Osmotic pressure electrolyte solutions

Numerous measurements of the conductivity of aqueous solutions performed by the school of Friedrich Kohhansch (1840-1910) and the investigations of Jacobns van t Hoff (1852-1911 Nobel prize, 1901) on the osmotic pressure of solutions led the young Swedish physicist Svante August Arrhenius (1859-1927 Nobel prize, 1903) to establish in 1884 in his thesis the main ideas of his famous theory of electrolytic dissociation of acids, alkalis, and salts in solutions. Despite the sceptitism of some chemists, this theory was generally accepted toward the end of the centnry. 

How are we to understand this odd result The answer is easy when we remember that osmotic pressure counts solute particles. The macroion cannot pass through the semiperme-able membrane. In the absence of added salt, its counterions will not pass through the membrane either since the electroneutrality of the solution must be maintained. Therefore the equilibrium pressure is that associated with z + 1) particles. Failure to consider the presence of the counterions will lead to the interpretation of a low molecular weight for the colloid. As we already saw, the presence of increasing amounts of salt leads to a leveling off of the ion concentrations on the two sides of the membrane. The effect of the charge on the macroion is essentially swamped out with increasing electrolyte. 

Data collected with membranes of this type played an important part in the formulation of present-day solution theory—so much so that the authors have used this theory without hesitation to compute osmotic pressures of solutions whose osmotic pressures have never been precisely measured. Such a solution is sea water. The copper ferrocyanide membrane is leaky to solutions of strong electrolytes. Some data have been obtained on weak solutions of strong electrolytes by the Townend method (16), but no one has made precise measurements on the osmotic pressure of sea water. 

Until the middle of the nineteenth century colloidal systems were regarded as being outside the realm of well behaved chemical systems because they did not behave in a manner expected of an aqueous solution. Such physico-chemical properties of colloidal solutions as the exhibition of osmotic pressure, electrolytic conductance, lowering of vapour pressure, elevation of boiling point, depression of freezing point etc. were different. Howwer, colloids must constitute extremely well behaved systems because life is a manifestation of various colloidal states. All protoplasm is in colloidal form. Most of the biologieal fluids, notably blood, lymph, milk, bile, and digestive secretions are colloidal solutions. Moreover, the biomembranes may themselves be considered to be a manifestation of the colloidal state. 

These results show more clearly than Fq. (8.126)-of which they are special cases-the effect of charge and indifferent electrolyte concentration on the osmotic pressure of the solution. In terms of the determination of molecular weight of a polyelectrolyte by osmometry. ... 

No experiments appear to have been made with such cells, although the equation has been verified with oxygen at different partial pressures in admixture with nitrogen, with platinum electrodes and hot solid glass as electrolyte (Haber and Moser). A similar case is that of two amalgams of a metal, of different concentrations, as electrodes, and a solution of a salt of the metal as electrolyte (G. Meyer, 1891). Here we must take the osmotic pressures of the metals in the amalgams, Pi, P2, and, for an 7i-valent metal ... 

Chapters 7 to 9 apply the thermodynamic relationships to mixtures, to phase equilibria, and to chemical equilibrium. In Chapter 7, both nonelectrolyte and electrolyte solutions are described, including the properties of ideal mixtures. The Debye-Hiickel theory is developed and applied to the electrolyte solutions. Thermal properties and osmotic pressure are also described. In Chapter 8, the principles of phase equilibria of pure substances and of mixtures are presented. The phase rule, Clapeyron equation, and phase diagrams are used extensively in the description of representative systems. Chapter 9 uses thermodynamics to describe chemical equilibrium. The equilibrium constant and its relationship to pressure, temperature, and activity is developed, as are the basic equations that apply to electrochemical cells. Examples are given that demonstrate the use of thermodynamics in predicting equilibrium conditions and cell voltages. 

It remains to evaluate the quantity c — Cs. Since an explicit general solution is not to be had, we resort to the consideration of special cases. First, suppose that the external electrolyte concentration Cs is very small compared with the concentration ic /z- of the ge-gen ions belonging to the polymer and occurring in the gel. Then the second term in the left-hand member of Eq. (45) may be neglected in comparison with the first. Furthermore, the very large ionic osmotic pressures developed in such cases will cause V2m to be very small, thus justifying adoption of the dilute solution approximations (see, for example, Eq. 40) for the right-hand member. The equilibrium relation reduces in this case to... 

If 0.6 N lithium bromide is added to the solution of the polyelectrolyte and also to the solvent on the opposite side of the osmometer membrane, the lowermost set of points in Fig. 145 (lower and left scales) is observed. The anion concentration inside and outside the coil is now so similar that there is little tendency for the bromide ions belonging to the polymer to migrate outside the coil. Hence the osmotic pressure behaves normally in the sense that each poly electrolyte molecule contributes essentially only one osmotic unit. The izjc intercept is lower than that for the parent poly-(vinylpyridine) owing to the increase in molecular weight through addition of a molecule of butyl bromide to each unit. 

Fluids can be classified further according to their tonicity. Isotonic solutions (i.e., normal saline or 0.9% sodium chloride [NaCl]) have a tonicity equal to that of the ICF (approximately 310 mEq/L or 310 mmol/L) and do not shift the distribution of water between the ECF and the ICF. Because hypertonic solutions (i.e., hypertonic saline or 3% NaCl) have greater tonicity than the ICF (greater than 376 mEq/L or 376 mmol/L), they draw water from the ICF into the ECF. In contrast, hypotonic solutions (i.e., 0.45% NaCl) have less tonicity than the ICF (less than 250 mEq/L or 250 mmol/L) leading to an osmotic pressure gradient that pulls water from the ECF into the ICF. The tonicity, electrolyte content, and glucose content of selected fluids are shown in Table 24—3. 

The activity coefficient of the solvent remains close to unity up to quite high electrolyte concentrations e.g. the activity coefficient for water in an aqueous solution of 2 m KC1 at 25°C equals y0x = 1.004, while the value for potassium chloride in this solution is y tX = 0.614, indicating a quite large deviation from the ideal behaviour. Thus, the activity coefficient of the solvent is not a suitable characteristic of the real behaviour of solutions of electrolytes. If the deviation from ideal behaviour is to be expressed in terms of quantities connected with the solvent, then the osmotic coefficient is employed. The osmotic pressure of the system is denoted as jz and the hypothetical osmotic pressure of a solution with the same composition that would behave ideally as jt. The equations for the osmotic pressures jt and jt are obtained from the equilibrium condition of the pure solvent and of the solution. Under equilibrium conditions the chemical potential of the pure solvent, which is equal to the standard chemical potential at the pressure p, is equal to the chemical potential of the solvent in the solution under the osmotic pressure jt,... 

Van t Hoff introduced the correction factor i for electrolyte solutions the measured quantity (e.g. the osmotic pressure, Jt) must be divided by this factor to obtain agreement with the theory of dilute solutions of nonelectrolytes (jt/i = RTc). For the dilute solutions of some electrolytes (now called strong), this factor approaches small integers. Thus, for a dilute sodium chloride solution with concentration c, an osmotic pressure of 2RTc was always measured, which could readily be explained by the fact that the solution, in fact, actually contains twice the number of species corresponding to concentration c calculated in the usual manner from the weighed amount of substance dissolved in the solution. Small deviations from integral numbers were attributed to experimental errors (they are now attributed to the effect of the activity coefficient). 

The osmotic pressure of an electrolyte solution jt can be considered as the ideal osmotic pressure jt decreased by the pressure jrel resulting from electric cohesion between ions. The work connected with a change in the concentration of the solution is n dV = jt dV — jrel dV. The electric part of this work is then JteldV = dWcl, and thus jzc] = (dWei/dV)T,n. The osmotic coefficient 0 is given by the ratio jt/jt, from which it follows that... 

Let us assume a solution of a non-electrolyte in water, separated from the pure solvent—water—by a semiper-meable membrane forming a piston (Fig. 8). Water enters the solution through the membrane and raises the piston, i.e., the solution can do work or possesses potential energy owing to its osmotic pressure. If the membrane is removed, the osmotic pressure causes diffusion until (if no other forces are active) the solute is uniformly distributed through the solvent. Osmotic pressure is, therefore, a factor tending to bring about uniform concentration. 

All eukaryote cells are faced with differences in intracellular solute composition when compared with the external environment. Many eukaryotes live in seawater, and have cells which are either bathed in seawater directly, or have an extracellular body fluid which is broadly similar to seawater [3]. Osmoregulation and body fluid composition in animals has been extensively reviewed (e.g. [3,15-21]), and reveals that many marine invertebrates have body fluids that are iso-osmotic with seawater, but may regulate some electrolytes (e.g. SO2-) at lower levels than seawater. Most vertebrates have a body fluid osmotic pressure (about 320mOsmkg 1), which is about one-third of that in seawater (lOOOmOsmkg ), and also regulate some electrolytes in body fluids at... 

This latter expression allows us to compute all the excess properties of dilute electrolytic solutions for instance, the excess osmotic pressure is determined by Eq. (138). The most remarkable result is of course that all these thermodynamic properties are non-anaiytic functions of the concentration ... 

Moreover the ionisation of the electrolyte groups adds to a number of unusual effects in the presence of small amounts of added salt. The intensity of light scattering decreases due to the ordering of the molecules in solution and the Osmotic pressure and ultracentrifugation behaviour are determined predominantly by the total charge on the molecule. 

Since osmotic pressure depends upon the number of particles of solute(s) in solution, the osmotic pressure of an electrolyte is directly proportional to the degree (or extent) of dissociation. The dissociation factor, symbolized by the letter i, can be calculated by dividing the total number of particles (which include undissociated molecules and ions) in a solution by the number of particles before dissociation, i.e.,... 

Electrolytes regulate body water volumes by establishing osmotic pressure which is proportional to the total number of particles in solution. The osmotic pressure of a solution is expressed in units of milliosmoles (mOsm). Osmolar concentrations reflects the number of particles (molecules as well as ions) of total solutes per volume of solution, which in turn determines the osmotic pressure of the solution. 

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