of the system whose value depends upon the amount of substance present in the system
Intensive
(p, T, C, viscosity, surface tension, vapour pressure)
The properties of the system whose value does not depend upon the amount of substance present in the system

8. Process is the change of all or individual parameters of the system during the length of time (the period of time)
Classification
of a process according to the constant parameter of a system are:
Isothermic process – temperature is constant, T=const
Isochoric process – volume is constant V = const.
Isobaric process – pressure of the system is constant, p = const
Adiabatic process – the system is completely isolated from the surroundings. For an adiabatic (Q=0) system of constant mass, ▲U=W

9. Classification of a process according to the releasing energy
Exothermic process is a process that releases energy as heat into its surroundings. We say that in an exothermic process energy is transferred ‘as heat’ to the surroundings. For example: a reaction of neutralization (acid + basic).
Endothermic process is a process in which energy is acquired from its surroundings as heat. Energy is transferred ‘as heat’ from the surroundings into the system. For example: the vaporization of water

10. Classification of a process according to the direction of reaction
Reversible process is a process in which the direction may be reversed at any stage by merely a small change in a variable like temperature, pressure, etc.
Irreversible process is a process which is not reversible. All natural process are irreversible

11. State function (thermodynamic function)
State of a system means the condition of the system , which is described in terms of certain observable (measurable) properties such as temperature (T), pressure (p), volume (V)
State function (thermodynamic function)
Internal energy U [J/mol]
Enthalpy H [kJ/mol] or [kJ]
Entropy S [J/mol K] or [J/K]
Gibbs energy G [J/mol] or [J]
ΔU = Uproducts - Ureactants

12. State function depends only upon the initial and final state of the system and not on the path by which the change from initial to final state is brought about.

13. Internal energy U
It is the sum of different types of energies associated with atoms and molecules such as electronic energy, nuclear energy, chemical bond energy and all type of the internal energy except potential and kinetic energies.

14. Heat (Q) is a form of energy which the system can exchange with the surroundings. If they are at different temperatures, the heat flows from higher temperature to lower temperature. Heat is expressed as Q.

15. Work (W) is said to be performed if the point of application of force is displaced in the direction of the force. It is equal to the distance through which the force acts. There are two main types of work electrical and mechanical. Electrical work is important in systems where reaction takes place between ions. Mechanical work is important specially in systems that contains gases. This is also known as pressure-volume work.

16. The work of the expanded ideal gass
WORK = Force and Distance h F s

17. If the volume of the system changes by a finite quantity from volume V1 to V2, then total work done can be obtained by integrating:
If the gas expands, V2 › V1 and work is done by the system and W is negative (we will use sign positive)
V2‹ V1 and the work is done on the system and W is positive
Note. It may be noted that many books use the opposite sign convention for work!!! (according to the IUPAC recommendation)

18. Enthalpy H Chemical reactions are generally carried out at constant pressure. ΔU gives the change in internal energy at constant volume. To express the energy changes at constant pressure, a new term called enthalpy was used. Enthalpy cannot be directly measured, but changes in it can be.

19. Enthalpy H
A thermodynamic function of a system, equivalent to the sum of the internal energy of the system plus the product of its volume multiplied by the pressure exerted on it by its surroundings.
▲H = ▲U + p▲V

20. The meaning of ▲H and ▲U 1) ▲H = ▲U in solid or liquid systems 2) ▲H >>▲U in gas system

21. The meaning of the state functions in the thermodynamic processes
Heat absorbed by the system = H positive (Q negative
Heat evolved by the system = H negative (Q positive)
The signs of W or Q are related to the internal energy change.
The meaning of the state functions in the thermodynamic processes
Exothermic process
Qv > 0, ▲U < 0
Qp > 0, ▲H < 0
2) Endothermic process
Qv < 0, ▲U > 0
Qp < 0, ▲H > 0

22. The first law of thermodynamics
Energy can neither be created nor destroyed although it may be converted from one form to another.
The given heat for the system spents on the change of the internal energy and producing the work:
Q = ▲U + W
The internal energy of the system can be changed In two ways:
By allowing heat to flow into the system or out of the system
By work done on the system or work done by the system.

23. Determination according to the process:
1) Isochoric process, V-constant
W= p▲V (▲V = V2- V1) = 0
Qv = ▲U;
2) Isobaric process, p-constant
Qp = ▲U + W , Qp = ▲U + p▲V , W=nRT (according Mendeleyev Klapeyron equation V=nRT/P)
Qp = ▲ H
3) Isothermic process, T – constant
▲U = n Cv ▲T = 0
QT = W =nRT lnV2/V1 = nRT lnP1/P2

24. Bomb calorimeter for the determination of change in internal energy
The process is carried out at constant volume, i.e., ΔV=0, then the product PΔV is also zero.
Thus, ΔU=Qv
The subscript v in Qv denotes that volume is kept constant.
Thus, the change in internal energy is equal to heat absorbed or evolved at constant temperature and constant volume

25. Thermochemistry ▲H = ▲U + p▲V ▲H = ▲U + ▲nRT
The study of the energy transferred as heat during the course of chemical reactions.
Thermochemical reactions:
H2(g) + Cl2(g) = 2HCl; ▲ H = -184,6 kJ
1/2 H2(g) + 1/2 Cl2(g) = HCl; ▲ H = -92,3 kJ/mol
▲ H is calculated for 1 mole of product
▲H = ▲U + p▲V
▲H = ▲U + ▲nRT
Energy change at constant P = Energy change at constant V + Change in the number of geseous moles * RT

26. Correlation U і Н:
If υ0, то НU: СаО + СО2 → СаСО3 If υ0, то НU: Na + H2O → NaOH + H2 If υ=0, то Н=U: H2 + Cl2 → 2HCl

27. Calculation of standard enthalpies of reactions
▲ H = ( Sum of the standard enthalpies of formation of products including their stoichiometric coefficients ) – ( Sum of the standard enthalpies of formation of reactants including their stoichiometric coefficients )
For elementary substances Н0298 = 0

29. The Hess’s law Н1 = Н2 + Н3 + Н4
Initial reactants
Н1
The products
of reaction
Н2
Н4
Н3
Н1 = Н2 + Н3 + Н4
If the volume or pressure are constant the total amount of evolved or absorbed heat depends only on the nature of the initial reactants and the final products and doesn’t depend on the passing way of reaction.

30. Conclusions from the Hess law
Нc298(the standard enthalpy of combustion) =-Нf298(the standard enthalpy of formation)
2. Н = ΣnНf298(prod.) - ΣnНf298(reactants)
3. Н = ΣnНс298(reactants) - ΣnНс298(prod.)
4.Н3=Н1-Н2
5.Н1=Н3-Н2 1 1 2 3 2 3

31. The standard enthalpy of formation ▲ H2980
is defined as the enthalpy change that takes place when one mole of the product is formed under standart condition.
C + O2 = CO2 ▲ Hf1 - x
C + ½ O2 = CO ▲ Hf2
CO + ½ O2 = CO2 ▲ Hf3
▲ Hf1 = ▲ Hf2 + ▲ Hf3

32. 1) Heat cannot be transferred from one body to a second body at a higher temperature without producing some other effect.
2) The entropy of a closed system increases with time
3) The entropy of the universe always increases in the course of every spontaneous (natural) change
The second law of thermodynamics introduces the concept of entropy and its relation with spontaneous processes
In an isolated system such as mixing of gases, there is no exchange of energy or matter between the system
And the surroundings. But due to increase in randomness, there is increase entropy ΔS›0

33. Free energy and free energy change
The maximum amount of energy available to a system during a process that can be converted into useful work
It’s denoted by symbol G and is given by
▲G = ▲H - T ▲S
where ▲G is the change of Gibbs energy (free energy)
This equation is called Gibbs equation and is very useful in predicting the spontaneity of a process.
N.B. Gibbs equation exists at constant temperature and pressure

34. Effect of Temperature on Feasibility of a Process
Exothermic reactions
For exothermic reactions ΔH is always negative, and therefore, it is favourable.
b) Endothermic reactions
For exothermic reactions ΔH is positive and always opposes the process.

35. 1) Spontaneous (irreversible) process :
▲ G < 0, ▲S > 0, ▲H < 0
2) Unspontaneous (reversible) process :
▲ G > 0, ▲S < 0, ▲H > 0
3) Equilibrium state
▲ G = 0

36. Ludwig Boltzmann’ equation:
THIRD LAW OF THERMODYNAMICS: the entropy of a perfectly pure crystal at absolute zero T=273 K is zero.
Ludwig Boltzmann’ equation:
S = klnW,
where k is Boltzmann’s constant k= R/Na
k=1.38 *10-23 J/K.
At Т→0: W = 1; S = 0.

37. Definition
The branch of science, which deals with the study oxidation-reduction reaction to produce the interconversion of chemical and electricl energy. of transition chemical energy to electrical energy is known as electrochemistry.

38. М = Мn+ + ne-
Мn+ + nе- = М
(i) Мn+ ions reflected back after colliding without any change;
(ii) Мn+ ions gaining electrons to form М (i.е. Мn+ get reduced);
(iii) Metal atoms losing electrons to form Мn+ (i.е. М gets oxidized)

39. Nernst’s equation
The dependence of cell voltage upon concentration can also be described quantitatively. The free-energy change G for any reaction is:
G =G 0+ RT ln Q
Where: Q represents the mass-action expression for an oxidation-reduction reaction
G = - nFE, and G0 = - nFE0
- nFE = - nFE0+ RT ln Q
E = E0 - RT/ nF x ln Q
R = J/K .mol
F = 96,485 С /mol

40. Мn++nе = М
Then the Nernst eqn. is applied as follows:
E = E0 – (RT/ nF) ln ([M]/ [Mn+])
where Е = electrode potential under given concentration of Мn+ ions and temperature Т
Е0 – standard electrode potential
R – gas constant
Т – temperature in К
n – number of electrons involved in the electrode reaction.

41. Standard (normal) hydrogen electrode
Pt, Н2 (g)/Н+ (Concentration)
H2 = 2H+ + 2е-
2H+ + 2е- = H2
E = E0 – (RT/ 2F) ln (pH2/ [H+]2), E0H+/H2 = 0V.
In the standard hydrogen gas electrode, hydrogen at atmospheric pressure is passed into 1 М НС1 in which foil of the platinized platinum remains immersed through which inflow or outflow of electrons takes place.

42. Since а cathode reaction is а reduction, the potential produced at such an electrode is called а reduction potential. Similarly, the potential produced at an anode is called an oxidation potential. These are known as standard reduction potentials or standard electrode potentials. They are usually tabulated for 25 С.

43. Types of electrodes
1. Metal-metal ion electrodes
2. Gas-ion electrodes
3. Metal-insoluble salt-anion electrodes
4. Inert "oxidation-reduction" electrodes
5. Membrane electrodes

44. Electrodes of the first kind.
An electrode of the first kind is а piece of pure metal that is in direct equilibrium with the cation of the metal. А single reaction is involved. For example, the equilibrium between а metal Х and its cation Х+n is:
Х+n + ne- = X (s)
for which
Еnd = Е0X+n – log ---- = Е0X+n log aX+n
n aX+n n

45. The metal - metal ion electrode consists of а metal in contact with its ions in solution. An example: silver metal immersed in а solution of silver nitrate
As a cathode: the diagram: Ag+(aq)  Ag(s)
half-reaction equation is: Ag+ (aq) + e-Ag(s)
as an anode: the diagram: Ag(s)  Ag+(aq)
half-reaction equation is: Ag(s)  Ag+(aq) + е-
Nernst’s equation:
E = E0 – (RT/ nF) ln ([Ag]/ [Agn+])

46. Electrodes of the Second Kind.
Metals not only serve as indicator electrodes for their own cations but also respond to the concentration of anions that form sparingly soluble precipitates or stable complexes with such cations.
AgCl + e- = Ag (s) + Cl- E0AgCl = V
The Nernst expression for this process is:
EAgCl = E0AgCl – log [Cl-] = pCl

47. In the metal-insoluble salt-anion electrode, а metal is in contact with one of its insoluble salts and also with а solution containing the anion of the salt. An example is the so-called silver - silver chloride electrode, written as а cathode as:
Cl- (aq)  AgCl(s)  Ag(s)
for which the cathode half-reaction is:
AgCl (s) + е-  Ag(s) + Cl- (aq)
Nernst’s equation:
E = E0 – (RT/ 1F) ln ([Ag] [Cl-]/ [AgCl])

48. An inert oxidation-reduction electrode consists of а strip, wire, or rod of an inert materiel, say, platinum, in contact with а solution, which contains ions of а substance is two different oxidation states. In the operation of this electrode the reactant not supplied by the electrode itself, nor is it introduced from outside the cell. And the product neither plates out nor leaves the cell. Instead, both reactant and product are present in solution. Thus, for the ferric - ferrous ion electrode functioning as а cathode,
Fe3+, Fe 2+(aq)  Pt(s)
the iron(III), or ferric, ion, Fe+3(aq), is reduced to the iron(II), or ferrous, ion, Fe+2(aq):
Fe+3(aq) + е- Fe+2(aq)
Nernst’s equation:
E = E0 – (RT/ 1F) ln ([Fe+2]/ [Fe+3])

49. а membrane electrode - the glass electrode. This can be depicted as:
Pt(s) Ag(s)  AgC1(s)  HC1(aq,1M)  glass 
Cell can be depicted as:
reference electrode  salt bridge  analyte solution  indicator electrode
Ecell = Eind + Eref + Ej

50. Cell potential or EMF of a cell.
The difference between the electrode potentials of the two half cell is known as electromotive force (EMF) of the cell or cell potential or cell voltage.
The EMF of the cell depends on the nature of the reactants, concentration of the solution in the two half cells, and temperature.

51. Reference electrode is electrode potential which stabile
А hydrogen electrode is seldom used as а reference electrode for day-to-day potentiometric measurements because it is somewhat inconvenient and is also а fire hazard.

52. Calomel Electrodes. A calomel electrode can be represented schematically as
Hg Hg2Cl2 (saturated), КС1 (saturate)
The electrode reaction in calomel half-cells is:
Нg2Cl2(s) + 2 е- = 2 Нg (1) + 2 Cl-
The "saturated" in а saturated calomel electrode refers to the KCl concentration. All calomel electrodes are saturated with Hg2CI2 (calomel). At 250С, the potential of the saturated calomel electrode versus the standard hydrogen electrode is V;

53. Silver/silver chloride electrodes
Silver/silver chloride electrodes. А system consists of а silver electrode immersed in а solution that is saturated in both potassium chloride and silver chloride:
Аg АgС1(saturated),KC1(saturated)
The half-reaction is AgC1(s) + е = Аg (s) + Сl-
The potential of this electrode is V at 25 0С.

54. An ideal indicator electrode responds rapidly and reproducibly to changes in the concentration of an analyte ion (or group of ions). Although no indicator electrode is absolutely specific in its response, а few are now available that are remarkably selective. There are two types of indicator electrodes: metallic and membrane.
Metallic indicator electrodes:
Electrodes of the first kind.
Electrodes of the Second Kind.
Membrane Electrodes

55. The relationship between pH and the voltage of the hydrogen elect
calomel electrode cell at 250С can be written as
Ecell Ecalomel
pH = ( log pH2)
Ecell
pH = = constant
0.0592

56. 1.Glass electrode – indicator electrode;
diagram which is:
Ag(s)  AgC1(s)  HC1(aq,1M)  glass 
2.Bulb of glass electrode.
3.Solution of unknown pH.
4.Silver-silver chloride electrode - reference electrode;diagram which is:
Cl- (aq)  AgCl(s)  Ag(s)
5.Amplifying potentiometer.

57. Thanks for attention