Methodical instructions for the 1^st year students

Medical faculty

LESSON 5 (6 HOURS)

Themes:

1. Chemical thermodynamics. Heat effect of dissolving reactions of salts and neutralization reactions.

2. Electrode processes. Determination of redox potentials.

3. Measurement of pH of solutions with the help of potentiometric method.

Aim: To be able to determine heat effect of dissolving reactions of salts and neutralization reactions; to measure redox potentials and pH of biological liquids with the help of potentiometer; to calculate theoretical value of redox potentials of these systems.

PROFESSIONAL ORIENTATION OF STUDENTS:

Bioenergetics studies of energy transformations in the organism. The chemical energy of food products is the main source of energy for the body. It is used in the internal processes: breathing, blood circulation, metabolism, secretion, temperature control. Chemical thermodynamics is the theoretical base for bioenergetics despite lots of specific characteristics of energy metabolism in the organism. Thermochemistry renders it possible to measure energy values of food products, which is important in nutritiology.

Electrochemical phenomena are observed in human organism. The movements of muscles, heartbeat, spreading of nerve impulses are accompanied by electrochemical phenomena. The processes of biological oxidation are the primary source of energy in the body. Knowledge of the basic concepts of reduction-oxidation processes, alterations of reduction-oxidation potentials and direction of electron (or hydrogen atom) transport is crucial for understanding biochemical reduction-oxidation processes. Electrochemical analysis is widely used in medicine for determination of pH of biological liquids.

BASIC LEVEL

1. Concept of energy, heat effect of reaction, thermochemical equations (secondary school course in chemistry)

2. Redox reactions. Concept of oxidantion and reduction (secondary school course in chemistry)

3. The basics of galvanic cells (secondary school course in physics)

4. Decimal logarithm (secondary school course in mathematics).

STUDENTS’ SELF-PREPARATION PROGRAM.

Theme 1. Chemical thermodynamics. Heat effect of dissolving reactions of salts and neutralization reactions.

1. Basic concepts of chemical thermodynamics: thermodynamic system, thermodynamic process (isobaric, isochoric, isothermal).

2. The internal energy of system. Enthalpy. The first law of thermodynamics.

3. Thermochemical equations. Standard enthalpy (heat) of formation, standard enthalpy (heat) of combustion.

4. Hess’s law and its consequences. Calorimetry. Thermochemical calculations.

5. The second law of thermodynamics. Entropy.

6. Thermodynamic criteria of spontaneous processes.

7. Energetic characteristics of biochemical processes. Macroergic compounds and their importance in the organism.

Theme 2. Electrode processes. Determination of redox potentials.

1. The mechanism of electrode potential formation. Nernst equation.

2. The half-cells. Potential of redox systems.

3. Standard hydrogen electrode. Standard electrode potential.

4. Galvanic cells. Electromotive force.

5. Measurement of electrode potentials.

Theme 3. Measurement of pH of solutions with the help of potentiometric method.

1. Classification of electrodes: electrodes of the first and second kind.

2. Reference electrode: hydrogen electrode, silver-silver chloride electrode, structure and their standard electrode potential.

3. Ion-selective (membranes) electrodes. Glass electrode, its structure. Determination of pH.

4. Inert oxidation-reduction electrodes, their structure and standard electrode potential.

5. Measurement of pH by using pH meter.

6. Potentials in biological systems.

METHODOLOGY OF PRACTICAL CLASS (9⁰⁰-12⁰⁰)

Theme 1. The chemical thermodynamics. Heat effect of dissolving reactions of salts and neutralization reactions.

Work 1. Determination of the heat effect of copper sulfate dissolving

Materials. copper sulfate crystalline, calorimeter.

Protocol. Add 25 ml of distilled water (m_H2O) to the calorimeter internal cup, immerse the thermometer in water and record the temperature (t₁). Quickly add 1g of CuSO₄ crystalline to the cup and mix with stiller. Measure the maximal temperature of the obtained solution (t₂) with the thermometer.

Calculate the heat effect of salt dissolving by the formula:

Calculate the theoretic heat effect of salts dissolving using Hess’s law, if:

CuSO₄ + 5H₂O = CuSO₄·5H₂O DH₁ = - 78.2 kJ/mol

CuSO₄ + 5H₂O + aq = CuSO₄·aq DH₂ = 11.7 kJ/mol

DH (theoretic) = DH₂ + DH₁

Work 2. Determination of the heat effect of neutralization reaction

NaOH (aq) + HCl(aq) = NaCl (aq) + H₂O DH_theoretic = - 57.3 kJ/mol

Materials. 1mol/l HCl solution, 1 mol/l NaOH solution, calorimeter.

Protocol. Add 30 ml of 1 mol/l HCl solution to the calorimeter internal cup, immerse the thermometer in solution and record the temperature (t₁). Quickly add 30 ml of 1 mol/l NaOH solution to the cup, assuming that it has the same temperature and mix with stiller. Measure the maximal temperature of the obtained solution (t₂) with the thermometer.

Suppose, that densities of the solutions, both of HCl and NaOH, are equal (1 g/ml). Then the mass of the solution in the cup is the following:

Record the m_cup, t₁ and t₂ in the table:

Parameter	Name	Value
1. mass of the internal cup	mcup
2. mass of the solution	msolution	60 g
3. heat capacity of the cup	Ccup	0.753 J/ g ^0C
4. heat capacity of the solution	Csolution	4.184 J/ g ^0 C
5. initial temperature	t1
6. maximal temperature	t2

Calculate the total heat capacity of the cup and solution:

Calculate the heat effect of neutralization reaction by the formula:

Where:

M – molarity of solution;

V – volume of solution.

Theme 2. Electrode processes. Determination of redox potentials

Work 3. Measurement of redox potential of the Fe⁺³/Fe⁺² system

Redox potential of the redox system is calculated with EMF values and measured by the ionometer.

Materials. 0,1 mol/l Mohr's salt solution (FeSO₄), 0,001 mol/l iron ammonium alum solution (Fe₂(SO₄)₃)

Protocol. Collect galvanic cells with a measuring electrode (the platinum electrode) and a reference electrode (the silver-silver chloride electrode with a constant potential 0,202 ±0,002 V)

Add 10 ml of 0,1 mol/l Mohr's salt solution (contain Fe²⁺) and 10 ml of 0,001 mol/l iron ammonium alum solution (contain Fe³⁺) to the flask (or beaker). Immerse the electrodes in the solution of redox system. Measure electromotive force (EMF) of this system by the potentiometer.

Calculate the experimental value of the redox potential for redox system Fe3+, Fe2+(aq) | Pt(s) by the formula:

where: EMF value, V (measured)

e_ref = 0,202 V (silver-silver chloride electrode potential)

Calculate the theoretical value of the redox potential by the Nernst equation for this system:

half-reaction: Fe⁺³(aq) + е^- ®Fe⁺²(aq)

Work 4. Measurement of redox potential of the MnO₄^-/Mn²⁺ system

Materials. 0,1 mol/l KMnO₄ solution, 0,01 mol/l MnSO₄ solution, H₂SO₄ conc.

Protocol. Collect galvanic cells with a measuring electrode (the platinum electrode) and a reference electrode (the silver-silver chloride electrode with a constant potential 0,202 ±0,002 V)

Add 25 ml of 0,1 mol/l KMnO₄ solution, 5 ml of 0,01 mol/l MnSO₄ solution and 5 drops of H₂SO₄ conc. (pH=1) to the flask (or beaker). Immerse the electrodes in the solution of redox system. Measure electromotive force (EMF) of this system by the potentiometer.

Calculate the experimental value of the redox potential for redox system MnO₄^-, Mn²⁺| Pt(s) by the formula:

where: EMF value, V (measured)

e_ref = 0,202 V (silver-silver chloride electrode potential)

Calculate the theoretical value of the redox potential by the Nernst equation for this system:

half-reaction: MnO₄^- + 8H⁺ + 5e ® Mn²⁺ + 4H₂O

Theme 3. Measurement of pH of solutions with the help of potentiometric method.

Work 5. Measurement of the pH of solutions and biological liquids

The pH value of the solution is measured using pH electrode. It essentially consists of a pair of electrodes: measuring (glass electrode) and reference electrode (silver-silver chloride electrode), both dipped in the solution of unknown pH.

Protocol. Measure pH of solutions: acidic solution, gastric juice, urine, intestinal juice, alkaline solution using pH meter. Record the pH value to the table, calculate the pOH values, [H₃O⁺] and [OH^-].

Example of calculation:

pH + pOH = 14

pOH = 14 - pH

The hydronium ion concentration ([H₃O⁺]) can be found from pH value by the reverse mathematical operation used to determine pH.

[H₃O⁺] = 10^-pH  or  [H₃O⁺] = antilog (- pH)

pH = 5,46

5,46 = - log [H₃O⁺]

- 5,46 = log [H₃O⁺]

[H₃O⁺] = 10^-5.,46 = 3,47.10^-6 mol/l

(On a calculator, calculate: press the number -0.54, press button 10^x (or two buttons INV and LOG), press button =, get the value antilogarithm 3,47·10^-6).

If you know the value of [H₃O⁺], find the value [OH^-]:

K_w= [H₃O⁺][OH-] = 1,0 ·10^-14

Solution	pH	pOH	[H3O^+]	[OH^-]	Medium
acid solution
gastric juice
urine
intestinal juice
alkaline solution

SEMINAR DISCUSSION OF THE THEORETICAL ISSUES (12³⁰-14⁰⁰).

TEST EVALUATION AND SITUATIONAL TASKS.

Multiple choice tests

1. Choose standard state of the system:

A. 101,3 kPa, 0⁰K.

B. 101,3 kPa, 273⁰K.

C. 101,3 kPa, 298⁰K.

D. 50 kPa, 273⁰K.

E. 50 kPa, 298⁰K.

2. The concept “In any process of an isolated system, the total energy remains the same” is:

A. The first law of thermodynamics.

B. The second law of thermodynamics.

C. The third law of thermodynamics.

D. Hess’s law.

E. Kirchhoff’s law.

3. Thermodynamics systems are classified by the number of phases into:

A. Exogenous and endogenous.

B. Solid, liquid, gas.

C. Homogeneous and heterogeneous.

D. Isolated, opened, closed

E. Reversible and irreversible.

4. Diagram of silver-silver chloride electrode is:

A. Pt, H₂/2H⁺.

B. Ме/Меⁿ⁺.

C. Zn/Zn²⁺.

D. Ag/AgCl, KCl.

E. Hg/Hg₂Cl₂, KCl.

5. Choose the formula of Nernst equation:

A. E = E⁰catode - E⁰anode.

B. E = E⁰ - (RT/ nF) ln ([M]/ [Mn+]).

C. pH = - log [H+].

D. pH = 14- pOH.

E. E = 14 - RT/nF

Examples of tasks solving

Task 1. The heat of formation of carbohydrates in the human body is 4.1 kcal/g. Daily needs for carbohydrates for students are 383g. Calculate the daily needs for carbohydrates for the student to get energy.

Answer:

∆Нform (carb) = 4,1 kcal/g m (carb) = 383 g	1. To solve this task, we use direct proportion: 4,1 kcal — 1 g carbohydrate X kcal — 383 g carbohydrates Answer: 1570,3 kcal
∆Н (day) = ?

Task 2. Calculate the heat effect of reaction of glucose oxidation, if ∆Н_formation of glucose - 1272.45 kJ/mol, carbon dioxide - 393.6 kJ/mol, water - 285.9 kJ/mol.

Answer:

∆Нform (C6H12O6) = -1272,4 kJ/mol ∆Нform (CO2) = -393,6 kJ/mol ∆Нform (H2O) = -285,9 kJ/mol	1. The reaction of oxidation of glucose: C6H12O6 + 6O2 → 6CO2 + 6H2O; 2. Use the equation of the first consequences of the Hess’s law: ∆Нformation = ∑ ∆Нf.prod ─ ∑ ∆Нf. reactant 3. Substitute these ∆Нformation reaction products and reactants, taking into account the stoichiometric coefficients: Answer: ∆Нreaction = - 2804,55 kJ / mol, exothermic reaction.
∆Нreaction = ?

Task 3. Can the following reaction occur spontaneously if Gibbs energy of SiO_2(solid) = -803,75 kJ/mol, NaOH_(liquid)= - 419,5 kJ/mol, Na₂SiO_3(solid) = -1427,8 kJ/mol, H₂O _(liquid) = -237,5 kJ/mol?

Answer:

∆G^0(SiO2(solid)) = -803,75 kJ/mol ∆G^0(NaOH(liquid))= - 419,5 kJ/mol, ∆G^0(Na2SiO3(solid))= -1427,8 kJ/mol, ∆G^0(H2O (liquid))= -237,5 kJ/mol	Since ∆G<0, it is spontaneous Answer: ∆G= -22,5 kJ/mol, reaction spontaneous
∆Greaction=?

Task 4. The heat effect of evaporation of 1 mol of water is 40.7 kJ. How much heat is spent a day if 800.0 g of water is evaporated through the skin?

Answer:

ν(H2O) = 1 mol ∆Нevaporation (H2O) = 40,7 kJ/mol m (H2O) = 800 g	1. Find how many moles do 800 g of water have: 2. To determine heat released according the proportion: 1 mol — 40,7 kJ 44,4 mol — X kJ X= 1807,08 kJ Answer: 1807,08 kJ
∆Н = ?

Task 5. Calculate the potential at 25^oC for the following cell.

Oxidation: Cu Cu²⁺ + 2 e^- E^o_ox = 0.3370 V

Reduction:     Ag⁺ + e^- Ag E^o_red = 0.7996 V

[Cu^2+] = 0,024 mol/l [Ag^+] = 0,048 mol/l E^oox = 0,3370 V E^ored = 0,7996 V

1. Equation for the reaction. Cu(s) + 2 Ag^+(aq) Cu^2+(aq) + 2 Ag(s) 2. Calculate the standard-state cell potential E^o = E^ored - E^oox = 0,4626 V 3. Nernst equation for this cell, note that n is 2 because two electrons are transferred in the balanced equation for the reaction: 4. Substituting the concentrations of Cu^2+ and Ag^+ ions into this equation gives the value for the cell potential. Answer: E = 0,373 V

E = ?

Task 6. Concentration of hydrogen ions in solution is 0,001 mol/l. Calculate the pH of this solution.

Answer:

[H^+] = 0,001 mol/l	pH = -log 0,001= - log 10^-3 pH = 3 Answer: pH of solution is 3
pH=?

Task 7. Electrode potential of the hydrogen electrode in solution is 0,1186 V. Calculate the pH of this solution and hydrogen ion concentration.

Answer:

E = -0,1186 V	Nernst equation for hydrogen electrode: Therefore: E^0 hydrogen electrode = 0 Answer: pH=2; [H+] = 0,01 mol/l
pH=? [H^+] = ?

Tasks

Task 8. Calculate the heat effect of reaction if standard enthalpy of formation for CO = -110 kJ/mol, CH₄= -74,9 kJ/mol, H₂O = -241,8 kJ/mol.

Task 9. 100 g of codfish contain 11,6 g of proteins. The calorie content of 1 g of proteins is 4,1 kcal. Calculate the calorie content of proteins of codfish.

Task 10. Calculate the energetic value of 200 g of margarine, that contains 0,3 % of proteins, 82,3 % of fats, and 1 % of carbohydrates, if 1 g of protein or 1 g of carbohydrate releases 17,18 kJ in the organism, and 1 g of fat releases 38,97 kJ.

Task 11. The heat of combustion of fat in the human body is 9.3 kcal/g. Average daily needs for fats are106 g for male students. Calculate the daily energy of fats for male students?

Task 12. Calculate the change of Gibbs free energy of glucose oxidation under standard conditions, if the standard Gibbs energy of glucose is 910 kJ/mol, water - 237 kJ/mol, carbon dioxide - 394 kJ/mol. Can this reaction occur spontaneously?

Task 13. Calculate pH of solution, in which hydrogen ion concentration is equal to 0,005 mol/l.

Task 14. Calculate hydrogen ion concentration in solution, in which potential of hydrogen electrode is equal to 0,082 V.

Task 15. The glass electrode is placed in solution with pH=2. Calculate electrode potential of a glass electrode.

Task 16. Calculate pH of solution, in which hydroxide ion (OH^-) concentration is equal to 0,0003 mol/l

Student should know:

· the First and Second Laws of Thermodynamics;

· Hess’s law and its consequences;

· concept of internal energy and enthalpy, entropy, standard enthalpy of formation, standard enthalpy of combustion, Gibbs’ energy;

· thermodynamic criteria of spontaneous processes

· Nernst’s equation;

· structure of electrodes;

· structure of galvanic cell;

· principle of measurement of pH with the glass electrode;

Student should be able to:

· determine heat effects of reaction of dissolving of salts and neutralization reaction;

· calculate heat effects of reactions with standard enthalpies of formation and combustion;

· calculate food value of products from energetic value of proteins, fats and carbohydrates;

· measure pH of solutions and biological liquids;

· calculate values of electrode potential with the Nernst equation;

· calculate pH, pOH, [H⁺], [OH^-] of solution.

Correct answers of test evaluations and situational tasks:

Multiple choice tests

1. C; 2. A; 3. C; 4. D; 5. B

Tasks

8. -206,2 kJ/mol.

9. 47,56 kcal

10. 6459 kJ

11. 985,8 kcal

12. ∆G = -2876, reaction spontaneos

13. pH = 2,3

14. [H⁺] = 4·10^-2 mol/l

15. E = -0,118 V

16. pH = 10,48

References:

1. http://intranet.tdmu.edu.ua/ Student’s facilities /Lecture presentations /Department of General Chemistry/ English / medical /1 course/ medical chemistry/ 06. Theoretic bases of bioenergetics

2. http://intranet.tdmu.edu.ua/ Student’s facilities /Practical classes materials /Department of General Chemistry/ medical chemistry/ medical/ 1 course/ English/ 05. Chemical thermodynamics

3. Atkins, Peter W.; de Paula, Julio. Physical Chemistry (9th ed.). Oxford University Press. – 2010. –  ISBN 978-0-19-954337-3.

4. Atkins, Peter W.; de Paula, Julio. Physical Chemistry for the Life Sciences (2nd ed.). W.H. Freeman & Company. – 2011.

5. Petrucci, Ralph H., Harwood, William S., Herring, F. G., and Madura Jeffrey D. General Chemistry: Principles & Modern Applications. 9th Ed. – New Jersey: Pearson Education, Inc., – 2007.

6. Ebbing, D.D. General Chemistry. Fifth Edition. – Houghton Mifflin Co., Boston, MA, – 2002. chapter 19, pp. 786-810.

7. Chemistry” 3th ed. J. Mc Murry and R. Fay; – Prentice Hall, Upper Saddle River, New Jersey 07458. – 2001. – ISBN 0-13-087205-9;

Methodical instructions have been worked out by:

Associate Prof. Kyryliv M. V.

Methodical instructions were discussed and adopted at the Department sitting

27 August 2013. Minute № 1