# Buffers and titration

Buffers play an integral role in guaranteeing the success of a titration experiment. What are buffers, how do they function, and what is their importance in titrimetric analysis? If you are looking for all this valuable information, then this is the right place for you. Because in this article, we will discuss everything you need to know about buffers and titration, so continue reading!

## What are buffers -Definition

A buffer is defined as a chemical solution that resists changes in pH in a reaction mixture. The buffer solution usually consists of a weak organic acid or a weak base and a large concentration of its highly ionizable salt solution. An equilibrium is maintained between the weak acid/base and the dissociated ions. Small changes in pH are reversed as the equilibrium shifts its position as per Le Chatlier’s principles.

## Examples of buffers

### Acidic Buffer

Acetic acid and sodium acetate (CH3COOH + CH3COONa+).

CH3COO is the conjugate base of a weak organic acid, i.e., acetic acid or ethanoic acid (CH3COOH).

Acetic acid partially dissociates in an aqueous solution to produce acetate ions by losing protons.

Sodium acetate (CH3COONa+) completely dissociates into acetate ions. In this way, the CH3COO ions are released into the reaction mixture in a bulk concentration. A large concentration of CH3COO ions in the mixture shifts the equilibrium direction backward so that most of the acetic acid stays undissociated. This is known as the common ion effect.

### Neutral Phosphate buffer

Disodium hydrogen phosphate and sodium dihydrogen phosphate (Na2HPO4 + NaH2PO4).

## How do buffers work

• When a small concentration of an acid such as HCl is added to the mixture containing an acetate buffer, the pH of the solution decreases.
• HCl completely dissociates to yield a large concentration of H+ ions.
• A bulk of H+ ions in the reaction mixture shifts the buffer equilibrium backward as per Le Chatelier’s principles.
• CH3COO combines with H+ ions to form CH3COOH.
• Consequently, the concentration of H+ ions decreases, and the pH of the reaction mixture increases back to normal.
• Similarly, if a small concentration of a base such as NaOH is added to the reaction mixture, the pH increase as NaOH completely dissociates into Na+ and OH ions.
• The H+ ions in the solution combine with OH-, and their effect is neutralized.
• The above equilibrium shifts forward, and more H+ ions are produced to completely reverse the effect of OH ions released into the solution until the pH reverts to its initial level.
• This is known as the buffer action that maintains a fairly constant pH in the reaction mixture.
• The amount or concentration of an acid or a base required for the buffer to reach its maximum capacity of resisting pH changes is known as buffer capacity.
• In other words, the buffer capacity is defined as the number of moles of an acid or base required to raise the pH of a solution by 1 unit.
• Each conjugate acid-base pair has a definite pH range at which it works the best.
• After reaching its maximum capacity, the buffer fails to reverse any pH changes in the solution further.

## Calculation of pH of a buffer solution

The pH of a buffer solution can be calculated using the Henderson Hasselbalch equation.

As per acid-base chemistry, the dissociation of a weak organic acid can be represented as follows:

The equilibrium constant (Ka) for this reaction is:

$K_{a}&space;=&space;\frac{[H^{+}][A^{-}]}{[HA]}$

Rearranging and taking logarithms on both sides of the equation gives us:

$log&space;\frac{[HA]}{[A^{-}]}+log&space;\left&space;(&space;K_{a}&space;\right&space;)&space;=&space;log[H^{+}]$

$-log[H^{+}]&space;=&space;-log(K_{a})&space;-log\frac{[HA]}{[A^{-}]}$

$pH&space;=&space;pK_{a}&space;-log&space;\frac{[HA]}{[A^{-}]}$

Equation I

Henderson Hasselbalch equation

This equation denotes that the pH of the buffer solution is equal to the power of acid dissociation constant when [A] = [HA].

Similarly, the Henderson Hasselbalch equation for calculating the pOH of a basic buffer solution is:

Equation II

Now let us practice an example:

Question:  A buffer solution consists of 0.250 M NH3 (Kb = 1.80 x 10-5) and 0.400 M NH4Cl. Calculate the pH of this buffer solution.

Step I: Calculate the pKb of the buffered solution by taking the negative logarithm of the base dissociation constant (Kb).

pKb = – log (1.80 x 10-5) = 4.74

Step II: Apply the Henderson Hasselbalch equation (Equation II) to calculate the pOH of the basic buffer.

pOH = 4.74 + log (0.400/0.250) = 4.94

Step III: Calculate pH by applying the formula: pH + pOH = 14

pH = 14-4.94 = 9.05

Result: As it is a basic buffer solution, so its pH is above 7. This buffer solution maintains the pH of the reaction mixture close to 9.05.

Practice some more examples to calculate the pH of a buffer solution here.

Now let’s see how an ammonia buffer of pH 9.05 can be prepared in your chemistry laboratory.

## Preparation of buffer solutions

You mix the right amount of chemicals, the magic happens in a glass vessel, and the buffer is prepared. Yes, it is as easy as that. The general steps for preparing a buffer solution are:

1. Consult your chemistry laboratory’s database and calculate the accurate weight and volume that needs to be added for preparing a specific buffer solution.
2. Weigh compounds using an analytical mass balance.
3. Dissolve the compounds in a suitable solvent.
4. Check and adjust the required pH using a digital pH meter.
5. Transfer the solution into a volumetric flask.
6. Raise the volume of the solution up to the mark using the same solvent.
7. Transfer the solution to an airtight reagent bottle.
8. Label the buffer solution and document your results.

SOP for preparing an ammonia buffer of pH 9.5:

• 33.5 g ammonium chloride (NH4Cl) is dissolved completely in double distilled water.
• 42 mL of 10 M ammonia solution (NH4OH) is added to the above solution.
• This mixture is transferred in a 250 mL volumetric glass flask.
• The final volume is raised up to the mark in the glass flask using distilled water.

## What is a buffer region in titrimetry

In titrations, the buffer region is the zone where the pH of the titration mixture stays fairly constant. The buffer region can be determined from a titration curve. For instance, in the titration curve shown below, there is a flat region in the middle. This is the buffer region.

The pH of the titration mixture stays constant in this region. Also, the equivalence point of the particular titration experiment lies within the buffer region.

After crossing the buffer zone, the pH changes dramatically with just one drop of the titrant added from the burette.

The buffer region strongly influences the shape of a titration curve in acid-base titrations. Buffer solutions are also important to maintain a stable pH in complexometric titrations.

Acid-base indicators exhibit color changes due to their structural transformations under different pH conditions.

Besides their immense importance in titrimetry, buffer solutions are also important for resisting abrupt pH changes in biological fluids such as the blood.

There are three important buffer systems in the human body, namely:

• Carbonate/Bicarbonate buffer
• Phosphate buffer
• Plasma protein buffer