An acid-base titration is the determination of the concentration of an acid or base by exactly neutralizing the acid/base with an acid or base of known concentration. This allows for quantitative analysis of the concentration of an unknown acid or base solution. It makes use of the neutralization reaction that occurs between acids and bases and the knowledge of how acids and bases will react if their formulas are known.
Acid-Base titrations can also be used to find percent purity of chemicals.
Before starting the titration a suitable pH indicator must be chosen. The endpoint of the reaction, the point at which all the reactants have reacted, will have a pH dependent on the relative strengths of the acid and base used. The pH of the endpoint can be estimated using the following rules:
A strong acid will react with a strong base to form a neutral (pH=7) solution.
A strong acid will react with a weak base to form an acidic (pH<7) solution.
A weak acid will react with a strong base to form a basic (pH>7) solution.
When a weak acid reacts with a weak base, the endpoint solution will be basic if the base is stronger and acidic if the acid is stronger. If both are of equal strength, then the endpoint pH will be neutral. However weak acids are not often titrated against weak bases because the color change shown with the indicator is often quick, and therefore very difficult for the observer to see the change of color.
A suitable indicator should be chosen, preferably one that will experience a change in color close to the end point of the reaction.
First, the burette should be rinsed with the standard solution, the pipette with the unknown solution, and the conical flask with distilled water.
Secondly, a known volume of the unknown concentration solution should be taken with the pipette and placed into the conical flask, along with a small amount of the indicator chosen. The burette should always be filled to the top of its scale with the known solution for ease of reading.
The known solution should then be allowed out of the burette, into the conical flask. At this stage we want a rough estimate of the amount of this solution it took to neutralize the unknown solution. Let the solution out of the burette until the indicator changes color and then record the value on the buret. This is the first titre and should be discluded from any calculations.
Perform three more titrations, this time more accurately, taking into account we know roughly where the end point will occur. Take note of each of the readings on the burette at the end point, and average these at the end. Endpoint is reached when the indicator just changes color permanently. This is best achieved by washing a hanging drop from the tip of the burette into the flask right at the end of the titration to achieve a drop that is smaller in volume than what can usually be achieved by just dripping titre off the burette.
Acid-base titration is performed with a phenolphthalein indicator, when it is a weak acid – strong base titration, a bromthymol blue indicator in strong acid- strong base reactions, and a methyl orange indicator for strong acid – weak base reactions. If the base is off the scale, i.e. a pH of >13.5, and the acid has a pH >5.5, then an Alizarin yellow indicator may be used.On the other hand, if the acid is off the scale, i.e. a pH of <0.5, and the base has a pH <8.5, then an Thymol Blue indicator may be used.
When titrating a weak acid with a strong base, pH can be calculated by the following formula: [1]
![pH = pK_s + log( \frac{[HO^-]_{added}}{[HA]_{total}-[HO^-]_{added}} )](http://upload.wikimedia.org/math/3/3/0/3309d89a1ea8c65ad166dc043d0c487f.png)
where:
- pKs is the acid dissociation constant of the weak acid.
- [HO-]added is the concentration of added strong base in the final solution (not in original standard solution)
- [HA]total is the summed concentration of both the weak acid and its conjugate base in the final solution.
Thus, at an addition of strong base that is half the amount of weak acid in the solution ([HO-]added = 0.5[HA]total), pH becomes equal to pKs.
