Titration is a chemical analysis process to quantify the concentration of an unknown sample using an exactly known quantity of another substance. The titration process is usually carried out by dropwise addition of a standard solution (i.e., a solution of known concentration) of titrating reagent, or titrant, from a burette, on the unknown sample placed in Erlenmeyer flask. The addition is stopped when the endpoint or the equivalence point is reached. The experimental point at which the completion of the reaction is called the endpoint or the equivalence point. The endpoint can be picked by the color change of an indicator. On the other hand, the equivalence point could be measured by potentiometric analysis where the pH change during the titration is measured.
Generally, the data obtained from the titration is used to calculate the concentration or the molarity of the analyte. The volumes of the analyte and titrant solutions, the concentration of the titrant, and pH is essential in calculating the total number of moles of analyte present. The data collected through the titration process could be illustrated on a graph called the titration graph. The shape of the curve provides important information about what is occurring in the solution during the titration. There are many types of titrations with different procedures and goals. The most common types of titration are acid-base titrations and redox titrations. An acid-base titration is classified as a strong acid-strong base, strong acid-weak base or weak acid-strong base titration. In general, acid-base titrations depend on the neutralization between an acid and a base when mixed in solution. This process could be carried out using a pH indicator or a change in pH values during the titration (potentiometric titration). In an acid-base titration, the indicator is a substance that can exist in two forms, an acid form and a basic form, with different colors. The acid-base indicator indicates the endpoint of the titration by changing color. In other terms, a pH indicator is actually a Brønsted-Lowry conjugate acid-base pair in which the acid is a different color to the base. Choosing the appropriate pH indicator depends on the strength of the acid and base in the titration. If the acid is stronger than the base, the resulting solution will be acidic with a pH < 7. On the contrary, if the acid is weaker than the base, the resulting solution will be basic with a pH > 7. However, if the strength of acid and base is equal, the formed solution will be neutral with pH equals 7. Therefore, the appropriate pH indicator will be selected as following:
- Strong acid + strong base → salt (pH=7)
Choose an indicator that changes color at pH=7
- Strong acid + weak base → salt (pH < 7)
Choose an indicator that changes color at pH < 7
- Weak acid + strong base → salt (pH > 7)
Choose an indicator that changes color at pH > 7
NH4Cl is the salt of a strong acid and a weak base, so a solution of NH4Cl will have a pH < 7 (NH4+ is a weak acid). Thus, the suitable indicator would be methyl red (pH range 4.4 - 6.0). The titration curve shown below represents the changes in pH that occur as HCl(aq) is added to NH3(aq).
The equivalence point for the reaction is represented by the blue line at pH=5.28. The background color represents the color of the solution containing the methyl red indicator over the same range of pH values. At pH < 4.4 the solution appears to be pink while at pH > 6.0 the solution appears to be yellow. At the endpoint, between pH 4.4 and 6.0, the solution appears to be orange (an equimolar mixture of pink and yellow). Since the equivalence point for the titration (pH=5.28) occurs within the pH range for the visible color change of the indicator (the endpoint between pH 4.4 and 6.0), this indicator can be used for this titration.
In a strong acid-weak base titration, there are four parts to the titration curve of a weak base (analyte) with a strong acid (titrant) as follow:
- Initial pH (pH of a weak base, alkaline pH)
- Buffer Equation (Henderson Hasselbach Equation)
- Equivalence Point (salt of weak base)
- Excess acid (pH based on the concentration of excess titrant)
In a strong acid-weak base titration, the pH will change rapidly at the beginning and then have a gradual slope until near the equivalence point. The gradual slope results from a buffer solution being produced by the addition of the strong acid, which resists rapid change in pH until the added acid or base exceeds the buffer's capacity and the rapid pH change occurs near the equivalence point.
Hydrochloric acid (strong acid) will react with ammonia (NH3) as a weak base to form aqueous ammonium chloride solution (NH4Cl) according to the following equation:
NH3(aq)+ HCl(aq) → NH4Cl(aq)
In this case, the ammonia is the Lewis base which donates its lone pair of electrons to the proton, which acts as the Lewis acid. The result is the creation of the ammonium ion that is the conjugate acid to the ammonia base. It's worth mentioning that because you're titrating a strong acid with a weak base, the pH of the resulting solution will be lower than 7 at the equivalence point. That could be attributed to the production of ammonium cation, NH4+, during the neutralization reaction which acts as a weak acid in an aqueous solution.
Initially, the pH is alkaline due to pure ammonia. Upon addition of HCl, it reacts with the ammonia forming its salt, ammonium chloride. This is an ammonia/ammonium buffer and the pH is determined by the ratio of the un-neutralized to neutralized ammonia. At the equivalence point, all the ammonia has been neutralized and converted to its conjugate acid ammonium, and so the pH is dictated by the concentration of the ammonium chloride salt. After the equivalence point, there is no more ammonia to react with the HCl and so it accumulates and the pH is decreased by the excess HCl.
The process of determining the exact concentration (molarity) of a solution is called standardization. Afterwards, the concentration of the unknown ammonia solution could be determined.
