Dissociation of strong electrolyte

7.
1.
1 Ionic atmosphere

The experimental results of the colligation of the solution show that in the non-electrolyte solution of 1dm ³ 0.
1mol·dm ^-3 , the solute particles that can independently function are 0.
1 mol, that is, the colligation of the non-electrolyte dilute solution and the solute particles in the solution The quantity is directly proportional, and has nothing to do with the nature of the solute
.

For the electrolyte solution, the situation becomes more complicated.

Accordingly, the Swedish chemist Arrhenius believes that the electrolyte dissociation occurs in the aqueous solution, but this dissociation is incomplete and there is a dissociation equilibrium
.

Experiments have proved that for weak electrolytes such as acetic acid and ammonia, the Arrhenius theory is applicable.

In 1923, Debey and Huckel put forward the theory of strong electrolyte solution.
They believed that the strong electrolyte was completely dissociated in aqueous solution and there were no molecules
.

Due to the electrostatic interaction between the ions, the positive ions are surrounded by negative ions, and the negative ions are surrounded by positive ions.

Obviously, the greater the concentration of the solution and the greater the concentration of ions in the solution, the stronger the ionic atmosphere interaction between the ions; similarly, the greater the number of charges carried by the ions in the solution, the greater the effect of the ionic atmosphere.
If it is stronger, the true concentration of ions will not be able to function normally
.

The concept of ionic strength is usually used to measure the influence of the solution on the ions present in it, namely

In the formula, I is the ionic strength of the solution; zi is the charge number of the i kinds of ions in the solution; bi is the mass molar concentration of the i kinds of ions
.

7.

1.

In a strong electrolyte solution, due to the effect of the ionic atmosphere, the ions in the solution cannot all play the role of particles.
The concentration of the ions that actually play a role in the solution is usually called the effective concentration of ions, also known as the activity, which is often represented by a
.

If the actual concentration of ions in the strong electrolyte solution is c, then

a=f·c

In the formula, f is called the activity coefficient
.

Obviously, the actual concentration of ions a is smaller than the actual concentration of ions c, so the activity coefficient f is a value less than 1

The main factors affecting the activity coefficient f are the concentration of the solution and the number of charges of the ions in the solution
.

The greater the concentration of the solution, the stronger the effect of the ionic atmosphere in the solution, and the less able the ions will be.

When the concentration of the solution is high, the activity coefficient f and activity a must be considered
.

Because the ion concentration in the solution is higher at this time, the ion intensity is higher, and the deviation between the activity a and the actual concentration c is larger

However, in most cases (such as the dissociation equilibrium of the weak electrolyte solution and the precipitation dissolution equilibrium system to be discussed later), because the concentration of the solution in the system is generally low, the effect of the ionic atmosphere in the solution is small, and the activity a is different from the actual The concentration c is relatively close
.

Therefore, without special instructions, the influence of the ion atmosphere is not considered, and f = 1 and a = c are approximated, and the concentration is used instead of the activity
.
However, it must be clear that in the calculation of various equilibriums that will be encountered later, the concentration in the expression of the equilibrium constant should actually be the activity a, and it is precisely because of the above approximation that the concentration is used to express
.

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