(G S Manku: email: manku.gs@gmail.com)

[10.3].1 AMMONIA                                                                                                  

Ammonia is a hydrogen bonded liquid with a low boiling point (240 K) and low melting point (195.5 K). It has a lower dielectric constant (ε = 22 at 240 K) than water. Generally, ionic solutes have a lower solubility in ammonia. However, many salts dissolve in ammonia to appreciable extent especially (i) ammonium salts irrespective of anions, (ii) metal ions that can form ammines (Cu²⁺, Ni²⁺, Co²⁺, Cd²⁺, Ag⁺, Pd²⁺, Pt²⁺  and (iii) polarizable molecules like iodine. Most of nitrates, nitrites and cyanides and thiocyanates are soluble. Highly charged ions and ions with high charge density are generally insoluble (fluoride, oxide, hydroxide, sulphate, carbonate, phosphates) are insoluble. Solubility of halides decreases from highly polarizable iodides to chlorides, only NaCl and BeCl₂ are insoluble chlorides.

The studies in ammonia have to be made at low temperatures because of small liquid range and low B.P. and low M.P. Special apparatus is required for studies as ammonia is hygroscopic and absorbs moisture.

[10.3].1.1 Acid –Base Reactions in Ammonia

As given earlier, in ammonia as solvent, NH₄⁺ is acid whereas NH­₂^–, NH^2– and  N^3– ions (amide, imide and nitride respectively) are bases. Typical neutralization reactions in ammonia (solvent) are:

NH₄Cl + NaNH₂ → NaCl + 2NH₃
2NH₄Cl + PbNH → PbCl₂ + 3NH₃
3NH₄Cl + Na₃N → NaCl + 4NH₃
NH₄Cl + NaOH → NaCl + NH₃ + H₂O.

The neutralization reactions can be studied conductometrically or by using phenolphthalein as indicator.

Ammonium salts dissolve metal oxides and hydroxides:

2NH₄NO₃ + MgO → Mg(NO₃)₂ + 2NH₃ + H₂O

Metal ions exist in ammonia solutions as ammoniated ions or as ammine complexes.

Amphoteric behavior is shown by zinc(II), aluminum(III), tin(II), etc.: In ammonia solution, or by adding an amide solution to zinc chloride solution, a white precipitate of zinc amide (Zn(NH₂)₂ ) is obtained which dissolves in excess of solvent or addition of amide to form a clear solution containing [Zn(NH₂)₄]^{2 –} ions:

Zn(NO₃)₂ + 4NH₃ ⇌ Zn(NH₂)₂ (↓)+ NH₄NO₃
Zn(NH₂)₂ (↓)+ 4NH₃ ⇌ (NH₄)₂[Zn(NH₂)]

All aqueous acids with pK_a ≤ 12 are leveled off to NH₄⁺ ions and act as strong acids. Compounds showing no acidic or weak basic behavior in water may act as weak acids, e.g. amides, anilides, etc.

CH₃CONH₂ + NH₃ ⇌ NH₄⁺ + CH­₃CONH ^–

Bases in water are either insoluble or behave as weak bases. Only extremely strong bases, stronger than amide ions, are strong bases in ammonia:

H ^– + NH₃ ⇌ H₂ + NH₂ ^–
O^{2 –} + NH₃ ⇌ OH ^–  + NH₂^–

Ammonium salts dissolve active metals like Mg, Zn, etc., with the evolution of hydrogen

2NH₄Cl + Zn → ZnCl₂ + 2NH₃ + H₂’

[10.3].1.2 Metathetical Reaction

: When two ions giving a low solubility product are brought in contact in solutions, precipitation takes place. Some examples of metathetical precipitation reactions in ammonia are given below.

NaCl + KI → KCl ↓ + NaI
2NH₄I + Zn(NO₃)₂ → ZnI₂ ↓ + NH₄NO₃
2NH₄Br + Sr(NO₃)₂ →  SrBr₂↓ + 2NH₄NO₃
(NH₄)₂S + 2AgNO₃ → Ag₂S ↓ + 2 NH₄NO₃

Many compounds are precipitated as complex ammines from ammonia solutions. These may correspond to hydrated compounds in aqueous solutions, e.g. CaBr₂.2NH₃ or CaCl₂.8NH₃.

[10.3].1.3 Solvolysis in Ammonia – Ammonolysis

The solvolysis reactions in ammonia are called ammonolysis of ammolytical reactions and resemble the hydrolysis reactions in aqueous media:

Cl₂ + 2NH₃ → NH₂Cl + NH₄⁺ + Cl ^–
SiCl₄ + 8NH₃ → Si(NH₂)₄ + 4NH₄⁺ + 4Cl ^–
HgCl₂ + 2NH₃ → HgNH₄Cl + NH₄⁺ + Cl ^–
Hg₂Cl₂ + 2NH₃ → Hg + HgNH₄Cl + NH₄⁺ + Cl ^–
3ZrBr₄ + (17 + x)NH₃ →3Zr(NH₂)₄.7NH₄Br.xNH₃ + 5NH₄⁺ + 5Br ^–

In a manner similar to aqueous solutions, a pH scale can be set up in ammonia in which the activity of NH₄⁺ ions is 1.00 M, pH = 13.5 is neutral solution where [NH₄⁺] = [NH₂ ^–], and pH = 27 is 1.00 M NH₂^–.

[10.3].1.4 Complex Formation Reactions

Though not studied much, complex formation reactions in ammonia and water appear similar.

AgNO₃ + NH₄CNS ⇌ AgCNS + NH₄NO₃
AgCNS + NH₄CNS ⇌ NH₄[Ag(SCN)₂]

[10.3].1.5 Redox Reactions

In ammonia, strong oxidizing agents do not exist. Nitric acid is converted into ammonium nitrate and permanganate and dichromate are only weak oxidizing agents). On the other hand, reducing agents show enhanced activity. This is due to the presence of solvated electrons in ammonia. The oxidizing agents have to compete against the solvent to take up electrons from the reducing agents, and their activity gets reduced. For reducing agents, transfer of electrons is facilitated by the solvent, and their activity gets increased. Thus, alkali metals are very powerful reducing agents in ammonia. Some of the redox reactions in ammonia are given below.

4Na + 2NH₃ → 2NaOH + 2NaNH₂
2Na + 2NH₄Br → 2NaBr + 2NH₃ + H₂
4NaNH₂ + O₂ → 2NaOH + 2NaNO₂ + 2NH₃
K + N₂O + NH₃ → KOH + KNH₂ + N₂
K + KMnO₄ → K₂MnO₄
[Sn(NH₂)₆]^4– + I₂ → [Sn(NH₂)₆]^2- + 2I^–

[10.3].1.6 Solutions of metals in ammonia

If a small piece of alkali metal or alkaline earth metal is placed in ammonia, it dissolves forming a deep blue solution. Addition of more metal deepens the blue color till a bronze phase separates out and floats on the surface. Cesium does not give the bronze phase. Further addition of metal increases the concentration to bronze phase till a saturation point is reached and no more of the metal dissolves.

The blue colored phase has the following properties:

1. The blue color is independent of the amount of metal dissolved and shows a broad maximum at around 1430 nm.
2. The density of solution is similar to that of ammonia.
3. The conductance is in the range of the electrolytes, the difference fro different metal solutions being due to the conductance of the metal cations formed.
4. With increasing metal concentration, the conductivity first increases to a maximum, then decreases and passes through a minimum before increasing again.
5. The dilute solutions are paramagnetic, paramagnetism decreases with increasing concentration of metal. Lande’s g factor is similar to that of free electrons.
6. Blue solutions are stable. Their rate of decomposition is even less than 1 % per day. Catalyzed decomposition ultimately gives sodamide NaNH₂ with the evolution of hydrogen:

2Na + 2NH₃ → 2NaNH₂ + H₂

These properties can be explained by assuming that the metals in blue solution are present as solvated (ammoniated) cations [Na(NH₃)₆]⁺ and the valence shell electrons are trapped in the vacant spaces (size = 300 – 340 pm) (size of cavities) of ammonia molecules. The size of cavity is determined by the partial molar volumes of alkali metals in ammonia solutions. These are called solvated electrons ([e(NH₃)₆]^– and give the blue color to solutions due to the 1s→2p change.

The paramagnetism is due to the presence of free solvated electrons. The paramagnetism increases first due to increase in their concentration, but later start decreasing due to the formation of electron pairs in which two electrons with opposite spin get paired up ([e_2­(NH₃)_y]^2–).

Due to the presence of free electrons, the blue solutions are powerful reducing agents.

The bronze colored phase has the following properties:

1. The bronze color has special metallic luster.
2. The solutions have a low density.
3. The conductivity is of the order of metals.
4. Magnetic susceptibilities are similar to the metals.
5. Upon evaporation, the solutions leave a residue of metals as such or as ammoniated metals [Li(NH₃)­₄] or [M(NH₃)­₄] (M = Ca, Sr, Ba, Eu, Sm)

These properties are consistent with a model describing the bronze phase as alloys of these metals with ammoniated metals ions bound by solvated electrons in the same manner as in alloys or in molten metals.

In addition to alkali and alkaline earth metals (except beryllium), europium, samarium and ytterbium also form blue solutions in ammonia (all these metals have a low ionization energies).

Cathodic reduction of AlI₃, BeCl₂ and NH₄⁺X⁺ in ammonia also gives blue solutions containing solvated electrons. Other similar solvent like amines, ethers and hexamethylphosphoramide also form similar solutions.

Though solvated electron is known in aqueous solutions also, its half life is very low (about 10^–3 s).

[10.3].2 SULPHURIC ACID

Sulphuric acid has a very high dielectric constant (110 ± 10). This should make it a better solvent than water for ionic compounds. But its high chemical reactivity and high viscosity (245.4. mill poise, about 25 times that of water) makes it a much poor solvent. The high viscosity makes dissolution of solutes slow and difficult for the solvent to be removed from crystals.

In sulphuric acid, only important reactions are acid–base reactions. Sulphuric acid autoionizes as:

2H₂SO₄ ⇌  H₃SO₄⁺ + HSO₄^–

It is an acidic solvent. All the species basic in water are basic in sulphuric acid also. Even species neutral in water (alcohols, ethers, thiols, esters amides, etc.) accept proton in sulphuric acid solutions and act as bases:

OH^– + 2H₂SO₄ ⇌ H₃O⁺ + 2HSO₄^–
NH₃ + 2H₂SO₄ ⇌ NH₄⁺ + HSO₄^–
H₂O + H₂SO₄ ⇌ H₃O⁺ + HSO₄^–
EtOH + H₂SO₄ ⇌ EtOH₂⁺ + 2HSO₄^–
CO(NH₂)₂ + H₂SO₄ ⇌ H₂NCONH₂⁺ + HSO₄^–

Acetic acid and nitric acid also behave as bases in sulphuric acid medium:

CH₃COOH + H₂SO₄ ⇌ CH₃COOH₂⁺ + HSO₄^–
HNO₃ + 2H₂SO₄ ⇌ NO₂⁺ + H₃O⁺ + 2HSO₄^–

Due to the very strong tendency of sulphuric acid to donate protons, only a few substances act as acids. Examples include disulphuric acid (H₂S₂O₇), fluorosulphonic acid (FSO₃H) and chlorosulphonic acid (ClSO₃H).

H₂S₂O₇ + H₂SO₄ ⇌ H₃SO₄⁺ + HS₂O₇^–

 [10.3].3 HYDROGEN FLUORIDE

10.4.3.1 Solubility of Compounds: The high dielectric constant, (83.6), stronger hydrogen bonds, other physical properties and slight acidic character of hydrogen fluoride make it similar to water. But its high reactivity and solvating power dissolves only s few substances unreacted and restricts its comparison with water or ammonia as a a solvent.

Only fluorides, fluoroborates and perchlorates dissolve in hydrogen fluoride unreacted and ionize into ions:

KF + HF ⇌ K⁺ + HF₂^–

Most of the halides are insoluble and stable, or dissolve forming fluorides and evolving hydrogen halides. Oxides and hydroxides react violently to form fluoride and fluoro complexes forming water and fluoride.

TiBr₄ + 6HF ⇌ H₂[TiF₆] + 4HBr
ZrO₂ + 7HF ⇌ H₃[ZrF₇] + 2H₂O

Potassium and sodium sulphate dissolve in HF forming sulphuric acid (other sulphates are insoluble).The solution then reacts slowly to form fluorosulphonates:

K₂SO₄ + 4HF ⇌ 2K⁺ + H₂SO₄ + 2HF₂ ^–
H₂SO₄ + 2HF → H₂O + FSO₃^– + H₂F⁺

Active metals (zinc, tin, magnesium, etc.) react with HF evolving hydrogen gas and forming fluorides.

[10.3].3.2 Acid – Base Reactions

 Auto-ionization of hydrogen fluoride is

3HF ⇌ H₂F⁺ + HF₂^–

Bases in HF are the species that give fluoride ions, and those protonating HF are acids. Thus fluorides are bases. Alcohols, ethers, ketones and anhydrides are bases as the oxygen atom in these compounds gets protonated in HF medium:

ROH + 2HF ⇌ ROH₂⁺ + 2HF₂^–
CH₃COCH₃ + 2HF ⇌ (CH₃)COH⁺ + 2HF₂^–
R–O–Y + 2HF ⇌ R – OH⁺–Y + HF₂^–

Solutions of strong Lewis acids like SbF₃ or BF₃ are strong acids with apparent protonation of solvent and are called super acids as they protonate even the alkanes:

SbF₃ + 2HF ⇌ H₂F⁺ + [SbF₄] ^–
Me₃CMe + superacid ⇌ [Me₃CMeH]⁺ → Me₃C⁺ + MeH

Neutralization reaction in hydrogen fluoride can be

2KF + 2HBF₄ → 2KBF₄ + 2HF
KF + SbF₃ → KSbF₄

[10.3].3.3 Metathetical Reactions

Sulphates and periodates are sufficiently stable in hydrogen fluoride medium and precipitate metal fluorides when these solutions are added to metal fluoride solution:

KIO₄ + AgF → AgIO₄ (↓) + KF
Na₂SO₄ + CoF₂ → CoSO₄ (↓) + 2NaF
NaClO₄ + TlF → TlClO₄ (↓) + NaF

[10.3].3.4 Solvolysis Reactions

Formation of covalent halides in anhydrous hydrogen fluoride to form metal fluorides can be termed as solvolysis reactions:

TiCl₄ + 4HF → TiF₄ + 4HCl

[10.3]4 ACETIC ACID

Though acetic acid is not like water, its melting point (289.8 K) , boiling point (391 K), stability and non toxic nature make it a convenient solvent. Anhydrous acetic acid is highly hygroscopic. It has low dielectric constant (7.14) and zero dipole moment due to dimerization in liquid state make it a poor solvent.

[10.3].4.1 Solubility of Salts: Nitrates of Li⁺, NH₄⁺ and Ca²⁺, acetates of LI⁺, K⁺, NH₄⁺, Cd²⁺ and Pb²⁺, chlorides of Ca²⁺, Zn²⁺, Sb³⁺ and Fe³⁺, BaI₂, ZnI₂, NH₄CNS, KCNS are extremely soluble. Sparingly soluble salts include AgNO₃, AlCl₃, HgCl₂, HgI₂, CoCl₂; while halides of Na⁺, K⁺ and NH₄⁺, KNO₃, NaNO₃, BaCl₂, KClO_3,(NH₄)₂SO₄ are slightly soluble. Silver and lead halides, CdI₂, phosphates and CaCO₃ are insoluble in acetic acid medium.

[10.3].4.2 Acid – Base Reactions: Due to the autoionization of acetic acid, protonic acids act as acids whereas metal acetates function as bases:

2CH₃COOH ⇌ CH₃COOH₂⁺ + CH₃COO^–

Typical neutralization reaction in acetic acid is

HCl + CH₃COONa → NaCl + CH₃COOH

Amphoteric behavior is shown by zinc(II), Aluminium(III), tin(II), etc:

ZnCl₂ + 2CH₃COONa ⇌ Zn(OCOCH₃)₂(↓) + 2NaCl
Zn(OCOCH₃)₂(↓) + 2CH₃COONa ⇌ Na₂[Zn(OCOCH₃)₄] (soluble)

[10.3].4.3 Metathetical Reactions: Large number of metathetical reactions proceed in acetic acid medium due to large variations in solubility of compounds.

2AgNO₃  + ZnI₂ → 2AgI (↓) + Zn(NO₃)₂
BaI₂ + 2NaNO₃ → Ba(NO₃)₂ (↓) + 2NaI

[10.3].4.4 Complex Formation Reactions: These are similar to those in aqueous medium:

Fe³⁺ + CNS ^– ⇌ FeCNS²⁺ (deep red color), etc.
Co(OCOCH₃)₂ + 4NH₄CNS ⇌
(NH₄)₂[Co(CNS)₄] (deep blue color) + 2CH₃COONH₄₌

[10.3].5 HYDROGEN CYANIDE

[10.3].5.1 Solubility: Due to high boiling point (299 K), low molecular weight, high dielectric constant (106.8), high dipole moment (2.93 D) and high hydrogen bonding, hydrogen cyanide should be as good solvent as water. However, HCN is far inferior to water for ionic compounds, and dissolves mainly the covalent compounds. Only a few metal salts (NaCl, KCl, NH₄Cl, KI, NaNO₃, K­_2­SO­₄) and covalent chlorides (HCl, POCl₃, SOCl₂, BiCl₃, SbCl₃) dissolve and ionize in HCN medium. Covalent compounds do not ionize, examples include SnX₄ (X = Cl, Br, I), I₂, benzene, methanol, urea and aniline.

[10.3].5.2 Acid – Base Reactions: Due to the autoionization of hydrogen cyanide, protonic acids (HCl, HNO₃, H₂SO₄, ClCH₂COOH) act as acids whereas metal cyanides function as bases:

2HCN ⇌ H₂CN⁺ + CN^–

Typical neutralization reaction in hydrogen cyanide medium is

HCl + NaCN → NaCl + HCN

Amphoteric behavior is shown by zinc(II), aluminium(III), tin(II), etc:

SnI₄ + 4NaCN ⇌ Sn(CN)₄(↓) + 4NaI
Sn(CN)₄ (↓) + 2NaCN ⇌ Na₂[Sn(CN)₆] (soluble)

[10.3].5.3 Metathetical Reactions: Due to limited solubility of ionic salts, not many metathetical reactions are possible. Some examples are:

2NaCl + K₂SO₄ → Na₂SO₄ (↓) + 2KCl
AgI + NaCl → NaI (↓) + AgCl

[10.3].5.4 Complex Formation Reactions: Formation of cyanide complexes can be regarded as solvolysis reactions or acid – base reactions, not many complex formation reactions are known because cyanide is a very strong ligand for metal ions.

[10.3].5.5 Solvolysis Reactions: A typical solvolysis reaction in hydrogen cyanide medium is precipitation of metal cyanides on dissolving metal salts in hydrogen cyanide:

Ag₂SO₄ + 2HCN → 2AgCN + H₂SO₄
CuCl₂ + 2HCN → Cu(CN)₂ + 2HCl

[10.3].6 HYDROGEN SULPHIDE

[10.3].6.1 General: Though sulphur is the next element to oxygen in periodic table, its hydride, hydrogen sulphide (H₂S) is not a good solvent. This is because of its narrow liquid range (187.5 – 212.9 K), low dipole moment (1.1 D) and low dielectric constant (10.2 at 213 K) along with its toxic nature, which together do not make it a good or a convenient solvent. Solubility of compounds in hydrogen sulphide may be due to solvation, protonation, solvolysis or chemical reaction.

[10.3].6.2 Acid – Base reactions: Because of its autoionization, hydrogensulphides ionizing to HS ^– ions act as strong bases in hydrogen sulphide, and can be titrated against HCl, HBr, etc. conductometrically.

2H₂S ⇌ H₃S⁺ + HS^–
[Pt(NH₃)]⁺HS^–  + HCl → [Pt(NH₃)]⁺Cl^– + H₂S
[Et₂NH]⁺HS^–  + HBr → [Et₂NH]⁺Br^–  + H₂S

[10.3].6.3 Metathetical Reactions: A few of metathetical reactions in hydrogen sulphide are:

2[Et₂NH]⁺HS^– + SnCl₄ ⇌ SnS₂ (↓) + 4[Et₂NH]⁺Cl^–  + 2H₂S
2[Et₂NH]⁺HS^–  + HgCl₂ →  Hg(SH)₂ (↓) + 2[Et₂NH]⁺Cl^–

[10.3].6.4 Solvolysis Reactions: Many covalent chlorides undergo solvolysis in hydrogen sulphide to form a hydrogensulphide, sulphide, or a thio hydrolyzed chloro compound:

2AgNO₃ + H₂S → Ag₂S (↓) + 2HNO₃
2AsCl₃ + 3H₂S → As₂S₃ (↓) + 6HCl
2PCl₃ + 3H₂S → P₂S₃  + 6HClHgCl₂ + H₂S → HgS (↓) + 2HCl
Hg₂Cl₂ + H₂S → Hg(SH)₂ + 2HCl
PCl₅ + H₂S → PSCl₃ + 2HCl
SbCl₅ + H₂S → SbSCl₃ + 2HCl

[10.3].6.5 Redox Reactions: Highly electropositive alkali metals and copper can reduce hydrogen sulphide evolving hydrogen and forming sulphides:

Cu + H₂S → CuS + H₂

Hydrogen sulphide can be oxidized by SO₃ or H_2­S₂O₇ to sulphur dioxide and other complicated products:

SO₃ + H₂S  →  H­₂O + SO₂ + S
H₂S + nS → H₂S_(n+1) (n = 2 to 6)

Sulphur dissolves in hydrogen sulphide to form polysulphides as shown above.