Compare the experimental results obtained by different groups (with different initial concentrations of the reactants and same temperature).

3. Reaction heat determination

In this case it is necessary to determine the equilibrium constant at two temperatures (recommended temperatures are 25 and 40оС). The calculation is to be done in accordance with the isobar equation (17):

2.2.TEST QUESTIONS

1.What do you know about the term “chemical equilibrium”?

2.The main characteristics of chemical equilibrium and its different

types.

3.Conditions of chemical equilibrium.

4.Mass action law and its thermodynamic background.

5.The definition of equilibrium constant and its characteristics. What are the applications of the equilibrium constant?

6.What is the difference between the equilibrium and nonequilibrium constants? What are the units of equilibrium and nonequilibrium constants?

7.The types of constants and correlation between them.

8.Equilibrium in heterogeneous processes. How one can calculate the equilibrium constant of heterogeneous reaction?

9.Isotherm equation of chemical reaction and its application.

10.The main factors that can influence on equilibrium constants.

11.Le Chatelier's principle, explain it with several examples.

12.Explain the effect of change of pressure, temperature and addition of reactants on equilibrium.

13.What do you know about the free energy ∆rG of a chemical reaction?

14.Dependence of chemical equilibrium on temperature. Isobar and isochore of a chemical reaction.

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2.3.REFERENCES

1.Ippolitov, E.G. Physical chemistry: textbook for chemistry students

of higher education institutions / E.G. Ippolitov, A.V. Artemov, V.V. Batrakov. − М.: Akademiya, 2005. – 448 p. [In Russian]

2.Stromberg, A.G. Physical chemistry: textbook for chemistry students of higher education institutions / A.G. Stromberg, D.P. Semchenko. – 5th ed., corrected. − М.: Vysshaya shkola, 2003. – 527 p. [In Russian]

3.Krasnov, K.S. Physical chemistry: textbook for higher education institutions: in 2 vol. / K.S. Krasnov. – 3rd ed., corrected. − М.: Vysshaya shkola, 2001. [In Russian]

4.Zimon, A.D. Physical chemistry: textbook for technology students of higher education institutions with non-chemical profile / A.D. Zimon, N.F. Leshchenko. – М.: Khimiya, 2000. – 318 p. [In Russian]

5.Atkins, P. Physical chemistry / P. Atkins, J. de Paula. – 9th ed. – New York: W.H. Freeman and company, 2010. – 1060 p.

6.Kudryashov, I.V. Collection of exercises and problems in physical chemistry / I.V. Kudryashov, G.S. Karetnikov. − 6th ed., corrected and revised. − М.: Vysshaya shkola, 1991. – 527 p. [In Russian]

7.Baron, N.М. Short handbook of physicochemical values: reference book / N.М. Baron [et al.]; ed. by A.A. Ravdel, A.M. Ponomaryova. – 10th ed., corrected and revised. – SPb.: Ivan Federov, 2002. – 238 p. [In Russian]

8.Ponomaryova, A.M. Phase equlibria and theory of solutions: study guidance / A.M. Ponomaryova. – Saint Petersburg: Saint Petersburg state technical university, 1992. – 159 p. [In Russian]

9.Zenin, G.S. Physical chemistry. I Part. Chemical thermodynamics: text of the lectures / G.S. Zenin, T.A. Privalova, N.V. Penkina. – SPb.: North-west state extra-mural technical university, 2001. – 77 p. [In Russian]

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3.LABORATORY PRACTICUM: SOLUTIONS

3.1.THEORETICAL INTRODUCTION

Solutions are homogeneous systems consisting of two or more components, the composition of which can vary within certain limits. Solutions cannot be considered simply as evenly distributed mixtures of the smallest particles of solutes, i.e. molecules, atoms, and ions, since these particles interact with solvent molecules. Such interaction sometimes leads to the formation of stable compounds, for example, hydrates. In particular, the above is confirmed by the evolution or absorption of heat upon dissolving various substances in a solvent. However, solutions differ from chemical compounds by irregular compositions and the absence of simple proportions. Thus, solutions are intermediate between mechanical mixtures and chemical compounds.

The general theory of solutions, which would allow one to determine the properties of solutions from the properties of its components, has not been developed yet because of complexity of these systems. Although, many relationships connect one property of solutions with others are established up today.

Generally, solutions can be liquid, solid, and gaseous (gas mixtures). One of the components of solution is considered as solvent, and the others as solutes. Liquid solutions are mostly used in chemistry; therefore a liquid component often is taken as the solvent, in which the other components (solid, liquid, or gases) are dissolved. In case of uniform mixing of two or several liquid components, any of them can be considered as solvent, but usually it is a liquid component contained in a larger amount.

One of the major characteristics of solutions is concentration. Concentration shows the ratio of a solute and solvent (by weight or volume) in a solution. There are many ways to express the concentration of solutions. Some of them are given below.

Percentage (weight %) is determined by the number of grams of a solute in 100 g of a solution. Molar concentration (molarity, sometimes simply called concentration) is determined by the number of moles (grammolecules) of a solute in 1 L of a solution. Normal concentration (normality) is determined by the number of gram-equivalents of a dissolved substance in 1 L of a solution. Molality (molar concentration) is determined by the number of moles of a solute per 1000 g of solvent. Molality has such an advantage over molarity that it does not depend on

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temperature. Molar fraction (and also molar percent) shows the number of moles of a solute (or solvent) divided by the total number of moles present in a solution. There are also other ways to express numerically the concentration of a solution, and one form of concentration can be converted into another by simple calculations. In case of gas mixtures it is convenient to express the content of one or another component via the fractional pressure, which is defined by the following expression:

where Ni is the molar fraction of the component in a gas mixture, and P is the total pressure of the mixture.

If the amount of solvent exceeds by many times the amount of a solute then such a solution is called dilute (infinitely dilute solution). This term is rather indefinite in a quantitative sense, therefore it was agreed to consider a solution infinitely dilute when interaction between the solute particles can be neglected. In case of dilute solutions, many properties of solutions depend only on the concentration of solutes and do not depend on their nature. As accepted, such properties of solutions are called colligative properties. They can be demonstrated by the following examples given below.

If a solute is characterized by its vapor pressure greater than the vapor pressure of solvent (pB >> pA), and, at the same time, both components of a solution are chemically inert, then the dissolution of such a gas in liquid follows the Henry’s law. The pressure of a volatile (gaseous) component, pB, over a solution at a constant temperature is in direct proportion to its molar fraction, NB, in a solution:

where KH is the Henry’s constant.

The dissolution of a mixture of gases in liquids follows the more general law (Dalton’s 2nd law), according to which the concentration of each gas in a solution is proportional to its partial pressure in the gas phase, as follows:

where Ci is the concentration of a gaseous component in a solution in mol/L, ki is the proportionality constant which depends on the nature of a gas, and рi is the partial pressure of a gas over the solution.

Relative drop of partial vapor pressure of solvent over a solution does not depend on the nature of a solute and equals its molar fraction (Raoult’s law):

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where p0A is the pressure of pure solvent, pA is the solvent’s pressure over a solution, and NB is the molar fraction of a solute.

The following two important relationships, which are connected with reducing the freezing temperature of solutions and increasing their boiling point, arise from the solvent’s vapor pressure drop over a solution in line with the Raoult’s law. A change in the freezing and boiling points of a solution in comparison with pure solvent is proportional to the molar concentration of a solute, as follows:

Тfreez = Т – Т0

is a decrease in the freezing temperature of a solution,

Тb.p. = Т0 – Т

is an increase in its boiling temperature, Т0 and Т are crystallization temperatures of the pure solvent and of the solution, respectively, K is the cryoscopic constant, Е is the ebullioscopic constant, and mB is the molality of a solute.

The above specified and many other relationships connected with the dependences of various properties of solutions from the concentration are obeyed only for the case of dilute or ideal solutions. An increase in the concentration of solutes (an increase in the total pressure for gas mixtures) gives rise to deviating their properties from the direct proportional dependence, therefore the concept of activity, a, of components (fugitiveness, f, or fugacity in case of gases) is introduced for the description of the properties of real solutions. If the values of concentrations in Eqs. (3.1)–(3.6) are substituted by activities and the pressures by fugacities, then the resulting equations can be applied also for the description of properties of concentrated solutions. The difficulty of using such a scheme consists in the fact that the activity and fugacity values depend on the concentration of a component, temperature and pressure, therefore they cannot be calculated but can only be determined experimentally. The concept of activity coefficient (fugacity coefficient is said rarely) is widely applied, which is defined by the following relations:

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