Naslov
Chemical Equilibrium via Dissolution of Solids

Autori
Stilinović, Vladimir

Vrsta, podvrsta i kategorija rada
Sažeci sa skupova, sažetak, stručni

Skup
1. Croatian Workshop on Chemical Education

Mjesto i datum
Split, Hrvatska, 10.11.2010. - 14.11.2010

Vrsta sudjelovanja
Predavanje

Vrsta recenzije
Domaća recenzija

Ključne riječi
solutions; saturation; chemical equilibrium

Sažetak
Chemical equilibrium is one of the most important subjects encountered in high-school chemistry curricula. Conceptual understanding of dynamic equilibrium is essential for the understanding of numerous chemical phenomena.1 Unfortunately, this conceptual understanding is often found to be absent, not only among high-school students, but also among university chemistry students, even among students of senior years.2 One of the reasons why chemical equilibrium causes so many problems to chemistry students is the way it is introduced. There are two main problematic points. On one hand the main idea – that a reaction will proceed until equilibrium is reached is contrary to the notion that reactions proceed until all of the reactant(s) is spent, the notion which is tacitly employed throughout the teaching of chemistry (in particular in stoichiometric problems). On the other hand, equilibrium is often introduced through new reactions which were not previously introduced to the students. Worse still, many of these reactions are inappropriate for experimental demonstration (e. g. thermal decomposition of PCl5 or HI). This leads to students’ understanding of chemical equilibrium as an unusual phenomenon which is encountered in some exotic reactions, rather than a fundamental principle in the vast majority of chemical processes. A possible solution to this problem is introducing chemical equilibrium by examples which are familiar to students and also simple to demonstrate experimentally. One type of equilibrium process students are certainly familiar with is the dissolution of solids.3 The idea that a solid cannot be dissolved in a given amount of solvent at will, but that at a certain point the solution becomes saturated, is the only generally understood example of a process which does not continue until the “reactants” are entirely used up, and as such presents a natural path for experimental introduction of chemical equilibrium. This is achieved by studying solutions and their properties in order to demonstrate the behaviour characteristic of chemical equilibrium. The most obvious property of a system that has reached equilibrium is that there is no visible change in macroscopic properties of the system over time. Saturated solution is in equilibrium with the undissolved precipitate since the amount of precipitate does not change over time (if the solvent does not evaporate). On the other hand, a system comprising of unsaturated or supersaturated solution and solid precipitate is not in equilibrium since the amount of precipitate changes (decreases in the first case, increases in the second). When the amount of precipitate stops changing, the system has reached equilibrium – the solution became saturated. A number of experiments can be devised which utilise the behaviour of solutions to demonstrate the basics of chemical equilibrium. Many of these include measurement of solubility, which is easily feasible by weighing the residue after evaporation of a sample of solution of known volume. For example, it can be demonstrated that the equilibrium concentration for a given substance is unique at any given temperature. Adding a substance containing one same ion as the monitored solute will reduce its solubility (common ion effect). Measurement of solubility of NaCl in a number of HCl solutions of varying concentrations not only demonstrates this point, but also allows an additional observation – if for each case the concentrations of sodium and chloride ions are calculated, it can be shown that the increase in c(Cl–) is coupled with a the decrease of c(Na+), so that the value of c(Cl–)c(Na+) is the same in each case. In other words, c(Cl–)c(Na+) = K, where K is a constant (Fig 1). Thus the concept of equilibrium constant is reached – rather than postulating its existence as is usually done, students may discover it by themselves, analysing the data gathered from simple measurements done in the classroom. The definition of equilibrium constant can further be generalised, either by theoretical considerations or by additional experiments. Another piece of information which can be attained by solubility measurements is that on the thermal dependence of equilibrium. Having established the connection between the solubility and the equilibrium constant, it is easy to demonstrate that if solubility of a substance changes with temperature, the equilibrium constant (ant thus obviously equilibrium itself) also changes. This can easily be demonstrated by measuring the solubility of an appropriate salt (such as KNO3) at two or more different temperatures. 1. G. M. Bodner, Chem. Educ. Res. Pract., 8 (2007), 93–100. 2. M. A. Pedrosa and M. H. Dias, Chem. Educ. Res. Pract. Eur., 1 (2000) 227–236. 3. K. L. Cacciatore, J. Amado and J. J. Evans, J. Chem. Ed., 85 (2008) 251–253.

Izvorni jezik
Engleski

Znanstvena područja
Kemija

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