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# Salt Solutions

Salt solutions are aqueous medium where certain soluble metallic salts are mixed in water which finally results in either acidic or basic depending upon the composition of the soluble salt. These salts are nothing but ionic solids which consists of both cations and anions arranged in crystalline lattice. These ionic salts dissociate into ions when they are mixed into water.

The anions of the salts which are derived from weak acids are basic in nature while the ones which has weak base and strong acid will have an acidic outlook.

 Related Calculators Calculate the Molality of the Solution Equation Solution Calculator Ionic Strength of a Solution Mole Fraction of a Solution

## Salt Solution Definition

These are aqueous medium having soluble metallic salts mixed in water resulting in either having basic characteristics or acidic characteristics depending upon the derived salt. Strong acids and strong base will never hydrolyze as they neutralise each other and eventually dissociate into just spectator ions.

These salts could made up of weak acid and strong base, weak base and strong acid, weak acid and weak base or even strong base and strong acid. But in case both are strong the resulting pH will be neutral as neither $H_{+}$ are altered by the cations nor $H_{+}$ are attracted by anions.

In case the salt is basic in nature the concentration of hydroxide ions get enhanced and when the salt is acidic in nature the concentration of hydrogen ions gets enhanced when dissolved in water. The solutions of salts and water remain as either conjugate base or conjugate acid. For any neutral salt solutions the salt has to be from a strong acid and strong base. For example Sodium chloride which is a salt of a strong base and strong acid.

$HCl (aq) + NaOH (aq) \rightarrow NaCl (aq) + H_{2}O (l)$.

In this case $Na_{+}$ is weak conjugate acid of strong base, while $Cl^{-}$ is weak conjugate base of a strong acid.

## Acidity or Basicity of Salt Solutions

So now what exactly decides the acidity and basicity of the solutions? In order to understand this we need to understand the acidity and basicity of the ions which comprise the salt. If the salt has a strong base and weak acid then the solution thus formed will be basic in nature while a salt comprising a strong acid and weak base will have an acidic look.

For example Potassium acetate having Potassium from group I the ion does not hydrolyze but as the acetate ion being conjugate of weak acid has a basic outlook. The solution thus formed is basic in nature.

$K^{+} (aq) + H_{2}O (l) \rightarrow No \ reaction$

$CH_{3}COO^{-} (aq) + H_{2}O (l) \leftrightarrow CH_{3} COOH + OH^{-} (aq)$

Rules that governs the nature of salt solution being acidic, basic or neutral in nature are as follows.

 Type of Salt Weather ions hydrolyze or not Nature of salt solution Salt of strong base and strong acid  Example KCl None of the ions hydrolyze Salt of strong base KOH and strong acid HCl Solution is neutral Salt of strong base and weak base  Example $CH_{3}COONa$ Salt of strong base NaOH and weak acid $CH_{3}COOH$ Solution is basic Salt of weak base and strong acid  Example $Mg(NO_{3})_{2}$ Salt of weak base $Mg(OH)_{2}$ and strong acid $HNO_{3}$ Solution is acidic

For a salt which has been derived from weak base and weak acid the deciding factor for type of solution is little different. The rules mentioned above does not cover entirely for salts which has acidic cation and basic anion and in such cases the cation Ka and anion Kb needs to be compared. If the Ka (acid dissociation constant) is greater than Kb, the salt solution is acidic nature while in case where the Kb (base dissociation constant) is greater than Ka the salt solution is basic nature.
• When Ka > Kb the pH will be somewhere hovering above 6 but below 7 and hence acidic; example Ammonium Fluoride $NH_{4}F$
• When Kb > Ka the pH will be little above 7 and hence making it basic; Ammonium hypochlorite, $NH_{4}ClO$

## Conductivity of Salt Solutions

Conductivity of salt solutions or electrolyte depends mainly on two factors:
• The concentration of the free charge carriers
• The ability of the charge carriers or ions to move in an electric field
The charge carriers tend to be present at high concentration when the dissolved salt concentration in an electrolyte phase is quite high and hence higher the salt concentrations the greater is the conductivity. This might fail only at a very high salt concentrations when it approaches the solubility limit as the property of electrolyte phase changes drastically due to limitations of the available solvent. As the viscosity rises the limits the ionic mobility and finally diminish conductivity.

The electrolytes could be broadly considered into two categories.
• Strong electrolytes
• Weak electrolytes
Weak electrolytes are usually exist in a partially dissociated state and the dissolved electrolyte remain in uncharged state or in neutral molecule form. Example of weak electrolyte is acetic acid in water. Strong electrolytes dissociate entirely and the dissolved electrolyte remain in a charged state.

## Is Dissolving Salt a Physical Change?

The categorisation of any change depends upon the factor of whether or not we could get back the original form or substances. If we can get back the original substance after a change has taken place then we categorise them as physical and if the change that has taken place is irreversible then we call them as chemical change.

In case of dissolving a salt in water, the original salt can be recovered after the salt solution is heated. Hence, the dissolving salt is a physical change.

## pH of Salt Solutions

The hydrogen and hydroxyl ion concentration in dilute solution are often small. The $H_{3}O^{+}$ and $OH^{-}$ ions enter into many equilibria in addition to the ionization of water and hence it is essential to specify the concentration of aqueous solutions and these may vary from high values to very small values.

The concentration of these ions in mole per litre are expressed as negative power of 10.

In order to express hydrogen concentration in a more compact and convenient manner without involving negative exponent the pH scale was introduced.

The pH is thus the negative logarithm of hydrogen ion concentration.

The relation is thus, pH =  $- log_{10} [H_{3}O^{+}]$ = $log_{10}$ $\frac{1}{[H_{3}O^{+}]}$

And, pOH = $- log10 [OH-]$ = $log_{10}$ $\frac{1}{[OH^{-}]}$

The product of water, Kw = $[H_{3}O^{+}] [OH^{-}]$

Log [H3O+] = $\frac{1}{2}$ $[log K_{w} + log K_{a} – log K_{b}]$

Hence, pH = $\frac{1}{2}$ $pK_{w}$ + $\frac{1}{2}$ $pK_{a}$ – $\frac{1}{2}$ $pK_{b}$

The pH of salt solutions depend upon the pK values of acid and the base.

If the pKa < pKb then pH of the solution will be less than $\frac{1}{2}$ pKw and ultimately result in acidic solution.

If the pKa > pKb the pH of solution will be more than $\frac{1}{2}$ pKw and hence the solution will take an alkaline nature.

However, when pKa = pKb then the pH of the solution will be equal to $\frac{1}{2}$ pKw and the solution will take a neutral outlook.

Thus, it’s clear that pH of salt solution is independent of the salt concentration.

Example:
Calculate the pH of a 0.1 M ammonium acetate solution where, Ka = 1.85 $\times 10^{-5}$, Kb = 1.85 $\times 10^{-5}$ at 298 K.

Solution:
According to the relation, pH = $\frac{1}{2}$ $pK_{w}$ + $\frac{1}{2}$ $pK_{a}$ – $\frac{1}{2}$ $pK_{b}$

The pH of the solution is $\frac{1}{2}$ (14) + $\frac{1}{2}$ (4.73) – $\frac{1}{2}$ (4.73) = 7.0
The solution thus is neutral.
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