Inorganic nonaqueous solvent: Difference between revisions
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The [[Acid-base reaction theories|Brönsted theory]] of acids can be extended to non-aqueous |
The [[Acid-base reaction theories|Brönsted theory]] of acids can be extended to non-aqueous |
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solvents which possess one or more hydrogen atom which can dissociate: |
solvents which possess one or more hydrogen atom which can dissociate: |
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such solvent are known as '''protic |
such solvent are known as '''protic solvents'''. |
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=== Strong acids and weak acids === |
=== Strong acids and weak acids === |
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A [[strong acid]] is an acid which exists mostly or entirely in its |
A [[strong acid]] is an acid which exists mostly or entirely in its |
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dissociated form, that is to say that the equilibrium |
dissociated form, that is to say that the equilibrium |
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:HA {{unicode| |
:HA {{unicode|⇌}} H<sup>+</sup>([[Solvation|solvated]]) + A<sup>−</sup> |
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is far to the right. In water, a strong acid is normally taken to be one with a |
is far to the right. In water, a strong acid is normally taken to be one with a |
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[[Acid dissociation constant|p''K''<sub>a</sub> value]] of less than that of hydronium, taken as zero by convention, e.g. [[hydrochloric acid]]. |
[[Acid dissociation constant|p''K''<sub>a</sub> value]] of less than that of hydronium, taken as zero by convention, e.g. [[hydrochloric acid]]. |
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=== Limiting bases === |
=== Limiting bases === |
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The limiting base in a given solvent is the ion derived from [[deprotonation]] |
The limiting base in a given solvent is the ion derived from [[deprotonation]] |
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of a solvent molecule. In water, this is the [[hydroxide]] ion, OH<sup> |
of a solvent molecule. In water, this is the [[hydroxide]] ion, OH<sup>−</sup>. |
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A base which has more affinity for protons than the limiting base cannot exist in solution, |
A base which has more affinity for protons than the limiting base cannot exist in solution, |
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as it will react with the solvent. |
as it will react with the solvent. |
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=== Liquid ammonia === |
=== Liquid ammonia === |
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The limiting acid in liquid ammonia is the [[ammonium]] ion, which has a p''K''<sub>a</sub> value |
The limiting acid in liquid ammonia is the [[ammonium]] ion, which has a p''K''<sub>a</sub> value |
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in water of 9.25. The limiting base is the [[amide]] ion, NH<sub>2</sub><sup> |
in water of 9.25. The limiting base is the [[amide]] ion, NH<sub>2</sub><sup>−</sup>. This is a |
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stronger base than the hydroxide ion and so cannot exist in aqueous solution. The p''K''<sub>a</sub> value |
stronger base than the hydroxide ion and so cannot exist in aqueous solution. The p''K''<sub>a</sub> value |
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of ammonia is estimated to be approximately 34 (''c.f.'' water, 14). |
of ammonia is estimated to be approximately 34 (''c.f.'' water, 14). |
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A [[superacid]] is a medium in which the hydrogen ion is only very weakly solvated. The classic example is a |
A [[superacid]] is a medium in which the hydrogen ion is only very weakly solvated. The classic example is a |
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mixture of [[antimony pentafluoride]] and liquid [[hydrogen fluoride]]: |
mixture of [[antimony pentafluoride]] and liquid [[hydrogen fluoride]]: |
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:SbF<sub>5</sub> + HF {{unicode| |
:SbF<sub>5</sub> + HF {{unicode|⇌}} H<sup>+</sup> + SbF<sub>6</sub><sup>−</sup> |
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The limiting base, the hexfluoroantimonate anion SbF<sub>6</sub><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: |
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hence, this mixture can be used to protonate organic molecules which would not be considered bases in other solvents. |
hence, this mixture can be used to protonate organic molecules which would not be considered bases in other solvents. |
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=== Acid-Base in none-protonic solvents === |
=== Acid-Base in none-protonic solvents === |
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In this kind of solvents acid base reactions goes different from a hydrogen based solvent. In those solvents |
In this kind of solvents acid base reactions goes 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 such 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. |
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Example, OPCl<sub>3</sub> |
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OPCL<sub>3</sub> <-> OPCL<sub>2</sub><sup>+</sup> + CL<sup>-</sup> |
OPCL<sub>3</sub> <-> OPCL<sub>2</sub><sup>+</sup> + CL<sup>-</sup> |
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This is the [[self-ionization]] of OPCl<sub>3</sub>. When |
This is the [[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 acctualy a OPCL<sub>3</sub> with an extra CL<sup>-</sup> on it. OPCL<sub>4</sub><sup>-</sup>. In this solvent yo would use pCl instant of pH |
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Revision as of 19:06, 27 July 2006
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.
Acid-base chemistry
The Brönsted theory of acids can be extended to non-aqueous solvents which possess one or more hydrogen atom which can dissociate: such solvent are known as protic solvents.
Strong acids and weak acids
A strong acid is an acid which exists mostly or entirely in its dissociated form, that is to say that the equilibrium
- HA ⇌ H+(solvated) + A−
is far to the right. In water, a strong acid is normally taken to be one with a pKa value of less than that of hydronium, taken as zero by convention, e.g. hydrochloric acid.
A weak acid may exist mostly in its undissociated form: this is the case for acetic acid in water.
Limiting acids
The limiting acid in a given solvent is the solvated form of the hydrogen ion. In water, this is usually denoted H3O+ 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.
Limiting bases
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−. A base which has more affinity for protons than the limiting base cannot exist in solution, as it will react with the solvent.
Liquid ammonia
The limiting acid in liquid ammonia is the ammonium ion, which has a pKa value in water of 9.25. The limiting base is the amide ion, NH2−. This is a stronger base than the hydroxide ion and so cannot exist in aqueous solution. The pKa value of ammonia is estimated to be approximately 34 (c.f. water, 14).
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.
Bases can exist in solution in liquid ammonia which cannot exist in aqueous solution: this is the case for any base which is stronger than the hydroxide ion but weaker than the amide ion. Many carbon anions can be formed in liquid ammonia solution by the action of the amide ion on organic molecules (see sodium amide for examples).
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:
- SbF5 + HF ⇌ H+ + SbF6−
The limiting base, the hexfluoroantimonate anion SbF6−, 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.
Comparisons of acidity and basicity between solvents
There exists a large corpus of data concerning acid strengths in aqueous solution (pKa 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 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 pKa 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, 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.
None-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 none-protonic solvents
In this kind of solvents acid base reactions goes 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 such 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, OPCl3 OPCL3 <-> OPCL2+ + CL- This is the self-ionization of OPCl3. When an acid for this solvent is added OPCL2+ is created, and a base gives rise to more CL- ions. In reality the CL- ion is acctualy a OPCL3 with an extra CL- on it. OPCL4-. In this solvent yo would use pCl instant of pH