Editing Titration (section)

===Acid–base titration===
{{Main|Acid–base titration}}
[[File:Acidobazna titracija 004.jpg|thumbnail|Methyl orange]]
{|border="2" cellpadding="5" align="center" style="text-align: center;" class=wikitable
|-
!style="background:#efefef;"|Indicator
!style="background:#efefef;"|Color on acidic side
!style="background:#efefef;"|Range of color change<br>(pH)
!style="background:#efefef;"|Color on basic side
|-
!style="background:#efefef;"|[[Methyl violet]]
| Yellow || 0.0—1.6 || Violet
|-
!style="background:#efefef;"|[[Bromophenol blue]]
| Yellow || 3.0—4.6 || Blue
|-
!style="background:#efefef;"|[[Methyl orange]]
| Red || 3.1—4.4 || Yellow
|-
!style="background:#efefef;"|[[Methyl red]]
| Red || 4.4—6.3 || Yellow
|-
!style="background:#efefef;"|[[Litmus]]
| Red || 5.0—8.0 || Blue
|-
!style="background:#efefef;"|[[Bromothymol blue]]
| Yellow || 6.0—7.6 || Blue
|-
!style="background:#efefef;"|[[Phenolphthalein]]
| Colorless || 8.3—10.0 || Pink
|-
!style="background:#efefef;"|[[Alizarine Yellow R|Alizarin yellow]]
| Yellow || 10.1—12.0 || Red
|}

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]] is added to the titration chamber, representing 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>
{{Cite web
  | title = pH measurements with indicators
  | url = http://www.ph-meter.info/pH-measurements-indicators
  | access-date = 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.

For very strong bases, such as [[organolithium reagent]], [[metal amides]], and [[hydride]]s, water is generally not a suitable solvent and indicators whose [[pKa]] are in the range of aqueous pH changes are of little use. Instead, the titrant and indicator used are much weaker acids, and anhydrous solvents such as [[THF]] are used.<ref>{{cite web |website=shenvilab.org/education |url= https://www.shenvilab.org/_files/ugd/24e834_d95df1b2e78146659e0e20809de02a9e.pdf | title=Titrating Soluble RM, R<sub>2</sub>NM and ROM Reagents }}</ref><ref>{{cite web |url= https://www.chem.tamu.edu/rgroup/gladysz/documents/alkylreagents.pdf |title= Methods for Standardizing Alkyllithium Reagents (literature through 2006) |access-date= 2014-06-04 }}</ref>

[[File:Faint pink color of Phenolphthalein.jpg|thumb|Phenolphthalein, a commonly used indicator in acid and base titration.]]

The approximate pH during titration can be approximated by three kinds of calculations. Before beginning of titration, the concentration of <chem>[H+]</chem> is calculated in an aqueous solution of weak acid before adding any base. When the number of moles of bases added equals the number of moles of initial acid or so called [[equivalence point]], one of hydrolysis and the pH is calculated in the same way that the conjugate bases of the acid titrated was calculated. Between starting and end points, <chem>[H+]</chem> is obtained from the [[Henderson–Hasselbalch equation|Henderson-Hasselbalch equation]] and titration mixture is considered as buffer. In Henderson-Hasselbalch equation the {{chem|[acid]}} and {{chem|[base]}} are said to be the molarities that would have been present even with dissociation or hydrolysis. In a buffer, <chem>[H+]</chem> can be calculated exactly but the dissociation of {{chem|HA}}, the hydrolysis of <chem>A-</chem> and self-ionization of water must be taken into account.<ref name=Harris>{{cite book|title = Quantitative Chemical Analysis |edition = Seventh|first= Daniel C.|last = Harris|publisher = Freeman and Company |date=2007|isbn =978-0-7167-7041-1|url = https://www.academia.edu/32945832}}</ref> Four independent equations must be used:<ref>
{{Cite book
  | last1 = Skoog
  | first1 = D.A.
  | last2 = West
  | first2 = D.M.
  | last3 = Holler
  | first3 = F.J.
  | title = Analytical Chemistry: An Introduction, seventh edition
  | publisher = Emily Barrosse
  | year = 2000
  | pages = [https://archive.org/details/isbn_9780030202933/page/265 265-305]
  | isbn = 0-03-020293-0
  | url = https://archive.org/details/isbn_9780030202933/page/265
  }}</ref>
:<math chem>[\ce{H+}][\ce{OH-}] = 10^{-14}</math>
:<math chem>[\ce{H+}] = K_a\ce{\frac{[HA]}{[A^{-}]}}</math>
:<math chem>[\ce{HA}] + [\ce{A-}] = \frac{(n_\ce{A} + n_\ce{B})}{V}</math>
:<math chem>[\ce{H+}]  + \frac{n_\ce{B}}{V} = [\ce{A-}] + [\ce{OH-}]</math>
In the equations, <math chem>n_\ce{A}</math> and <math chem>n_\ce{B}</math> are the moles of acid ({{chem|HA}}) and salt ({{chem|XA}} where X is the cation), respectively, used in the buffer, and the volume of solution is {{mvar|V}}. The [[law of mass action]] is applied to the ionization of water and the dissociation of acid to derived the first and second equations. The mass balance is used in the third equation, where the sum of <math chem>V[\ce{HA}]</math> and <math chem>V[\ce{A-}]</math> must equal to the number of moles of dissolved acid and base, respectively. Charge balance is used in the fourth equation, where the left hand side represents the total charge of the cations and the right hand side represents the total charge of the anions: <math chem>\frac{n_\ce{B}}{V}</math> is the molarity of the cation (e.g. sodium, if sodium salt of the acid or sodium hydroxide is used in making the buffer).<ref>
{{Cite book
  | last = Henry
  | first = N.
  |author2= M.M. Senozon
   | title = The Henderson-Hasselbalch Equation: Its History and Limitations
  | publisher = Journal of Chermical Education
  | year = 2001
  | pages = 1499–1503
}}</ref>