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| {{distinguish2|the mathematical notion of [[tetration]]}}
| | <br><br>My web-site; [http://volvo780.hu/modules.php?name=Your_Account&op=userinfo&username=EEVFT forum ecigarette] |
| {{about|''volumetric titration''||Titration (disambiguation)}}
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| [[File:Winkler Titration Prior Titration.jpg|thumb|A [[Winkler test for dissolved oxygen|Winkler titration]] to determine the concentration of dissolved oxygen in a water sample]]
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| '''Titration''', also known as '''titrimetry''',<ref>
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| {{Cite encyclopedia
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| | title = titrimetry
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| | encyclopedia = The Century Dictionary and Cyclopedia
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| | last = Whitney
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| | first = W.D.
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| | coauthors = B.E. Smith
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| | pages = 6504
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| | publisher = The Century co.
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| | year = 1911
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| }}</ref> is a common laboratory method of [[Quantitative research|quantitative]] [[Analytical chemistry|chemical analysis]] that is used to determine the unknown [[concentration]] of an identified [[analyte]]. Since [[volume]] measurements play a key role in titration, it is also known as '''volumetric analysis'''. A [[reagent]], called the ''titrant'' or ''titrator''<ref>
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| {{Cite book
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| | title = Compendium for Basal Practice in Biochemistry
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| | publisher = Aarhus University
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| | year = 2008
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| }}</ref> is prepared as a [[standard solution]]. A known concentration and volume of titrant reacts with a solution of ''analyte'' or ''titrand''<ref>
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| {{Cite web
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| | title = titrand
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| | work = Science & Technology Dictionary
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| | publisher = McGraw-Hill
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| | url = http://www.answers.com/topic/titrand
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| | accessdate = 30 September 2011
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| }}</ref> to determine concentration.
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| | |
| ==History and etymology==
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| The word "titration" comes from the Latin word ''titulus'', meaning inscription or title. The French word ''titre'', also from this origin, means rank.
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| Volumetric analysis originated in late 18th-century France. [[François-Antoine-Henri Descroizilles]] ([[:fr:François-Antoine-Henri Descroizilles|fr]]) developed the first burette (which was similar to a graduated cylinder) in 1791.<ref>
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| {{cite book
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| |last=Szabadváry |first=F.
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| |year=1993
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| |title=History of Analytical Chemistry
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| |pages=208–209
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| |publisher=[[Taylor & Francis]]
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| |isbn=2-88124-569-2
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| }}</ref> [[Joseph Louis Gay-Lussac]] developed an improved version of the burette that included a side arm, and coined the terms "[[pipette]]" and "[[burette]]" in an 1824 paper on the standardization of indigo solutions. A major breakthrough in the methodology and popularization of volumetric analysis was due to [[Karl Friedrich Mohr]], who redesigned the burette by placing a clamp and a tip at the bottom, and wrote the first textbook on the topic, ''Lehrbuch der chemisch-analytischen Titrirmethode'' (''Textbook of analytical-chemical titration methods''), published in 1855.<ref>
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| {{cite book
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| |last=Rosenfeld |first=L.
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| |year=1999
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| |title=Four Centuries of Clinical Chemistry
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| |pages=72–75
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| |publisher=[[CRC Press]]
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| |isbn=90-5699-645-2
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| }}</ref>
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| | |
| ==Procedure==
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| [[File:Titration for Soil.JPG|thumb|Analysis of soil samples by titration]]
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| A typical titration begins with a [[beaker (glassware)|beaker]] or [[Erlenmeyer flask]] containing a very precise volume of the analyte and a small amount of indicator placed underneath a calibrated [[burette]] or [[Pipette|chemistry pipetting syringe]] containing the titrant. Small volumes of the titrant are then added to the analyte and indicator until the indicator changes, reflecting arrival at the endpoint of the titration. Depending on the endpoint desired, single drops or less than a single drop of the titrant can make the difference between a permanent and temporary change in the indicator. When the [[equivalence point|endpoint]] of the reaction is reached, the volume of reactant consumed is measured and used to calculate the concentration of analyte by
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| :<math>\mathbf{C}_a=\frac{\mathbf{C}_t\times\mathbf{V}_t\times\mathbf{M}}{\mathbf{V}_a}</math>
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| | |
| where '''C'''<sub>a</sub> is the concentration of the analyte, typically in [[molarity]]; '''C'''<sub>t</sub> is the concentration of the titrant, typically in molarity; '''V'''<sub>t</sub> is the volume of the titrant used, typically in liters; '''M''' is the mole ratio of the analyte and reactant from the balanced [[chemical equation]]; and '''V'''<sub>a</sub> is the volume of the analyte used, typically in liters.<ref>
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| {{Cite book
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| | last = Harris
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| | first = D.C.
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| | title = Quantitative Chemical Analysis
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| | publisher = W. H. Freeman and Company
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| | edition = 7
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| | year = 2007
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| | pages = 12
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| | isbn = 9780716770411
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| }}</ref>
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| | |
| ===Preparation techniques===
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| Typical titrations require titrant and analyte to be in a liquid (solution) form. Though solids are usually dissolved into an aqueous solution, other solvents such as [[Acetic acid|glacial acetic acid]] or [[ethanol]] are used for special purposes (as in [[petrochemistry]]).<ref>
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| {{Cite book
| |
| | last = Matar
| |
| | first = S.
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| | coauthors = L.F. Hatch
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| | title = Chemistry of Petrochemical Processes
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| | publisher = Gulf Professional Publishing
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| | edition = 2
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| | year = 2001
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| | isbn = 0-88415-315-0
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| }}</ref> Concentrated analytes are often diluted to improve accuracy.
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| | |
| Many non-acid-base titrations require a constant [[pH]] throughout the reaction. Therefore a [[buffer solution]] may be added to the titration chamber to maintain the pH.<ref>
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| {{Cite book
| |
| | last = Verma
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| | first = Dr. N.K.
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| | coauthors = S.K. Khanna, Dr. B. Kapila
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| | title = Comprehensive Chemistry XI
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| | publisher = Laxmi Publications
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| | location = New Dehli
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| | pages = 642–645
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| | isbn = 81-7008-596-9
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| }}</ref>
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| | |
| In instances where two reactants in a sample may react with the titrant and only one is the desired analyte, a separate [[Masking agent|masking solution]] may be added to the reaction chamber which masks the unwanted ion.<ref>
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| {{Cite book
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| | last = Patnaik
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| | first = P.
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| | title = Dean's Analytical Chemistry Handbook
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| | publisher = McGraw-Hill Prof Med/Tech
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| | edition = 2
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| | year = 2004
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| | pages = 2.11–2.16
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| | isbn = 0-07-141060-0
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| }}</ref>
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| | |
| Some [[redox]] reactions may require heating the sample solution and titrating while the solution is still hot to increase the [[reaction rate]]. For instance, the oxidation of some oxalate solutions requires heating to {{convert|60|C|F}} to maintain a reasonable rate of reaction.<ref>
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| {{Cite book
| |
| | last = Walther
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| | first = J.V.
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| | title = Essentials of Geochemistry
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| | publisher = Jones & Bartlett Learning
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| | year = 2005
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| | pages = 515–520
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| | isbn = 0-7637-2642-7
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| }}</ref>
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| | |
| ==Titration curves==
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| {{main|Titration curve}}
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| | |
| [[Image:Oxalic acid titration grid.svg|thumb|A typical titration curve of a [[Diprotic acid|diprotic]] acid k titrated with a strong base. Shown here is [[oxalic acid]] titrated with [[sodium hydroxide]]. Each of the two equivalence points is visible.]]
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| A titration curve is a curve in the plane whose ''x''-coordinate is the volume of [[titrant]] added since the beginning of the titration, and whose ''y''-coordinate is the concentration of the analyte at the corresponding stage of the titration (in an acid-base titration, the ''y''-coordinate is usually the pH of the solution).<ref>
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| {{Cite book
| |
| | last = Reger
| |
| | first = D.L.
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| | coauthors = S.R. Goode, D.W. Ball
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| | title = Chemistry: Principles and Practice
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| | publisher = Cengage Learning
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| | edition = 3
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| | year = 2009
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| | pages = 684–693
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| | isbn = 0-534-42012-5
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| }}</ref>
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| | |
| In an acid-base titration, the titration curve reflects the strength of the corresponding acid and base. For a strong acid and a strong base, the curve will be relatively smooth and very steep near the equivalence point. Because of this, a small change in titrant volume near the equivalence point results in a large pH change and many indicators would be appropriate (for instance [[litmus test|litmus]], [[phenolphthalein]] or [[bromothymol blue]]).
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| | |
| If one reagent is a weak acid or base and the other is a strong acid or base, the titration curve is irregular and the pH shifts less with small additions of titrant near the equivalence point. For example, the titration curve for the titration between [[oxalic acid]] (a weak acid) and [[sodium hydroxide]] (a strong base) is pictured. The equivalence point occurs between pH 8-10, indicating the solution is basic at the equivalence point and an indicator such as [[phenolphthalein]] would be appropriate. Titration curves corresponding to weak bases and strong acids are similarly behaved, with the solution being acidic at the equivalence point and indicators such as [[methyl orange]] and [[bromothymol blue]] being most appropriate.
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| | |
| Titrations between a weak acid and a weak base have titration curves which are highly irregular. Because of this, no definite indicator may be appropriate and a [[pH meter]] is often used to monitor the reaction.<ref>
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| {{Cite book
| |
| | last = Bewick
| |
| | first = S.
| |
| | coauthors = J. Edge, T. Forsythe, and R. Parsons
| |
| | title = CK12 Chemistry
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| | publisher = CK-12 Foundation
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| | year = 2009
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| | pages = 794–797
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| | isbn =
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| }}</ref>
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| | |
| The type of function that can be used to describe the curve is called a [[sigmoid function]].
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| | |
| ==Types of titrations==
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| There are many types of titrations with different procedures and goals. The most common types of qualitative titration are [[acid-base titration]]s and [[redox titration]]s.
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| | |
| ===Acid–base titration===
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| {{Main|Acid–base titration}}
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| | |
| {|border="2" cellpadding="5" align="center" style="text-align: center;" class=wikitable
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| |-
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| !style="background:#efefef;"|Indicator
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| !style="background:#efefef;"|Color on acidic side
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| !style="background:#efefef;"|Range of color change
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| !style="background:#efefef;"|Color on basic side
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| |-
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| !style="background:#efefef;"|Methyl violet
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| | Yellow || 0.0–1.6 || Violet
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| |-
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| !style="background:#efefef;"|Bromophenol blue
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| | Yellow || 3.0–4.6 || Blue
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| |-
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| !style="background:#efefef;"|Methyl orange
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| | Red || 3.1–4.4 || Yellow
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| |-
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| !style="background:#efefef;"|Methyl red
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| | Red || 4.4–6.3 || Yellow
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| |-
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| !style="background:#efefef;"|Litmus
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| | Red || 5.0–8.0 || Blue
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| |-
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| !style="background:#efefef;"|Bromothymol blue
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| | Yellow || 6.0–7.6 || Blue
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| |-
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| !style="background:#efefef;"|Phenolphthalein
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| | Colorless || 8.3–10.0 || Pink
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| |-
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| !style="background:#efefef;"|Alizarin yellow
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| | Yellow || 10.1–12.0 || Red
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| |}
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| | |
| Acid-base titrations depend on the [[Neutralization (chemistry)|neutralization]] between an acid and a base when mixed in solution. In addition to the sample, an appropriate [[pH indicator|indicator]] is added to the titration chamber, reflecting the pH range of the equivalence point. The acid-base indicator indicates the endpoint of the titration by changing color. The endpoint and the equivalence point are not exactly the same because the equivalence point is determined by the stoichiometry of the reaction while the endpoint is just the color change from the indicator. Thus, a careful selection of the indicator will reduce the indicator error. For example, if the equivalence point is at a pH of 8.4, then the Phenolphthalein indicator would be used instead of Alizarin Yellow because phenolphthalein would reduce the indicator error. Common indicators, their colors, and the pH range in which they change color are given in the table above.<ref>
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| {{Cite web
| |
| | title = pH measurements with indicators
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| | url = http://www.ph-meter.info/pH-measurements-indicators
| |
| | accessdate = 29 September 2011
| |
| }}</ref> When more precise results are required, or when the reagents are a weak acid and a weak base, a [[pH meter]] or a conductance meter are used.
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| | |
| ===Redox titration===
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| {{Main|Redox titration}}
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| | |
| Redox titrations are based on a [[redox|reduction-oxidation reaction]] between an oxidizing agent and a reducing agent. A [[potentiometer]] or a [[redox indicator]] is usually used to determine the endpoint of the titration, as when one of the constituents is the oxidizing agent [[potassium dichromate]]. The color change of the solution from orange to green is not definite, therefore an indicator such as sodium diphenylamine is used.<ref>
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| {{Cite book
| |
| | last = Vogel
| |
| | first = A.I.
| |
| | coauthors = J. Mendham
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| | title = Vogel's textbook of quantitative chemical analysis
| |
| | publisher = Prentice Hall
| |
| | edition = 6
| |
| | year = 2000
| |
| | pages = 423
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| | isbn = 0-582-22628-7
| |
| }}</ref> Analysis of wines for [[sulfur dioxide]] requires iodine as an oxidizing agent. In this case, starch is used as an indicator; a blue starch-iodine complex is formed in the presence of excess iodine, signalling the endpoint.<ref>
| |
| {{Cite book
| |
| | last = Amerine
| |
| | first = M.A.
| |
| | coauthors = M.A. Joslyn
| |
| | title = Table wines: the technology of their production
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| | publisher = University of California Press
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| | volume = 2
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| | edition = 2
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| | year = 1970
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| | pages = 751–753
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| | isbn = 0-520-01657-2
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| }}</ref>
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| | |
| Some redox titrations do not require an indicator, due to the intense color of the constituents. For instance, in [[permanganometry]] a slight persisting pink color signals the endpoint of the titration because of the color of the excess oxidizing agent [[potassium permanganate]].<ref>
| |
| {{Cite book
| |
| | last = German Chemical Society. Division of Analytical Chemistry
| |
| | title = Fresenius' Journal of Analytical Chemistry
| |
| | publisher = J.F. Bergmann
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| | volume = 166-167
| |
| | year = 1959
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| | location = University of Michigan
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| | pages = 1
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| | language = German
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| }}</ref>
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| | |
| ===Gas phase titration===
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| Gas phase titrations are titrations done in the [[gas phase]], specifically as methods for determining reactive species by reaction with an excess of some other gas, acting as the titrant. In one common the gas phase titration, gaseous [[ozone]] is titrated with nitrogen oxide according to the reaction
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| | |
| :O<sub>3</sub> + NO → O<sub>2</sub> + NO<sub>2</sub>.<ref>
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| {{Cite book
| |
| | last = Hänsch
| |
| | first = T.W.
| |
| | title = Metrology and Fundamental Constants
| |
| | publisher = IOS Press
| |
| | year = 2007
| |
| | pages = 568
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| | isbn = 1-58603-784-6
| |
| }}</ref><ref>
| |
| {{Cite web
| |
| | title = Gas phase titration
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| | publisher = Bureau International des Poids et Mesures
| |
| | url = http://www.bipm.fr/en/scientific/chem/gas_titration.html
| |
| | accessdate = 29 September 2001
| |
| }}</ref>
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| | |
| After the reaction is complete, the remaining titrant and product are quantified (e.g., by [[Fourier transform spectroscopy|FT-IR]]); this is used to determine the amount of analyte in the original sample.
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| | |
| Gas phase titration has several advantages over simple [[spectrophotometry]]. First, the measurement does not depend on path length, because the same path length is used for the measurement of both the excess titrant and the product. Second, the measurement does not depend on a linear change in absorbance as a function of analyte concentration as defined by the [[Beer-Lambert law]]. Third, it is useful for samples containing species which interfere at wavelengths typically used for the analyte.<ref>
| |
| {{Cite journal
| |
| | last = DeMore
| |
| | first = W.B.
| |
| | coauthors = M. Patapoff
| |
| | title = Comparison of Ozone Determinations by Ultraviolet Photometry
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| and Gas-Phase Titration
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| | journal = Environmental Science & Technology
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| | volume = 10
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| | issue = 9
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| | pages = 897–899
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| | date = September 1976
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| | doi = 10.1021/es60120a012
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| }}</ref>
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| | |
| ===Complexometric titration===
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| {{Main|Complexometric titration}}
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| | |
| Complexometric titrations rely on the formation of a [[complex (chemistry)|complex]] between the analyte and the titrant. In general, they require specialized [[complexometric indicator|indicator]]s that form weak complexes with the analyte. Common examples are [[Eriochrome Black T]] for the titration of [[calcium]] and [[magnesium]] ions, and the [[chelating agent]] [[EDTA]] used to titrate metal ions in solution.<ref>
| |
| {{Cite book
| |
| | last = Khopkar
| |
| | first = S.M.
| |
| | title = Basic Concepts of Analytical Chemistry
| |
| | publisher = New Age International
| |
| | edition = 2
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| | year = 1998
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| | pages = 63–76
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| | isbn = 81-224-1159-2
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| }}</ref>
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| | |
| ===Zeta potential titration===
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| {{Main|Zeta potential titration}}
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| | |
| Zeta potential titrations are titrations in which the completion is monitored by the [[zeta potential]], rather than by an [[pH indicator|indicator]], in order to characterize [[heterogeneous]] systems, such as [[colloid]]s.<ref>
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| {{Cite journal
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| | last = Somasundaran
| |
| | first = P.
| |
| | title = Calculation of Zeta-Potentials from Electrokinetic Data
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| | journal = Encyclopedia of Surface and Colloid Science
| |
| | volume = 2
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| | edition = 2
| |
| | pages = 1097
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| | publisher = CRC Press
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| | year = 2006
| |
| | isbn = 0-8493-9607-7
| |
| }}</ref> One of the uses is to determine the [[iso-electric point]] when [[surface charge]] becomes zero, achieved by changing the [[pH]] or adding [[surfactant]]. Another use is to determine the optimum dose for [[flocculation]] or [[stabilizer (chemistry)|stabilization]].<ref>
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| {{Cite book
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| | last = Dukhin
| |
| | first = A.S.
| |
| | coauthors = P.J. Goetz
| |
| | title = Ultrasound for Characterizing Colloids: Particle sizing, Zeta potential, Rheology
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| | publisher = Elsevier
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| | series = Studies in Interface Science
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| | volume = 15
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| | year = 2002
| |
| | pages = 256–263
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| | isbn = 0-444-51164-4
| |
| }}</ref>
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| | |
| ===Assay===
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| {{Main|Assay}}
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| | |
| An assay is a form of biological titration used to determine the concentration of a [[virus]] or [[bacterium]]. Serial dilutions are performed on a sample in a fixed ratio (such as 1:1, 1:2, 1:4, 1:8, etc.) until the last dilution does not give a positive test for the presence of the virus. This value is known as the [[titer]], and is most commonly determined through [[ELISA|enzyme-linked immunosorbent assay]] (ELISA).<ref>
| |
| {{Cite book
| |
| | last = Decker
| |
| | first = J.M.
| |
| | title = Introduction to immunology
| |
| | publisher = Wiley-Blackwell
| |
| | series = Eleventh Hour
| |
| | edition = 3
| |
| | year = 2000
| |
| | pages = 18–20
| |
| | isbn = 0-632-04415-2
| |
| }}</ref>
| |
| | |
| ==Measuring the endpoint of a titration==
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| {{main|Equivalence point}}
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| Different methods to determine the endpoint include:<ref>
| |
| {{Cite encyclopedia
| |
| | title = Titration
| |
| | encyclopedia = Science & Technology Encyclopedia
| |
| | publisher = McGraw-Hill
| |
| | url = http://www.answers.com/topic/titration
| |
| | accessdate = 30 September 2011
| |
| }}</ref>
| |
| *Indicator: A substance that changes color in response to a chemical change. An [[pH indicator|acid-base indicator]] (e.g., [[phenolphthalein]]) changes color depending on the pH. [[Redox indicator]]s are also used. A drop of indicator solution is added to the titration at the beginning; the endpoint has been reached when the color changes.
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| *[[Potentiometer (measuring instrument)|Potentiometer]]: An instrument that measures the [[electrode potential]] of the solution. These are used for redox titrations; the potential of the working electrode will suddenly change as the endpoint is reached.
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| [[File:PHmeter basic.JPG|thumb|An elementary [[pH meter]] that can be used to monitor titration reactions]]
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| *[[pH meter]]: A potentiometer with an electrode whose potential depends on the amount of H<sup>+</sup> ion present in the solution. (This is an example of an [[ion-selective electrode]].) The pH of the solution is measured throughout the titration, more accurately than with an indicator; at the endpoint there will be a sudden change in the measured pH.
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| *[[electrical conductivity|Conductivity]]: A measurement of ions in a solution. Ion concentration can change significantly in a titration, which changes the conductivity. (For instance, during an acid-base titration, the H<sup>+</sup> and OH<sup>-</sup> ions react to form neutral H<sub>2</sub>O.) As total conductance depends on all ions present in the solution and not all ions contribute equally (due to [[electrophoretic mobility|mobility]] and [[ionic strength]]), predicting the change in conductivity is more difficult than measuring it.
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| *Color change: In some reactions, the solution changes color without any added indicator. This is often seen in redox titrations when the different oxidation states of the product and reactant produce different colors.
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| *[[Precipitation (chemistry)|Precipitation]]: If a reaction produces a solid, a precipitate will form during the titration. A classic example is the reaction between Ag<sup>+</sup> and Cl<sup>-</sup> to form the insoluble salt AgCl. Cloudy precipitates usually make it difficult to determine the endpoint precisely. To compensate, precipitation titrations often have to be done as "back" titrations (see below).
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| *[[Calorimeter#Isothermal titration calorimeter|Isothermal titration calorimeter]]: An instrument that measures the heat produced or consumed by the reaction to determine the endpoint. Used in [[biochemistry|biochemical]] titrations, such as the determination of how [[substrate (biochemistry)|substrate]]s bind to [[enzyme]]s.
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| *[[Thermometric Titration|Thermometric titrimetry]]: Differentiated from calorimetric titrimetry because the heat of the reaction (as indicated by temperature rise or fall) is not used to determine the amount of analyte in the sample solution. Instead, the endpoint is determined by ''the rate of temperature change''.
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| *[[Spectroscopy]]: Used to measure the absorption of light by the solution during titration if the [[spectrum]] of the reactant, titrant or product is known. The concentration of the material can be determined by [[Beer's Law]].
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| *[[amperometric titration|Amperometry]]: Measures the current produced by the titration reaction as a result of the oxidation or reduction of the analyte. The endpoint is detected as a change in the current. This method is most useful when the excess titrant can be reduced, as in the titration of [[halide]]s with Ag<sup>+</sup>.
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| | |
| ===Endpoint and equivalence point===
| |
| Though equivalence point and endpoint are used interchangeably, they are different terms. ''Equivalence point'' is the theoretical completion of the reaction: the volume of added titrant at which the number of [[Mole (unit)|moles]] of titrant is equal to the number of moles of analyte, or some multiple thereof (as in [[polyprotic]] acids). ''Endpoint'' is what is actually measured, a physical change in the solution as determined by an [[pH indicator|indicator]] or an instrument mentioned above.<ref>
| |
| {{Cite book
| |
| | last = Harris
| |
| | first = D.C.
| |
| | title = Quantitative Chemical Analysis
| |
| | publisher = Macmillan
| |
| | edition = 6
| |
| | year = 2003
| |
| | pages = 129
| |
| | isbn = 0-7167-4464-3
| |
| }}</ref>
| |
| | |
| There is a slight difference between the endpoint and the equivalence point of the titration. This error is referred to as an indicator error, and it is indeterminate.<ref>
| |
| {{Cite book
| |
| | last = Hannan
| |
| | first = H.J.
| |
| | title = Technician's Formulation Handbook for Industrial and Household Cleaning Products
| |
| | publisher = Lulu.com
| |
| | year = 2007
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| | pages = 103
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| | isbn = 0-615-15601-0
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| }}</ref>
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| | |
| ===Back titration===
| |
| Back titration is a titration done in reverse; instead of titrating the original sample, a known excess of standard reagent is added to the solution, and the excess is titrated. A back titration is useful if the endpoint of the reverse titration is easier to identify than the endpoint of the normal titration, as with [[Precipitation (chemistry)|precipitation]] reactions. Back titrations are also useful if the reaction between the analyte and the titrant is very slow, or when the analyte is in a non-[[Solubility|soluble]] solid.<ref>
| |
| {{Cite book
| |
| | last = Kenkel
| |
| | first = J.
| |
| | title = Analytical Chemistry for Technicians
| |
| | publisher = CRC Press
| |
| | volume = 1
| |
| | edition = 3
| |
| | year = 2003
| |
| | pages = 108–109
| |
| | isbn =
| |
| }}</ref>
| |
| | |
| ==Particular uses==
| |
| Specific examples of titrations include:
| |
| | |
| ;Acid-Base Titrations
| |
| | |
| *In [[biodiesel]]: Waste vegetable oil ([[Vegetable oil fuel|WVO]]) must be neutralized before a batch may be processed. A portion of WVO is titrated with a base to determine acidity, so the rest of the batch may be properly neutralized. This removes [[Fatty acid#Free fatty acids|free fatty acids]] from the WVO that would normally react to make soap instead of biodiesel.<ref>
| |
| {{Cite book
| |
| | last = Purcella
| |
| | first = G.
| |
| | title = Do It Yourself Guide to Biodiesel: Your Alternative Fuel Solution for Saving Money, Reducing Oil Dependency, Helping the Planet
| |
| | publisher = Ulysses Press
| |
| | year = 2007
| |
| | pages = 81–96
| |
| | isbn = 1-56975-624-4
| |
| }}</ref>
| |
| *[[Kjeldahl method]]: A measure of nitrogen content in a sample. Organic nitrogen is digested into [[ammonia]] with [[sulfuric acid]] and [[potassium sulfate]]. Finally, ammonia is back titrated with [[boric acid]] and then [[sodium carbonate]].<ref>
| |
| {{Cite book
| |
| | title = Remington: the science and practice of pharmacy
| |
| | publisher = Lippincott Williams & Wilkins
| |
| | volume = 1
| |
| | edition = 21
| |
| | year = 2005
| |
| | pages = 501
| |
| | isbn = 0-7817-4673-6
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| }}</ref>
| |
| *[[Acid value]]: The mass in milligrams of [[potassium hydroxide]] (KOH) required to neutralize [[carboxylic acid]] in one gram of sample. An example is the determination of [[Fatty acid#Free fatty acids|free fatty acid]] content. These titrations are achieved at low temperatures.
| |
| *[[Saponification value]]: The mass in milligrams of KOH required to saponify carboxylic acid in one gram of sample. Saponification is used to determine average chain length of fatty acids in fat. These titrations are achieved at high temperatures.
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| *Ester value (or ester index): A calculated index. Ester value = Saponification value – Acid value.
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| *Amine value: The mass in milligrams of KOH equal to the [[amine]] content in one gram of sample.
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| *[[Hydroxyl value]]: The mass in milligrams of KOH required to neutralize [[hydroxyl]] groups in one gram of sample. The analyte is [[acetylation|acetylated]] using [[acetic anhydride]] then titrated with KOH.
| |
| | |
| ;Redox titrations
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| | |
| *[[Winkler test for dissolved oxygen]]: Used to determine oxygen concentration in water. Oxygen in water samples is reduced using [[manganese(II) sulfate]], which reacts with [[potassium iodide]] to produce [[iodine]]. The iodine is released in proportion to the oxygen in the sample, thus the oxygen concentration is determined with a redox titration of iodine with [[thiosulfate]] using a starch indicator.<ref>
| |
| {{Cite book
| |
| | last = Spellman
| |
| | first = F.R.
| |
| | title = Handbook of Water and Wastewater Treatment Plant Operations
| |
| | publisher = CRC Press
| |
| | edition = 2
| |
| | year = 2009
| |
| | pages = 545
| |
| | isbn = 1-4200-7530-6
| |
| }}</ref>
| |
| *[[Vitamin C]]: Also known as ascorbic acid, vitamin C is a powerful reducing agent. Its concentration can easily be identified when titrated with the blue dye Dichlorophenolindophenol ([[Dichlorophenolindophenol|DCPIP]]) which turns colorless when reduced by the vitamin.<ref>
| |
| {{Cite book
| |
| | title = Biology
| |
| | publisher = Taylor & Francis
| |
| | volume = 3
| |
| | year = 1967
| |
| | location = London
| |
| | pages = 52
| |
| }}</ref>
| |
| *[[Benedict's reagent]]: Excess [[glucose]] in urine may indicate [[diabetes]] in the patient. Benedict's method is the conventional method to quantify glucose in urine using a prepared reagent. In this titration, glucose reduces [[Copper|cupric]] ions to cuprous ions which react with [[potassium thiocyanate]] to produce a white precipitate, indicating the endpoint.<ref>
| |
| {{Cite book
| |
| | last = Nigam
| |
| | title = Lab Manual Of Biochemistry
| |
| | publisher = Tata McGraw-Hill Education
| |
| | year = 2007
| |
| | pages = 149
| |
| | isbn = 0-07-061767-8
| |
| }}</ref>
| |
| *[[Bromine number]]: A measure of [[Saturation (chemistry)|unsaturation]] in an analyte, expressed in milligrams of bromine absorbed by 100 grams of sample.
| |
| *[[Iodine number]]: A measure of unsaturation in an analyte, expressed in grams of iodine absorbed by 100 grams of sample.
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| | |
| ;Miscellaneous
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| | |
| *[[Karl Fischer titration]]: A potentiometric method to analyze trace amounts of water in a substance. A sample is dissolved in [[methanol]], and titrated with Karl Fischer reagent. The reagent contains iodine, which reacts proportionally with water. Thus, the water content can be determined by monitoring the [[Electric potential|potential]] of excess iodine.<ref>
| |
| {{Cite book
| |
| | last = Jackson
| |
| | first = M.L.
| |
| | coauthors = P. Barak
| |
| | title = Soil Chemical Analysis: Advanced Course
| |
| | publisher = UW-Madison Libraries Parallel Press
| |
| | year = 2005
| |
| | pages = 305–309
| |
| | isbn = 1-893311-47-3
| |
| }}</ref>
| |
| | |
| ==See also==
| |
| *[[Acid]]
| |
| *[[Base (chemistry)|Base]]
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| *[[Nonaqueous titration]]
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| | |
| ==References==
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| {{reflist|2}}
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| | |
| ==External links==
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| {{Commons category|Titration}}
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| *[http://www.wikihow.com/Perform-a-Titration Wikihow: Perform a Titration]
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| *[http://www.avogadro.co.uk/miscellany/titration/titreset.htm An interactive guide to titration]
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| *[http://scienceaid.co.uk/chemistry/applied/titration.html Science Aid: A simple explanation of titrations including calculation examples]
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| *[http://www2.iq.usp.br/docente/gutz/Curtipot_.html Titration freeware - simulation of any pH vs. volume curve, distribution diagrams and real data analysis]
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| {{Analytical chemistry}}
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| | |
| [[Category:Titration| ]]
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