Home :: Wet Chemistry Analysis
Wet chemistry is a term used to refer to chemistry generally done in the liquid phase. It is also known as bench chemistry because many of the tests performed are done at a lab bench.
Traditionally, it involves the use of laboratory glassware, such as beakers and flasks, and excludes quantitative chemical analysis using instrumentation. Many high school and college laboratories teach students basic wet chemistry methods.
Before the age of theoretical and computational chemistry it was the predominant form of scientific discovery in the chemical field. This is why it is sometimes referred to as classic chemistry or classical chemistry. Because of the high volume of wet chemistry that must be done in today's society and quality control requirements, many wet chemistry methods have been automated and computerized for streamlined analysis.
Wet chemistry techniques can be used for qualitative chemical measurements, such as changes in color (colorimetry), but often involves more quantitative chemical measurements, using methods such as gravimetry and titrimetry. Some uses for wet chemistry include tests for:
It can also involve the elemental analysis of samples, e.g., water sources, for items like:
[Courtesy of http://en.wikipedia.org/wiki/Wet_chemistry]
Gravimetric analysis describes a set of methods in analytical chemistry for the quantitative determination of an analyte based on the mass of a solid. A simple example is the measurement of solids suspended in a water sample: A known volume of water is filtered, and the collected solids are weighed.
In most cases, the analyte must first be converted to a solid by precipitation with an appropriate reagent. The precipitate can then be collected by filtration, washed, dried to remove traces of moisture from the solution, and weighed. The amount of analyte in the original sample can then be calculated from the mass of the precipitate and its chemical composition.
In other cases, it may be easier to remove the analyte by vaporization. The analyte might be collected -- perhaps in a cryogenic trap or on some absorbent material such as activated carbon -- and measured directly. Or, the sample can be weighed before and after it is dried; the difference between the two masses gives the mass of analyte lost. This is especially useful in determining the water content of complex materials such as foodstuffs.
[Courtesy of http://en.wikipedia.org/wiki/Gravimetric_analysis]
Titration is a common 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 quantity 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 color). In simple acid-base titrations a pH indicator may be used, such as phenolphthalein, which turns (and stays) pink when a certain pH (about 8.2) is reached or exceeded. Methyl orange can also be used, which is red in acids and yellow in alkalis.
Not every titration requires an indicator. In some cases, either the reactants or the products are strongly colored and can serve as the "indicator". For example, an oxidation-reduction titration using potassium permanganate (pink/purple) as the titrant does not require an indicator. When the titrant is reduced, it turns colorless. After the equivalence point, there is excess titrant present. The equivalence point is identified from the first faint pink color that persists in the solution being titrated.
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 change the pH significantly — leading to an immediate color 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.
[Courtesy of http://en.wikipedia.org/wiki/Titration]