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# Weak Acids

Svante Arrhenius was the first to find the essential characteristics of acids and were fist recognised as something which tastes sour. According to Arrhenius postulate, acids are substances which produced hydrogen H+ ions in aqueous medium and this was considered to be a major concept in identifying and also quantifying the chemistry of acids. Although this was considered as a path breaking idea, yet the limitations of this concept cropped up because of the application of only aqueous medium.

Later on the Bronsted Lowry model gave the proton (H+) donor which paved way to look at the acids from new perspective of strong and weak characteristics. The amount of hydrogen ions dissociating during any chemical change gives an idea about its strnegth. More the dissciation stronger the acid is categorised as while the partial disscoiation gives an idea of its weakness. Weak acids will have pH closer to 7.0 and dissociate partially.

## Weak Acids Definition

Lavoisier considered oxygen as a common component for all acids and named the element as such. Chemical analysis by Humphrey Davy proved beyond doubt that oxygen is not part of many of the strong acids around us. Later Lavoisier suggested that it’s Hydrogen which is part of all the acids that we get to see and he defined acids as compounds which contain hydrogen atom which usually gets replaced by metals when reacted. Arrhenius whereas defined acid as a substance which dissociates in water to provide hydrogen ions. Hydrogen ion is nothing but a hydrogen atom without its electron. Hydrogen ion thus is considered as equivalent to a proton. For any weak acid the pattern remains same but the only difference that we get to see is the partial loss of hydrogen ion. The acids will dissociate but only a part of hydrogen atoms in their combined state lose electron and dissociate, while the rest of the hydrogen atoms remain with the main radical.

Several water molecules may actually be associated with single hydrogen ion to produce ions like penta-hydro-dioxy $H_{5}O_{2}^{+}$ and $H_{7}O_{3} ^{+}$ in various acids. In aqueous medium the H+ should be interpreted as a hydronium ion rather than a hydrogen ion. As per the definition for strong acids, it’s the limits of ionisation in water which determines the strength of acids and hence the weaker acids can be considered to be solutes that react reversibly with water to form the hydronium ions. For strong acids the main ones are HCl, HBr, and $H_{2}SO_{4}$ etc. while weak ones are mostly organic acids found in animals and plants. We can see two specific categories containing an ionisable hydrogen atom like $HNO_{2}$ and cations like ammonium ions $(NH_{4}^{+})$.

Examples of weak acids are acetic acid $(CH_{3}COOH)$, Boric acid $(H_{3}BO_{3})$ which shows the hydrogen atoms remaining with the base compound and only partial hydrogen dissociation takes place. The strength of an acid completely depends upon the dissociative power of any species. The more quickly and fully dissociate a species, the stronger it is considered as an acid. Hence these organic compounds having hydrogen atoms as part of the acid species gives out only a part of the hydrogen atoms and hence are considered as weak acids. The relative strength of these acids is specified by their reaction with water. Sevral of these weak acids are carbon based or carbon atoms in main structure and having similar structure. To an extent we could say that weak and very weak acids are almost insoluble in water and hence solvent other than water is needed to turn them into a solution. Compared to this water react with HCl readily reacts with water while formic acid also reacts with water but is much weaker than HCl acid.

Weak acids partly ionise the water and strong acids ionises the solvent medium compltely.
Weak acid - partially dissociated ions
Strong acid - completely dissociated ions
A quantitative measure of the degree of dissociation is shown by equilibrium constant for acid and the higher the equilibrium constant, greater ther percent dissociation of acid. Which indicates that species with higher equilibrium a stronger acid and anything with lower equilibrium is weak acid.

## Weak Acids List

The following common weak acid list is provided below.

 Weak Acid Weak acid formula Carbonic acid $H_{2}CO_{3}$ Acetic acid $CH_{3}COOH$ Chromic acid $H_{2}C_{2}O_{4}$ Boric acid $H_{3}BO_{3}$ Oxalic acid $(COOH)_{2}$ Phosphoric acid $H_{3}PO_{4}$ Lactic acid $H(C_{3}H_{3}O_{3})$ Hydrocyanic acid HCN

## pH of Weak Acids

The pH of any weak acid is calculated from the extent of ionisation and hence the pH of any weak acid can completely depend upon its concentration and also its pKa value. While calculating the pH of a weak acid both these factors are very vital.

Steps of finding out the pH through these two factors:
• For any weak acid the equilibrium equation in aqueous form is essential
• The equilibrium expression for Ka for the acid is applied. Ka = [products] / [reactants]
• The proportion of molecules ionised being small in most cases for weak acids the value of non-ionised acid is effectively taken as equal to the original concentration of the solution.
• Once the hydronium ion concentration is calculated, the same is substituted into the expression to find pH
• pH = $- log10 [H_{3}O^{+}]$

## Determination of ka of Weak Acids

To determine the Ka of any weak acid the steps followed are as follows:

Acid reacts with water to produce the hydronium ion $(H_{3}O^{+})$ along with a conjugate base ion. The acid dissociation and subsequent equilibrium expression is necessary. Weak acids dissociates partially as it ionises small extent in water

The acid ionisation constant also known as equilibrium constant for ionisation of weak acid and is calculated from the following basic equation.

$HA (aq) + H_{2}O(l) \leftrightarrow H_{3}O^{+} (aq) + A^{-} (aq)$

The corresponding equilibrium constant would be

Ka = [H3O+] [A-] / [HA] [H2O]

Ka = [H2O] Kc = [H3O+] [A-] / [HA]

Hence, it’s proven the Ka (acid ionisation constant) is equal to $[H_{2}O]$ Kc

## Weak Acids and Conjugate Bases

According to Bronsted Lowry theory, the conjugate acid base pairs are always intricately related to each other. A weak acid reacts with a base in forward reaction and the products are also acids and bases. Where a basic ion can accept H+ ion from water while the water acts as acid because it donates a proton in the reverse reaction. This implies that the acid involved in reaction and the base so formed are called conjugate acid base pair. The conjugate acid base pairs are always found to be differing by just one H+.

$Conjugate acid \leftrightarrow conjugate base + H^{+}$

In any conjugate acid base pair the relative strength of conjugate acid and base are mainly determined by equilibrium position in equation. In case the position of equilibrium gives out more products than reactants then the conjugate acid is considered stronger than conjugate base. We can determine the relative strengths of conjugate acids and bases in complete chemical reaction. This gives us an idea that in $\frac{H_2O}{OH^-}$ conjugate acid base pair the water plays the role of a very weak acid and OH- a very strong base. To summarise the idea of Bronsted Lowry theory for acid base conjugates let us look at the equilibrium equation once more.

$Acid + base \leftrightarrow conjugate base of acid + conjugate acid of base$
• Anions categorised as conjugate base of strong acids are considered as weak bases and have no effect on pH of solution
• Acid base behaviour of all anions given out by poly-protic acids completely depend upon the idea of deprotonation.
• Acidic cations are restricted to only metal cations with either 2+ or 3+ charges and especially ammonium ions
• All metal cations are hydrated in solvent like water and hence form ions but these ions behave as acid only when these metals have +2 or +3 charge.
Acids differ in the tendency to lose a proton. Acids which donate protons quickly and completely are considered as string acids and each of these strong acids will have weak conjugate bases as these bases will have very less tendency to take the proton back. Weak acids in turn has strong conjugate base.
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