It is considered that Water is abundant and common We waste it, pollute it, let it run down the drain, flush it away………..we take it for granted. However chemically speaking, water is really not common at all When compared to other compounds of similar size, composition and structure- it is absolutely unique In fact its properties are so unusual that it would be irreplaceable. Let’s take a chemical look at these unusual properties how they arise and what their implications are.

Dissociation of Water

Dissociation of Water Water can both dissociate in to ions and can decompose in to Oxygen and Hydrogen Decomposition is a rare phenomenon in water.

But dissociation (H+ and OH- ion) is a normal process in chemical reactions. This is also called as ionization.

But dissociation (H+ and OH- ion) is a normal process in chemical reactions. This is also called as ionization.

Limits of Dissociation In pure state, water does get ionize To a very small Extent. While most of the water remains in water as un ionized H2O H+ + OH- ------I According to law of mass action Keq = [H+] [OH]----------------II [H2O]

It means, K eq is equal to the mathematical products of molar concentration written in [ ] of dissociated divided by [ undissociated ] amount of water. [H2O]. As the quantity of dissociated water molecules is very small so the quantity of intact water molecules can be taken as constant. Rearrange the equation II: Keq [H2O] = [H+] [OH-] [H+] [OH-] = Keq [H2O]---------------III From equation III, the [H+] can be calculated if K eq and quantity of water [H2O], in pure water is known, At 25oC, K =1.8X10-16 moles/L and [H2O]= Wt, of 1 litre water =1000 = 55.5 moles/ L molecular wt of water 18

Put the value of Keq and [H2O] in equation III [H+][OH] = 1.8x10-16 [55.5] =1x10-14 moles /L----------IV Since water is neutral because [H+] = [OH] ---------------------V Put [H+], in place of [OH] in equation IV. [H+][H+] = 1x10-14 moles/L or [H+]2 =1x10-14 moles /L or [H+]2 =[10-7]-2 or [H+] = [10-7] moles/L------------------VI

pH scale In 1909, Sorensen described the hydrogen ion concentration, [H+], in terms of pH in order to convert fraction into whole numbers. The symbole “P” stands for ions concentration [H+] . Thus pH = - log [H+]---------------VII Put the value of [H+] form equation VI into equation VII. pH =-log[10-7] =- [-7] =[7]

For pure water, [H+]=[OH-] = 1.0x10-7 mol/L so the pH =7 ---------------------VIII Similar to [H+] which is stated as pH, the [OH-] for water can be expressed as pOH. [H+] [OH-] = 1x10-14 moles/L

Take negative log of both sides of the above eq. pH +pOH =14--------------IX As pH =7 so pOH = 14-7 or pOH = 14-7 =7 -------------------X Hence pOH, the negative log of [OH] for pure water is also 7. A solution with highest [H+] will have lowest pH value and lowest [OH] .

Dissociation in acids and bases What are strong and weak acids / bases ? Completely ionized  Strong acids / bases HCl  H+ + Cl- NaoH  Na+ + OH- Also called as strong electrolytes

Weak acids or bases? Species which ionize only partially CH3COOH  CH3COO- + H+ Ca(OH)2  Ca++ + 2(OH-)

Many biochemical possess functional groups, which are weak acids or weak bases e.g. Carboxyl group – COOH H Amino group R – N These are present in all proteins, nucleic acids, coenzymes and intermediary metabolites. Hence their dissociation is fundamental to changes in intracellular pH, structure and biochemical activity of these compounds.

Biological system is quite influenced by [H+] pH of blood is maintained within a normal range 7.35 – 7.45. Fluctuation beyond these limit leads to acidosis or alkalosis Enzymes control many biological reactions and each enzyme function optimally at a particular pH. Oxygenated Hb in RBC is deoxygenated and reaction is influenced by Hydrogen ions. This reaction makes O2 available to tissues for use.

The equation which governs the dissociation process of weak acids and bases is known as Henderson – Hasselbalch Equation (HHE). Let us study the dissociation of a weak acid HA. HA H+ + A- (i) Dissociation constant for reaction (i) Ka = [H]+ [A]- [HA] or [H+] = Ka[HA] [A-] changing to log form

Contd. log [H+] = log Ka + log [HA] [A-] - log [H+] = - log Ka – log [HA] pH = pKa – log [HA] pH = pKa – log [Acid] [Base] or pH = pKa + log [Base] [Acid] This equation is known as Henderson – Hasselbalch Equation

Contd. Salt being proton acceptor can also be termed as a base and then pH = pKa + log [Salt] [Acid] This equation helps in the calculation of pH of a known solution and for determining the amounts of an acid and its salt to prepare a solution at a given pH.

Contd. If base / acid is 1:1 ratio then Henderson – Hasselbalch equation becomes pH = pKa + log 1/1 log 1/1 = 0 pH = pKa

Interpretation pH of an equi-molar solution (50:50) of a conjugate acid base pair is equal to its pKa And such solution will have maximum buffering capacity.

Let us study the dissociation of a base BOH BOH B+ + OH- The dissociation constant pKb and equilibrium constant for this reaction are similar to that of an acid.

Contd. like pH, of an acid pOH of a weak base can also be calculated as pOH = pKb + log [B+] [BOH] pOH = pKb

pH = pKa + log [Base] [Acid] pH = pKa + log [Salt]

Value of pH is kept in Henderson – Value of pH is kept in Henderson – Hasselbalch equation and pKa calculated This equation in its simplest form hold well for molecules having single titrable groups e.g. CH3COOH, HCOOH Molecules having more than one titrable group like citric acid, more complex mathematical calculations are required.

Relative strength Relative strengths of weak acids and bases are expressed quantitatively as their dissociation constant, which express their tendency to ionize.

Dissociation Constant and pK values of Some representative Carboxylic acids Acetic (Mono) 1st 1.76x10-5 4.75 Glutaric (Di) 1st 4.58x10-5 2nd 3.89x10-5 4.34 5.41 Citric (Tri) 1st 8.4x10-4 2nd 1.8x10-5 3rd 4.0x10-6 3.08 4.74 5.40

Determining pK value of an acid pKa of any dissociable group is characteristic property of a molecule. It is easy to determine by preparing titration curve. pH meter, burette and titrant are needed Known amounts of acids / alkali is added to a solution of unknown and pH is determined.

The Acid Dissociation Constant, Ka HA(aq) + H2O(l) H3O+(aq) + A-(aq) [H3O+][A-] [HA][H2O] Kc = Kc[H2O] = Ka = [H3O+][A-] [HA] The value of Ka is an indication of acid strength. Stronger acid larger Ka higher [H3O+] Weaker acid smaller Ka lower % dissociation of HA

The relations among [H3O+], pH, [OH-], and pOH.

General form of a titration curve calculated from the H.H. EQUATION.

Dissociation Curve of glycine

Some Common Acids and Bases and their Household Uses.

Methods for measuring the pH of an aqueous solution. pH paper pH meter