Strong electrolytes completely, or almost completely, ionizes or dissociates in solutions irrespective of their concentration. Molecules of strong electrolytes are completely divided on ions even in a solid condition, for example, as it is at sodium chloride. Physical-chemical (spectral and optic) researches prove the absence of neutral molecules in solutions of strong electrolytes. The constance of heat neutralization of strong acids with strong bases can be explained only by the presence of complete dissociation of the last ones. However, experimental values of osmotic pressure, decrease of freezing point and increase of boing point is lower at solutions of strong electrolytes than these values received for 100% electrolytic dissociation of molecules of given electrolyte.
In solutions of strong electrolytes each ion is surrounded by ions of opposite character that’s why the movement of ions is restricted. Ions of the strong electrolyte interact with each other in the solution because of the presence of great electrostatic forces. Owing to this ions of the same character make so called ionic atmosphere aroud the ion of opposite character.
It is important to consider also the solvation of ions. Ionic atmosphere and solvation sphere lower the movement of ions in the solution and is the cause of seeming not complete dissociation of electrolyte. Besides, in solutions of strong electrolytes at high concentrations the dissociation of ions can occur. For example, in water solution BaCI2 and AICI3 there was determined the formation of ions BaCI+, AICI2+ and others which broke apart at the dilution of a solution . To account for these influences instead of real concentration of ions C for strong electrolytes there was entered the concept of active concentration called activity –ά (Lewis, 1907). So, changes of solution’s characteristics of salt electrolytes are not connected with the change of real degree of dissociation as it takes place at weak electrolytes, and due to display of seeming degree of dissociation (άsem). A real degree of dissociation of strong electrolytes at all concentrations is equal to 1. Active concentration of electrolyte in the solution is the value showing the total influence of mutual gravitation of opposite ions, the influence of salvation of ions and others.
There is a connection between the activity and a real concentration of ions: ά=γC, where ά – activity, mol/L; C-concentration of electrolyte, mol/L; γ – factor of activity which depends on concentration. The factor of activity is usually less than 1. In diluted solution of strong electrolytes where interaction between ions is low, factor of activity is equal to 1, then ά=C.
Thus, we summarize characteristics of solutions of nonelectrolytes given above and also weak and strong electrolytes.
|
Nonelectrolytes |
Electrolytes |
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|
Weak |
Strong |
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|
The level of dissociation |
ά = 0 |
0‹ά‹1 |
ά=1, άsem≠1 |
|
Isotonic factor |
i=1 |
i=1+ά(υ-1) |
i=1+ άsem(υ-1) |
For electrolytes making the base of liquid spheres of organism are true the basic regularities of strong and weak electrolytes. Also constant’s values of dissociation (K) and limits of dissociation (ά) are calculated for weak biological electrolytes (amino acids, proteins, hemoglobin, salts of inorganic acids etc.).
The change of concentration of weak and, especially, strong electrolytes invariably involve the change of electrolytic description of the sphere and its acid-base characteristics (pH sphere). The basic weak electrolyte of organism is water. The most important electrolyte processes taking place in water spheres of organism are acid-alkaline transformations.
3. ACID-BASE CHARACTERISTICS OF SOLUTIONS
3.1. PROTOLYTIC ACID-BASE THEORY
The theory of electrolytic dissociation was created by Arrhenuis in 1864. According to this theory acids and bases are “substances dissociated in solutions with a formation of ions of hydrogen (H+) and hydroxide (OH+)”.
Indeed, typical characteristics of acids (sour taste, metal’s dilution with allocation of hydrogen, interaction with bases and etc.) and bases (colour change of indicator, reaction of neutralization with acids and etc.) are due to characteristics of ions H+ and OH-. As a result of interaction of these ions a weak electrolyte is formed – water. However studying of the reactions proceeding in nonaqueous and especially gaseous environments, our representations about acids and the bases have changed. For example, reaction NH3(g)+HCI(g)=NH4CI(t), where ammonia plays the role of the base, goes without ions OH-, instead of ions H+ in dissolvents are found: ions H3O+ - in water, C2H5OH2+ - in ethanol, NH4+ - in liquid ammonia. Further we see that according to the dissolvent many electrolytes can dissociate either like acids or like bases.
For explanation of these facts there was created a new acid-base theory – protolytic (protonic) cid-base theory (Bronsted-Lowry, 1923). According to protonic theory acids are compounds able to give protons away – donation of protons.
Bases are compounds able to accept protons - reception of protons. Acids and bases participating in the same balanced process are called conjugate acid and base:
HA+H2O 
H3O+ + A- B+
H2O BH+ + OH-
Where HA and A- - conjugate acid-base pair
BH+ and B – conjugate acid-base pair
By analogy to acid-base reactions protonic reactions are considered in a form of the sum of 2 processes, for example:
NH3 + H+ NH4+
Base 1 acid 1
HCI H+ + CI-
Acid 2 base 2
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