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An inorganic nonaqueous solvent is a solvent other than water, that is not an organic compound. These solvents are used in chemical research and industry for reactions that cannot occur in aqueous solutions or require a special environment. Inorganic nonaqueous solvents can be classified into two groups, protic solvents and aprotic solvents. Early studies on inorganic nonaqueous solvents evaluated ammonia, hydrogen fluoride, sulfuric acid, as well as more specialized solvents, hydrazine, and selenium oxychloride.^[1]

Protic inorganic nonaqueous solvents

Prominent members include ammonia, hydrogen fluoride, sulfuric acid, hydrogen cyanide. Ammonia (and several amines as well) are useful for the generating solutions of highly reducing species because the N-H bond resists reduction. The chemistry of electrides and alkalides relies on amine solvents.

The combination of HF and SbF₅ is the basis of a superacid solution. Using this mixture, the conjugate acid of hydrogen sulfide can be isolated:^[2]

H₂S + HF + SbF₅ → [H₃S]SbF₆

Autoionization

The limiting acid in a given solvent is the solvonium ion, such as H₃O⁺ (hydronium) ion in water. An acid which has more of a tendency to donate a hydrogen ion than the limiting acid will be a strong acid in the solvent considered, and will exist mostly or entirely in its dissociated form. Likewise, the limiting base in a given solvent is the solvate ion, such as OH^- (hydroxide) ion, in water. A base which has more affinity for protons than the limiting base cannot exist in solution, as it will react with the solvent.

For example, the limiting acid in liquid ammonia is the ammonium ion, NH₄⁺ which has a pK_a value in water of 9.25. The limiting base is the amide ion, NH₂^-. NH₂⁻ is a stronger base than the hydroxide ion and so cannot exist in aqueous solution. The pK_a value of ammonia is estimated to be approximately 34 (c.f. water, 14^[3]^[4]).

Aprotic inorganic nonaqueous solvents

Prominent members include sulfur dioxide, sulfuryl chloride fluoride, dinitrogen tetroxide, antimony trichloride, and bromine trifluoride. These solvents have proven useful for study highly electrophilic or highly oxidizing compounds or ions. Several (SO₂, SO₂ClF, N₂O₄) are gases near room temperature, so they are handled using vacuum-line techniques.

The generation of [IS₇]⁺ and [BrS₇]⁺ are illustrative. These highly electrophilic salts are prepared in SO₂ solution.^[5] The preparation of [SBr₃]⁺ salts also calls for a mixed solvent composed of SO₂ and SO₂FCl.^[6] Sulfuryl chloride fluoride is often used for the synthesis of noble gas compounds.^[7]

Autoionization

Many inorganic solvents participate in autoionization reactions. In the solvent system definition of acids and bases, autoionization of solvents affords the equivalent to acids and bases. Relevant autoionizations:

2BrF₃

{\ce {<=>>}}

BrF₂⁺ + BrF₄⁻

N₂O₄ ⇌ NO⁺ (nitrosonium) + NO₃⁻ (nitrate)

2SbCl₃ ⇌ SbCl₂⁺ + SbCl₄⁻

2POCl₃ ⇌ POCl₂⁺ + POCl₄⁻

According to the solvent-system definition, acids are the compounds that increase the concentration of the solvonium (positive) ions, and bases are the compounds that result in the increase of the solvate (negative) ions, where solvonium and solvate are the ions found in the pure solvent in equilibrium with its neutral molecules:

The solvent SO₂ is relatively uncomplicated^[how?], it does not autoionize.

References

^ Audrieth, Ludwig Frederick (1953). Non-aqueous Solvents; Applications as Media for Chemical Reactions. Wiley.
^ Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. p. 682. ISBN 978-0-08-037941-8.
^ Meister, Erich C.; Willeke, Martin; Angst, Werner; Togni, Antonio; Walde, Peter (2014). "Confusing Quantitative Descriptions of Brønsted-Lowry Acid-Base Equilibria in Chemistry Textbooks – A Critical Review and Clarifications for Chemical Educators". Helvetica Chimica Acta. 97 (1): 1–31. doi:10.1002/hlca.201300321. ISSN 1522-2675.
^ Silverstein, Todd P.; Heller, Stephen T. (2017-06-13). "pKa Values in the Undergraduate Curriculum: What Is the Real pKa of Water?". Journal of Chemical Education. 94 (6): 690–695. Bibcode:2017JChEd..94..690S. doi:10.1021/acs.jchemed.6b00623. ISSN 0021-9584.
^ Murchie, M. P.; Passmore, J.; Wong, C.-M. (1990). "Iodine and Bromine Polysulfur Hexafluoroarsenate(V) and Hexafluoroantimonate(V)". Inorganic Syntheses. Inorganic Syntheses. Vol. 27. pp. 332–339. doi:10.1002/9780470132586.ch67. ISBN 9780470132586.
^ Murchie, Mike; Passmore, Jack (1986). "Tribromosulfur(IV) Hexafluoroarsenate(V)". Inorganic Syntheses. Inorganic Syntheses. Vol. 24. pp. 76–79. doi:10.1002/9780470132555.ch23. ISBN 9780470132555.
^ Koppe, Karsten; Bilir, Vural; Frohn, Hermann-J.; Mercier, Hélène P. A.; Schrobilgen, Gary J. (2007). "Syntheses, Solution Multi-NMR Characterization, and Reactivities of [C₆F₅Xe]+Salts of Weakly Coordinating Borate Anions, [BY₄]⁻ (Y = CF₃, C₆F₅, CN, or OTeF₅)". Inorganic Chemistry. 46 (22): 9425–9437. doi:10.1021/ic7010138. PMID 17902647.

External links

Media related to Inorganic solvents at Wikimedia Commons

[1] Audrieth, Ludwig Frederick (1953). Non-aqueous Solvents; Applications as Media for Chemical Reactions. Wiley.

[2] Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. p. 682. ISBN 978-0-08-037941-8.

[3] Meister, Erich C.; Willeke, Martin; Angst, Werner; Togni, Antonio; Walde, Peter (2014). "Confusing Quantitative Descriptions of Brønsted-Lowry Acid-Base Equilibria in Chemistry Textbooks – A Critical Review and Clarifications for Chemical Educators". Helvetica Chimica Acta. 97 (1): 1–31. doi:10.1002/hlca.201300321. ISSN 1522-2675.

[4] Silverstein, Todd P.; Heller, Stephen T. (2017-06-13). "pKa Values in the Undergraduate Curriculum: What Is the Real pKa of Water?". Journal of Chemical Education. 94 (6): 690–695. Bibcode:2017JChEd..94..690S. doi:10.1021/acs.jchemed.6b00623. ISSN 0021-9584.

[5] Murchie, M. P.; Passmore, J.; Wong, C.-M. (1990). "Iodine and Bromine Polysulfur Hexafluoroarsenate(V) and Hexafluoroantimonate(V)". Inorganic Syntheses. Inorganic Syntheses. Vol. 27. pp. 332–339. doi:10.1002/9780470132586.ch67. ISBN 9780470132586.

[6] Murchie, Mike; Passmore, Jack (1986). "Tribromosulfur(IV) Hexafluoroarsenate(V)". Inorganic Syntheses. Inorganic Syntheses. Vol. 24. pp. 76–79. doi:10.1002/9780470132555.ch23. ISBN 9780470132555.

[7] Koppe, Karsten; Bilir, Vural; Frohn, Hermann-J.; Mercier, Hélène P. A.; Schrobilgen, Gary J. (2007). "Syntheses, Solution Multi-NMR Characterization, and Reactivities of [C₆F₅Xe]+Salts of Weakly Coordinating Borate Anions, [BY₄]⁻ (Y = CF₃, C₆F₅, CN, or OTeF₅)". Inorganic Chemistry. 46 (22): 9425–9437. doi:10.1021/ic7010138. PMID 17902647.

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+{{Short description|Type of nonaqueous solvents}}
-An ''inorganic nonaqueous solvent'' is a [[solvent]] other than water, that is not an [[organic compound]]. Common examples are liquid [[ammonia]] and liquid [[sulfur dioxide]]. These solvents are used industrially and in chemical research. These solvents are used for reactions that cannot occur in aqueous solutions.
+An '''inorganic nonaqueous solvent''' is a [[solvent]] other than water, that is not an [[organic compound]]. These solvents are used in chemical research and industry for reactions that cannot occur in aqueous solutions or require a special environment. Inorganic nonaqueous solvents can be classified into two groups, protic solvents and aprotic solvents.  Early studies on inorganic nonaqueous solvents evaluated ammonia, hydrogen fluoride, sulfuric acid,  as well as more specialized solvents, hydrazine, and selenium oxychloride.<ref>{{cite book|title=Non-aqueous Solvents; Applications as Media for Chemical Reactions| author= Audrieth, Ludwig Frederick|publisher=Wiley|year=1953}}</ref>
+==Protic inorganic nonaqueous solvents==
-== Acid-base chemistry ==
+Prominent members include [[ammonia]], [[hydrogen fluoride]], [[sulfuric acid]], [[hydrogen cyanide]].
-The [[Acid-base reaction theories|Brønsted-Lowry theory]] of acids can be extended to non-aqueous
+Ammonia (and several amines as well) are useful for the generating solutions of highly reducing species because the N-H bond resists reduction.  The chemistry of [[electride]]s and [[alkalide]]s relies on amine solvents.
-solvents which possess one or more hydrogen atom which can dissociate:
-such solvent are known as '''protic solvents'''.
+The combination of HF and SbF<sub>5</sub> is the basis of a [[superacid]] solution.  Using this mixture, the conjugate acid of [[hydrogen sulfide]] can be isolated:<ref>{{Greenwood&Earnshaw2nd|page=682}}</ref>
-=== Strong acids and weak acids ===
+:H<sub>2</sub>S  +  HF + SbF<sub>5</sub>  →  [H<sub>3</sub>S]SbF<sub>6</sub>
-A [[strong acid]] is an acid which exists mostly or entirely in its
-dissociated form, that is to say that the equilibrium
-:HA {{unicode|⇌}} H<sup>+</sup>([[Solvation|solvated]]) + A<sup>−</sup>
-is far to the right. In water, a strong acid is normally taken to be one with a
-[[Acid dissociation constant|p''K''<sub>a</sub> value]] of less than that of hydronium, -1.74. An example of a strong acid in water is [[hydrochloric acid]].
+===Autoionization===
-A [[weak acid]] may exist mostly in its undissociated form: this is the
+The limiting acid in a given solvent is the solvonium ion, such as H<sub>3</sub>O<sup>+</sup> ([[hydronium]]) ion in water. An acid which has more of a tendency to donate a [[hydrogen ion]] than the limiting acid will be a strong acid in the solvent considered, and will exist mostly or entirely in its dissociated form. Likewise, the limiting base in a given solvent is the solvate ion, such as OH<sup>-</sup> ([[hydroxide]]) ion, in water. A base which has more affinity for protons than the limiting base cannot exist in solution, as it will react with the solvent.
-case for [[acetic acid]] in water.
+For example, the limiting acid in liquid ammonia is the [[ammonium]] ion, NH<sub>4</sub><sup>+</sup> which has a p''K''<sub>a</sub> value in water of 9.25. The limiting base is the [[amide]] ion, NH<sub>2</sub><sup>-</sup>. NH<sub>2</sub><sup>−</sup> is [[Leveling effect|a stronger base than the hydroxide ion]] and so cannot exist in aqueous solution. The p''K''<sub>a</sub> value of ammonia is estimated to be approximately 34 (''c.f.'' water, 14<ref>{{Cite journal|last1=Meister|first1=Erich C.|last2=Willeke|first2=Martin|last3=Angst|first3=Werner|last4=Togni|first4=Antonio|last5=Walde|first5=Peter|date=2014|title=Confusing Quantitative Descriptions of Brønsted-Lowry Acid-Base Equilibria in Chemistry Textbooks – A Critical Review and Clarifications for Chemical Educators|journal=Helvetica Chimica Acta|language=en|volume=97|issue=1|pages=1–31|doi=10.1002/hlca.201300321|issn=1522-2675}}</ref><ref>{{Cite journal|last1=Silverstein|first1=Todd P.|last2=Heller|first2=Stephen T.|date=2017-06-13|title=pKa Values in the Undergraduate Curriculum: What Is the Real pKa of Water?|journal=Journal of Chemical Education|volume=94|issue=6|pages=690–695|doi=10.1021/acs.jchemed.6b00623|bibcode=2017JChEd..94..690S|issn=0021-9584}}</ref>).
-=== Limiting acids ===
-The limiting acid in a given solvent is the solvated form of the
-hydrogen ion. In water, this is usually denoted H<sub>3</sub>O<sup>+</sup> and
-known as the [[hydronium]] ion. An acid which has more of a tendency to
-donate a hydrogen ion than the limiting acid will be a strong acid in the
-solvent considered, and will exist mostly or entirely in its dissociated form.
+==Aprotic inorganic nonaqueous solvents==
-=== Limiting bases ===
+Prominent members include [[sulfur dioxide]], [[sulfuryl chloride fluoride]], [[dinitrogen tetroxide]], [[antimony trichloride]], and [[bromine trifluoride]]. These solvents have proven useful for study highly electrophilic or highly oxidizing compounds or ions.  Several (SO<sub>2</sub>, SO<sub>2</sub>ClF, N<sub>2</sub>O<sub>4</sub>) are gases near room temperature, so they are handled using [[Vacuum line|vacuum-line]] techniques.
-The limiting base in a given solvent is the ion derived from [[deprotonation]]
-of a solvent molecule. In water, this is the [[hydroxide]] ion, OH<sup>−</sup>.
-A base which has more affinity for protons than the limiting base cannot exist in solution,
-as it will react with the solvent.
+The generation of [IS<sub>7</sub>]<sup>+</sup> and [BrS<sub>7</sub>]<sup>+</sup> are illustrative.  These highly electrophilic salts are prepared in SO<sub>2</sub> solution.<ref>{{cite book |doi=10.1002/9780470132586.ch67|chapter=Iodine and Bromine Polysulfur Hexafluoroarsenate(V) and Hexafluoroantimonate(V)|series=Inorganic Syntheses|last1=Murchie|first1=M. P.|last2=Passmore|first2=J.|last3=Wong|first3=C.-M.|title=Inorganic Syntheses|pages=332–339|volume=27|year=1990|isbn=9780470132586}}</ref>
-=== Liquid ammonia ===
+The preparation of [SBr<sub>3</sub>]<sup>+</sup> salts also calls for a mixed solvent composed of SO<sub>2</sub> and SO<sub>2</sub>FCl.<ref>{{cite book |doi=10.1002/9780470132555.ch23|chapter=Tribromosulfur(IV) Hexafluoroarsenate(V)|series=Inorganic Syntheses|last1=Murchie|first1=Mike|last2=Passmore|first2=Jack|title=Inorganic Syntheses|pages=76–79|volume=24|year=1986|isbn=9780470132555}}</ref> Sulfuryl chloride fluoride is often used for the synthesis of [[noble gas compound]]s.<ref>{{cite journal |doi=10.1021/ic7010138|title=Syntheses, Solution Multi-NMR Characterization, and Reactivities of &#91;C<sub>6</sub>F<sub>5</sub>Xe&#93;+Salts of Weakly Coordinating Borate Anions, &#91;BY<sub>4</sub>&#93;<sup>−</sup> (Y = CF<sub>3</sub>, C<sub>6</sub>F<sub>5</sub>, CN, or OTeF<sub>5</sub>)|year =2007|last1=Koppe|first1 =Karsten|last2 =Bilir|first2 =Vural|last3=Frohn|first3=Hermann-J.|last4=Mercier|first4=Hélène P. A.|last5=Schrobilgen|first5=Gary J.|journal=Inorganic Chemistry|volume=46|issue=22|pages=9425–9437|pmid=17902647}}</ref>
-The limiting acid in liquid ammonia is the [[ammonium]] ion, which has a p''K''<sub>a</sub> value
+<!--The chemistry of [[xenon]] compounds is often conducted in [[hydrogen fluoride]] or [[bromine pentafluoride]], which dissolve readily both [[xenon difluoride]]s and its multiple derivatives,<ref>Pointner BE, Suontamo RJ, Schrobilgen GJ. Syntheses and X-ray crystal structures of alpha- and beta-{{chem|[XeO|2|F][SbF|6|]}}, {{chem|[XeO|2|F][AsF|6|]}}, {{chem|[FO|2|XeFXeO|2|F][AsF|6|]}}, and {{chem|[XeF|5|][SbF|6|)]·XeOF|4}} and computational studies of the {{chem|XeO|2|F|+}} and {{chem|FO|2|XeFXeO|2|F|+}} cations and related species. ''Inorg Chem.'' 2006 Feb 20;45(4):1517-34.</ref>  [[Sulfuryl chloride fluoride]] is also useful for strong oxidants.<ref>Mercier HP, Moran MD, Sanders JC, Schrobilgen GJ, Suontamo RJ. "Synthesis, structural characterization, and computational study of the strong oxidant salt {{chem|[XeOTeF|5|][Sb(OTeF|5|)|6|]·SO|2|ClF}}." ''Inorg Chem.'' 2005 Jan 10;44(1):49-60.</ref>
-in water of 9.25. The limiting base is the [[amide]] ion, NH<sub>2</sub><sup>−</sup>. This is a
-stronger base than the hydroxide ion and so cannot exist in aqueous solution. The p''K''<sub>a</sub> value
-of ammonia is estimated to be approximately 34 (''c.f.'' water, 15.74).
+[[Sulfuryl chloride fluoride]] is the solvent of choice for handling extreme oxidants. For example, it can be used to generate and study free [[carbocations]]<ref>Mercier HP, Moran MD, Schrobilgen GJ, Steinberg C, Suontamo RJ. The syntheses of carbocations by use of the noble-gas oxidant, {{chem|[XeOTeF|5|][Sb(OTeF|5|)|6|]}}: the syntheses and characterization of the {{chem|CX|3|+}} (X = Cl, Br, {{chem|OTeF|5}}) and {{chem|CBr(OTeF|5|)|2|+}} cations and theoretical studies of {{chem|CX|3|+}} and {{chem|BX|3}} (X = F, Cl, Br, I, {{chem|OTeF|5}}). ''J Am Chem Soc.'' 2004 May 5;126(17):5533-48.</ref> and [[arenium ion]]s.<ref>V D Shteingarts, Polyfluorinated Arenonium Ions, ''Russ. Chem Rev'' 1981;50(8):735-748.</ref>
-Any acid which is a stronger acid than the ammonium ion will be a strong acid in liquid ammonia. This is the
+-->
-case for acetic acid, which is completely dissociated in liquid ammonia solution. The addition of pure acetic
-acid and the addition of ammonium acetate have exactly the same effect on a liquid ammonia solution: the increase
-in its acidity: in practice, the latter is preferred for safety reasons.
+===Autoionization===
-Bases can exist in solution in liquid ammonia which cannot exist in aqueous solution: this is the case for any
+Many inorganic solvents participate in [[autoionization]] reactions.  In the solvent system definition of acids and bases, autoionization of solvents affords the equivalent to acids and bases.  Relevant autoionizations:
-base which is stronger than the hydroxide ion but weaker than the amide ion. Many carbon anions can be formed
+: 2BrF<sub>3</sub>  {{EqmR}}  BrF<sub>2</sub><sup>+</sup>  +  BrF<sub>4</sub><sup>−</sup>
-in liquid ammonia solution by the action of the amide ion on organic molecules (see [[sodium amide]] for examples).
+: N<sub>2</sub>O<sub>4</sub> ⇌ NO<sup>+</sup> ([[nitrosonium]]) + NO<sub>3</sub><sup>−</sup> ([[nitrate]])
+: 2SbCl<sub>3</sub> ⇌ SbCl<sub>2</sub><sup>+</sup>  + SbCl<sub>4</sub><sup>−</sup>
+: 2POCl<sub>3</sub> ⇌ POCl<sub>2</sub><sup>+</sup> + POCl<sub>4</sub><sup>−</sup>
+According to the [[Acid-base reaction#Solvent system definition|solvent-system definition]], acids are the compounds that increase the concentration of the '''solvonium''' (positive) ions, and bases are the compounds that result in the increase of the '''solvate''' (negative) ions, where solvonium and solvate are the ions found in the pure solvent in equilibrium with its neutral molecules:
-=== Superacids ===
-A [[superacid]] is a medium in which the hydrogen ion is only very weakly solvated. The classic example is a
-mixture of [[antimony pentafluoride]] and liquid [[hydrogen fluoride]]:
-:SbF<sub>5</sub> + HF {{unicode|⇌}} H<sup>+</sup> + SbF<sub>6</sub><sup>−</sup>
-The limiting base, the hexfluoroantimonate anion SbF<sub>6</sub><sup>−</sup>, is so weakly attracted to the hydrogen ion that virtually any other base will bind more strongly:
-hence, this mixture can be used to protonate organic molecules which would not be considered bases in other solvents.
+The solvent SO<sub>2</sub> is relatively uncomplicated{{How |date=February 2024}}, it does not autoionize.
-=== Comparisons of acidity and basicity between solvents ===
-There exists a large corpus of data concerning acid strengths in aqueous solution (p''K''<sub>a</sub> values), and it
-is tempting to transfer this to other solvents. Such comparisons are, however, fraught with danger, as they only consider the effect
-of solvation on the stability of the hydrogen ion, while neglecting its effects on the stability of the other species involved in the equilibrium.
-Gas phase acidities (normally known as [[Proton affinity|proton affinities]]) can be measured, and their relative order is often quite different
-from that of the aqueous acidities of the corresponding acids. Few quantitative studies on acidities in nonaqueous solvents have been carried out, although some qualitative data are available.
-It appears that most acids which have a p''K''<sub>a</sub> value of less than 9 in water are indeed strong acids in liquid ammonia. However, the hydroxide ion is often a much stronger base in
-nonaqueous solvents (e.g. liquid ammonia, [[dimethyl sulfoxide|DMSO]]) than in water.
-It should be noted that [[pH]] values are at present undefined in nonaqueous solvents, as the definition of pH assumes an aqueous solution.
-== Non-protonic solvents ==
-The most common solvents is made up by one or more hydrogen atoms. But solvents can also be made up by anything but hydrogen atoms.
-=== Acid-Base in aprotic solvents ===
-In [[aprotic solvent]]s (solvents without hydrogen), acid/base reactions are different from a hydrogen based solvent. In those solvents an acid gives away a hydrogen atom, and a base takes hydrogen atoms. To define the acid and base in aprotic solvents one must use [[Lewis acid]] and [[Lewis base]]. In such solvents an acid is the atom/molecule that accepts an electron pair, and a base is an electron pair giver.
-Example ([[phosphoryl chloride]]):
-:OPCl<sub>3</sub> ⇌ OPCl<sub>2</sub><sup>+</sup> + Cl<sup>-</sup>
-This is the [[Dissociation (chemistry)|self-ionization]] of OPCl<sub>3</sub>. When an acid for this solvent is added OPCl<sub>2</sub><sup>+</sup> is created, and a base gives rise to more Cl<sup>-</sup> ions. In reality the Cl<sup>-</sup> ion is actually a OPCl<sub>3</sub> with an extra Cl<sup>-</sup> on it, i.e. OPCl<sub>4</sub><sup>-</sup>. In this solvent you would use pCl instant of pH.
 ==See also==
 *[[Nonaqueous titration]]
+*[[Protic solvent]]
+== References ==
+{{reflist}}
+==External links==
+*{{Commons category-inline|Inorganic solvents}}
 {{Chemical solutions}}
+[[Category:Inorganic solvents| ]]
 [[Category:Solvents]]
-[[Category:Solutions]]

v t e Chemical solutions
Solution	Ideal solution Aqueous solution Solid solution Buffer solution Flory–Huggins Mixture Suspension Colloid Phase diagram Phase separation Eutectic point Alloy Saturation Supersaturation Serial dilution Dilution (equation) Apparent molar property Miscibility gap
Concentration and related quantities	Molar concentration Mass concentration Number concentration Volume concentration Normality Molality Mole fraction Mass fraction Isotopic abundance Mixing ratio Ternary plot
Solubility	Solubility equilibrium Total dissolved solids Solvation Solvation shell Enthalpy of solution Lattice energy Raoult's law Henry's law Solubility table (data) Solubility chart Miscibility
Solvent	(Category) Acid dissociation constant Protic solvent Polar aprotic solvent Inorganic nonaqueous solvent Solvation List of boiling and freezing information of solvents Partition coefficient Polarity Hydrophobe Hydrophile Lipophilic Amphiphile Lyonium ion Lyate ion