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In medicine, titration is the process of gradually adjusting the dose of a medication until the desired effect is achieved.

Titration is a standard laboratory method of quantitative/chemical analysis which can be used to determine the concentration of a known reactant. Because volume measurements play a key role in titration, it is also known as volumetric analysis. A reagent, called the titrant, of known concentration (a standard solution) and volume is used to react with a measured volume of reactant (Analyte). Using a calibrated burette to add the titrant, it is possible to determine the exact amount that has been consumed when the endpoint is reached. The endpoint is the point at which the titration is stopped. This is classically a point at which the number of moles of titrant is equal to the number of moles of analyte, or some multiple thereof (as in di- or tri- protic acids). In the classic strong acid-strong base titration the endpoint of a titration is when the pH of the reactant is just about equal to 7, and often when the solution permanently changes color due to an indicator. There are however many different types of titrations (see below).

Many methods can be used to indicate the endpoint of a reaction; titrations often use visual indicators (the reactant mixture changes colour). In simple acid-base titrations a pH indicator may be used, such as phenolphthalein, which turns (and stays) pink when a certain pH is reached or exceeded. Due to the logarithmic nature of the pH curve, the transitions are generally extremely sharp, and thus a single drop of titrant just before the endpoint can increase the pH by several points - leading to an immediate colour change in the indicator. That said, there is a slight difference between the change in indicator color and the actual equivalence point of the titration. This error is referred to as an indicator error, and it is indeterminate.

Etymology

The word "titration" comes from the Latin "titalus," meaning inscription or title. The French word, titre, also comes from this origin, meaning rank. Titration is by definition the determination of rank or concentration of a solution, so it corresponds.

Preparing a sample for titration

Titrations are performed "in solution". If the sample isn't a liquid or solution the samples must be solved. If the analyte is very concentrated in the sample it might be useful to dilute the sample.

Although the vast majority of titrations are carried out in aqueous solution, other solvents such as glacial acetic acid or ethanol (in petrochemistry) are used for special purposes.

A measured amount of the sample can be given in the flask and then be solved or diluted. The mathematical result of the titration can be calculated directly with the measured amount. Sometimes the sample is solved or diluted beforehand and a measured amount of the solution is used for titration. In this case the solving or diluting must be done accurately with a known coefficient because the mathematical result of the titration must be multiplied with this factor.

A lot of titrations require buffering to maintain a certain pH for the reaction. Therefore buffer solutions are added to the reactant solution in the flask.

Some titrations require "masking" of a certain ion. This can be necessary when two reactants in the sample would react with the titrant and only one of them must be analysed. Or when the reaction would be disturbed or inhibited by this ion. In this case another solution is added to the sample which "masks" the unwanted ion (for instance by a weak binding with it or even forming a solid insoluble substance with it).

Some redox reactions may require heating the solution with the sample and titration while the solution is still hot (to increase the reaction rate).

Procedure

N.B. Before starting, make sure that all of your glassware—especially the burette—is clean and dry. Alternatively, you can rinse out the glassware with the solution that is to be stored in the particular piece, as this will coat the glassware with the solution as to not affect the concentration of the solution.

1. Accurately measure a volume of the reactant into to a beaker or Erlenmeyer flask.
2. Add a suitable indicator to the flask.
3. Pour the titrant into the burette, read the start-point of the liquid on the burette.
4. Turn the tap of the burette to allow the titrant to fall slowly into the reactant. Swirl the flask with the other hand or with a magnetic "flea".
5. The indicator should change colour as the titrant is added, but then quickly return to its original colour.
6. As the end-point is approached, the indicator takes longer to turn back to its starting colour. Add the titrant more slowly at this point (one drop at a time).
7. When the indicator remains at its end colour, the reaction has reached the end point. Measure the amount of titrant liquid used, as shown on the scale of the burette.
8. Repeat as many trials as needed, then average the volumes.

Once the number of moles of reactant that have been neutralised has been determined then it is easy to calculate the concentration in moles per litre.

in the above formula

C represents the concentration (Molarity)

N represents the number of moles (mol)

V represents the volume of the solution (L)

Titration curves

Titrations are often recorded on titration curves, whose compositions are generally identical: the independent variable is the volume of the titrant, while the dependent variable is the pH of the solution (which changes depending on the composition of the two solutions). The equivalence point is a significant point on the graph (the point at which all of the starting solution, usually an acid, has been neutralized by the titrant, usually a base). It can be calculated precisely by finding the second derivative of the titration curve and computing the points of inflection (where the graph changes concavity); however, in most cases, simple visual inspection of the curve will suffice (in the curve given to the right, the only immediately visible equivalence point (actually the second equivalence point) occurs after roughly 5 mL of MMH has been titrated with the citric acid solution)Therefore, to calculate the other two Pka values, one must take the seen pKa value and devide his volume into the respective amounts to find the 2 other equivalence points, and therefore half-equivalence points.

In monoprotic acids, the point halfway between the beginning of the curve (before any titrant has been added) and the equivalence point is significant: at that point, the concentrations of the two solutions (the titrant and the original solution) are equal. Therefore, the Henderson-Hasselbalch equation can be solved in this manner:

pH = pK_a + log(1) = pK_a + 0

pH = pK_a

Therefore, one can easily find the acid dissociation constant of the monoprotic acid by finding the pH of the point halfway between the beginning of the curve and the equivalence point, and solving the simplified equation. In the case of the sample curve, the K_a would be approximately 1.78 * 10 ^{− 5} from visual inspection (the actual K_a2 is 1.7 * 10 ^{− 5}).

For polyprotic acids, calculating the acid dissociation constants is only marginally more difficult: the first acid dissociation constant can be calculated the same way as it would be calculated in a monoprotic acid. The second acid dissociation constant, however, is the point halfway between the first equivalence point and the second equivalence point (and so on for acids that release more than two protons, such as phosphoric acid).

Types

Different types of titration include:

• Acid-base titration is based on the neutralization reaction between the analyte and the titrant.
  • The most common method is with a pH indicator that changes the colour of the solution permanantly at the end point.
  • A Potentiometer or a ph meter (a specific type of potentiometer) can be used, measuring the voltage in the solution. At the end point there will be a recognisable jump in the voltage. It can be more accurate than the indicator method.
  • It is also possible to perform this as conductometric titration (see below) to determine the end point.

• Redox titration is based on a redox reaction between the analyte and titrant.
  • A redox indicator can be used to make the end point visible.
  • With some redox reactions the end point is visible because the reaction itself changes the colour of the solution (reactions with Potassium permanganate or Potassium dichromate).
  • A Potentiometer can be used as well.

• Precipitation: one ion of the reactant forms with one ion of the titrant an insoluble compound. The end point is visible when the first suspended solid substance forms. This can be rather inaccurate. So, colour changing indicators are also used or the reaction is recorded via change of electrical conductivity (see below).

• Complexometric titration is based on complex formation between the analyte and titrant. The reaction can take some time. In this case it's advised to add the titrant even slower shortly before the end point. An complexometric indicator changes the colour of the solution at the end point.

• Conductometry is not a different type of titration reaction but a different method to determine the end point. It requires the recording of the change of electrical conductivity of the sample solution while the titration. While the reaction of analyte and titrant takes place, the ions in solution change. Ions of the analyte are taken out of solution forming compounds, new ions from the titrant come into the solution. Because different ions have a different conductivity the conductivity of the solution changes. At the end point there is a recognizable "break" in the change of the conductivity (for instance first it was falling then it rises or first it was slowly rising and then it is very fast rising).

• A Isothermal titration calorimeter uses the heat of reaction to determine the end point.

• Spectroscopy measures the absorption of either photons or electrons.

• Back titrations are like normal titrations, except that a known excess of a standard reagent is added to the solution being titrated. The solution is then titrated back, taking into account the addition of the excess. Back titrations are useful if the end point of the reverse titration is easier to identify than the end point of the normal titration.

• Virus titration to determine the virus concentration. It may be based on TCID50, EID50, ELD50, LD50 or pfu.

Particular uses

• As applied to biodiesel, titration is the act of determining the acidity of a sample of WVO by the dropwise addition of a known base to the sample while testing with pH paper for the desired neutral pH=7 reading. By knowing how much base neutralizes an amount of WVO, we discern how much base to add to the entire batch.

• Titrations in the petrochemical or food industry to define oils, fats or biodiesel and similar substances. An example procedure for all three can be found here: [1].
  • Acid number: an acid-base titration with colour indicator is used to determine the free fatty acid content.
  • Iodine number: a redox titration with colour indication which indicates the amount of unsaturated fatty acids.
  • Saponification value: an acid-base back titration with colour indicator or potentiometric to get a hint about the average chain length of fatty acids in a fat.

See also

• Indicator

External links

• Titration - apparatus, technique and calculation

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