Nitric acid (HNO₃), also known as aqua fortis and spirit of nitre, is a highly corrosive and toxic strong acid that can cause severe burns. The synthesis of nitric acid was first recorded circa 800 AD by the Muslim alchemist Jabir ibn Hayyan.

Colorless when pure, older samples tend to acquire a yellow cast due to the accumulation of oxides of nitrogen. If the solution contains more than 86% nitric acid, it is referred to as fuming nitric acid. Fuming nitric acid is characterized as white fuming nitric acid and red fuming nitric acid, depending on the amount of nitrogen dioxide present.

Properties

Pure anhydrous nitric acid (100%) is a colorless liquid with a density of 1522 kg/m³ which solidifies at -42 °C to form white crystals and boils at 83 °C. When boiling in light, even at room temperature, there is a partial decomposition with the formation of nitrogen dioxide following the reaction:

4HNO3 2H2O + 4NO2 + O2 (72°C)

which means that anhydrous nitric acid should be stored below 0 °C to avoid decomposition. The nitrogen dioxide (NO2) remains dissolved in the nitric acid coloring it yellow, or red at higher temperatures. While the pure acid tends to give off white fumes when exposed to air, acid with dissolved nitrogen dioxide gives off reddish-brown vapours, leading to the common name "red fuming acid" or "fuming nitric acid".

Nitric acid is miscible with water in all proportions and distillation gives an azeotrope with a concentration of 68% HNO₃ and a boiling temperature of 120.5 °C at 1 atm. Two solid hydrates are known; the monohydrate (HNO3•H2O) and the trihydrate (HNO₃•3H₂O).

Nitrogen oxides (NOx) are soluble in nitric acid and this property influences more or less, all the physical characteristics depending on the concentration of the oxides. These mainly include the vapor pressure above the liquid and the boiling temperature, as well as the color mentioned above.

Nitric acid is subject to thermal or light decomposition with increasing concentration and this may give rise to some non-negligible variations in the vapour pressure above the liquid because the nitrogen oxides produced dissolve partly or completely in the acid.

Acidic properties

Being a typical acid, nitric acid reacts with alkalis, basic oxides, and carbonates to form salts, such as ammonium nitrate. Due to its oxidizing nature, nitric acid generally does not liberate hydrogen on reaction with metals and the resulting salts are usually in the higher oxidized states. For this reason, heavy corrosion can be expected and should be guarded against by the appropriate use of corrosion resistant metals or alloys.

Nitric acid has an acid dissociation constant (pKa) of −1.4: in aqueous solution, it almost completely (93% at 0.1 mol/L) ionizes into the nitrate ion NO₃^- and a hydrated proton, known as a hydronium ion, H₃O⁺.

HNO₃ + H₂O H₃O⁺ + NO₃^-

Oxidizing properties

Reactions with metals

Being a powerful oxidizing agent, nitric acid reacts violently with many organic materials and the reactions may be explosive. Depending on the acid concentration, temperature and the reducing agent involved, the end products can be variable. Reaction then takes place with all metals except the precious metal series and certain alloys. As a general rule of course, oxidizing reactions occur primarily with the concentrated acid, favouring the formation of nitrogen dioxide (NO₂).

 Cu + 4HNO₃ Cu(NO₃)2 + 2NO₃ + 2H₂O

The acidic properties tend to dominate with dilute acid, coupled with the preferential formation of nitrogen oxide (NO).

 3Cu + 8HNO₃ 3Cu(NO₃)2 + 2NO + 4H₂O

Since nitric acid is an oxidizing agent, hydrogen (H₂) is rarely formed. Only magnesium (Mg), Manganese (Mn) and calcium (Ca) react with cold, dilute nitric acid to give hydrogen:

Mg_(s) + 2HNO_{3 (aq)} Mg(NO₃)2 _(aq) + H_{2 (g)}

Passivation

Although chromium (Cr), iron (Fe) and aluminium (Al) readily dissolve in dilute nitric acid, the concentrated acid forms a metal oxide layer that protects the metal from further oxidation, which is called passivation.

Reactions with non-metals

Reaction with non-metallic elements, with the exception of silicon and halogens, usually oxidizes them to their highest oxidation states as acids with the formation of nitrogen dioxide for concentrated acid and nitrogen oxide for dilute acid.

C + 4HNO₃ CO₂ + 4NO₂ + 2 H₂O
or
3C + 4HNO₃ 3CO₂ + 4NO + 2 H₂O

Grades

White fuming nitric acid, also called 100% nitric acid or WFNA, is very close to the anhydrous nitric acid product. One specification for white fuming nitric acid is that it has a maximum of 2% water and a maximum of 0.5% dissolved NO₂.

Red fuming nitric acid, or RFNA, contains substantial quantities of dissolved nitrogen dioxide (NO₂) leaving the solution with a reddish-brown color. One formulation of RFNA specifies a minimum of 17% NO₂, another specifies 13% NO₂.

An inhibited fuming nitric acid (either IWFNA, or IRFNA) can be made by the addition of 0.6 to 0.7% hydrogen fluoride, HF. This fluoride is added for corrosion resistance in metal tanks (the fluoride creates a metal fluoride layer that protects the metal).

Industrial production

Nitric acid is made by mixing nitrogen dioxide (NO₂) with water in the presence of oxygen or air to oxidize the nitrous acid also produced by the reaction.

Dilute nitric acid may be concentrated by distillation up to 68% acid, which is an azeotropic mixture with 32% water. Further concentration involves distillation with sulfuric acid which acts as a dehydrating agent. In the laboratory, such distillations must be done with all-glass apparatus at reduced pressure, to prevent decomposition of the acid.

Commercial grade nitric acid solutions are usually between 52% and 68% nitric acid. Commercial production of nitric acid is via the Ostwald process, named after Wilhelm Ostwald.

The acid can also be synthesized by oxidizing ammonia, but the product is diluted by the water also formed as part of the reaction. However, this method is important in producing ammonium nitrate from ammonia derived from the Haber process, because the final product can be produced from nitrogen, hydrogen, and oxygen as the sole feedstocks.

Laboratory synthesis

In laboratory, nitric acid can be made from copper(II) nitrate or by reacting approximately equal masses of potassium nitrate (KNO₃) with 96% sulfuric acid (H₂SO₄), and distilling this mixture at nitric acid's boiling point of 83 °C until only a white crystalline mass, potassium hydrogen sulfate (KHSO₄), remains in the reaction vessel. The obtained red fuming nitric acid may be converted to the white nitric acid.

H2SO₄ + KNO₃ KHSO₄ + HNO₃

The dissolved NOx are readily removed using reduced pressure at room temperature (10-30 min at 200 mmHg or 27 kPa) to give white fuming nitric acid. This procedure can also be performed under reduced pressure and temperature in one step in order to produce less nitrogen dioxide gas.[citation needed]

Uses

Nitric acid in a laboratory.

IWFNA may be used as the oxidizer in liquid fuel rockets IRFNA was one of 3 liquid fuel components for the BOMARC missile. A solution of nitric acid and alcohol, Nital, is used for etching of metals to reveal the microstructure.

Commercially available aqueous blends of 5-30% nitric acid and 15-40% phosphoric acid are commonly used for cleaning food and dairy equipment primarily to remove precipitated calcium and magnesium compounds (either deposited from the process stream or resulting from the use of hard water during production and cleaning).

Nitration

Nitric acid is used in the manufacture of nitrate-containing explosives such as nitroglycerin, trinitrotoluene (TNT) and cyclotrimethylenetrinitramine (RDX), as well as fertilizers such as ammonium nitrate.

Digestion

In elemental analysis by ICP-MS and ICP-AES, dilute nitric acid (0.5 to 2.0 %) is used as a matrix compound for determining metal traces in solutions.[citation needed] Ultrapure acid is required for such determination, because small amounts of metal ions could affect the result of the analysis.

Wood working

In a low concentration (approximately 10%), nitric acid is often used to artificially age pine and maple. The color produced is a grey-gold very much like very old wax or oil finished wood (wood finishing).

Other uses

Alone, it is useful in metallurgy and refining as it reacts with most metals, and in organic syntheses. When mixed with hydrochloric acid, nitric acid forms aqua regia, one of the few reagents capable of dissolving gold and platinum

2 FeS₂ + 7 O₂ + 2 H₂O 2 Fe²⁺ + 4 SO₄^2- + 4 H⁺

The Fe²⁺ can be further oxidized to Fe³⁺, according to:

4 Fe²⁺ + O₂ + 4 H⁺ 4 Fe³⁺ + 2 H₂O

and the Fe3+ produced can be precipitated as the hydroxide or hydrous oxide. The equation for the formation of the hydroxide is

Fe³⁺ + 3 H₂O Fe(OH)₃ + 3 H⁺

The ironion ("ferric iron", in casual nomenclature) can also oxidize pyrite. When iron oxidation of pyrite occurs, the process can become rapid. pH values below zero have been measured in AMD produced by this process.

AMD can also produce Sulphuric acid at a slower rate, so that the Acid Neutralization Capacity (ANC) of the aquifer can neutralize the produced acid. In such cases, the Total Dissolved solids (TDS) concentration of the water can be increased form the dissolution of minerals from the acid-neutralization reaction with the minerals.

Extraterrestrial Sulphuric acid

Sulphuric acid is produced in the upper atmosphere of Venus by the sun's photochemical action on carbon dioxide, sulfur dioxide, and water vapor. Ultraviolet photons of wavelengths less than 169 nm can photodissociate carbon dioxide into carbon monoxide and atomic oxygen. Atomic oxygen is highly reactive. When it reacts with sulfur dioxide, a trace component of the Venerian atmosphere, the result is sulfur trioxide, which can combine with water vapor, another trace component of Venus's atmosphere, to yield Sulphuric acid.
CO₂ CO + O
SO₂ + O SO₃
SO₃ + H₂O H₂SO₄

In the upper, cooler portions of Venus's atmosphere, Sulphuric acid exists as a liquid, and thick Sulphuric acid clouds completely obscure the planet's surface when viewed from above. The main cloud layer extends from 45–70 km above the planet's surface, with thinner hazes extending as low as 30 and as high as 90 km above the surface.

Infrared spectra from NASA's Galileo mission show distinct absorptions on Jupiter's moon Europa that have been attributed to one or more Sulphuric acid hydrates. The interpretation of the spectra is somewhat controversial. Some planetary scientists prefer to assign the spectral features to the sulfate ion, perhaps as part of one or more minerals on Europa's surface.

Manufacture

Sulphuric acid is produced from sulfur, oxygen and water via the contact process.

In the first step, sulfur is burned to produce sulfur dioxide.

(1) S(s) + O2(g) SO2(g)

This is then oxidised to sulfur trioxide using oxygen in the presence of a vanadium(V) oxide catalyst.

(2) 2 SO₂ + O2(g) 2 SO₃(g) (in presence of V2O5)

Finally the sulfur trioxide is treated with water (usually as 97-98% H2SO4 containing 2-3% water) to produce 98-99% Sulphuric acid.

(3) SO₃(g) + H₂O(l) H₂SO₄(l)

Note that directly dissolving SO3 in water is not practical due to the highly exothermic nature of the reaction, forming a corrosive mist instead of a liquid. Alternatively, SO3 can be absorbed into H2SO4 to produce oleum (H2S2O7), which may then be mixed with water to form Sulphuric acid.

(3) H₂SO₄(l) + SO₃ H₂S₂O₇(l)

Oleum is reacted with water to form concentrated H2SO4.

(4) H₂S₂O₇(l) + H₂O(l) 2 H₂SO₄(l)

Physical properties

Forms of Sulphuric acid

Although nearly 100% Sulphuric acid can be made, this loses SO3 at the boiling point to produce 98.3% acid. The 98% grade (18M) is more stable in storage, and is the usual form of what is described as concentrated Sulphuric acid. Other concentrations are used for different purposes.

Some common concentrations are

10%, dilute Sulphuric acid for laboratory use,
33.5%, battery acid (used in lead-acid batteries),
62.18%, chamber or fertilizer acid,
77.67%, tower or Glover acid,
98%, concentrated acid.

Different purities are also available. Technical grade H₂SO₄ is impure and often colored, but is suitable for making fertilizer. Pure grades such as United States Pharmacopoeia (USP) grade are used for making pharmaceuticals and dyestuffs.

When high concentrations of SO₃(g) are added to Sulphuric acid, H₂S₂O₇, called pyroSulphuric acid, fuming Sulphuric acid or oleum or, less commonly, Nordhausen acid, is formed. Concentrations of oleum are either expressed in terms of% SO₃ (called% oleum) or as% H₂SO₄ (the amount made if H₂O were added); common concentrations are 40% oleum (109% H₂SO₄) and 65% oleum (114.6% H₂SO₄). Pure H₂S₂O₇ is a solid with melting point 36°C.

Polarity and conductivity

Anhydrous H2SO4 is a very polar liquid, having a dielectric constant of around 100. It has a high electrical conductivity, caused by dissociation through protonating itself, a process known as autoprotolysis, or autoionization.

2 H₂SO₄ H₃SO₄⁺ + HSO₄^-

The equilibrium constant for the autoprotolysis is

Kap(25°C)= [H₃SO₄⁺][HSSO₄^- ] = 2.7 × 10−4.

The comparable equilibrium constant for water, Kw is 10−14, a factor of 1010 (10 billion) smaller.

In spite of the viscosity of the acid, the effective conductivities of the H₃SO₄⁺ and HSO₄^- ions are high due to an intra-molecular proton-switch mechanism (analogous to the Grotthuss mechanism in water), making Sulphuric acid a good conductor. It is also an excellent solvent for many reactions.

The equilibrium is actually more complex than shown above; 100% H₃SO₄ contains the following species at equilibrium (figures shown as millimol per kg solvent): HSO₄^- (15.0), H₃SO₄⁺ (11.3), H₃O+ (8.0), HS₂O₇^- (4.4), H₂S₂O₇ (3.6), H₂O (0.1).

The hydration reaction of Sulphuric acid is highly exothermic. If water is added to the concentrated Sulphuric acid, it can react, boil and spit dangerously. One should always add the acid to the water rather than the water to the acid. The necessity for this safety precaution is due to the relative densities of these two liquids. Water is less dense than Sulphuric acid, meaning water will tend to float on top of this acid. The reaction is best thought of as forming hydronium ions, by
H₂SO₄ + H₂O H₃O+ + HSO₄^-,

and then

HSO₄^- + H₂O H₃O⁺ + SO42^- .

Because the hydration of Sulphuric acid is thermodynamically favorable, Sulphuric acid is an excellent dehydrating agent. The affinity of Sulphuric acid for water is sufficiently strong that it will remove hydrogen and oxygen atoms from other compounds; for example, mixing starch (C6H12O6)n and concentrated Sulphuric acid will give elemental carbon and water which is absorbed by the Sulphuric acid (which becomes slightly diluted): (C₆H₁₂O₆)n 6C + 6H₂O. The effect of this can be seen when concentrated Sulphuric acid is spilled on paper; the cellulose reacts to give a burned appearance, the carbon appears much as soot would in a fire. A more dramatic reaction occurs when Sulphuric acid is added to a tablespoon of white sugar; a rigid column of black, porous carbon will quickly emerge. The carbon will smell strongly of caramel.

Other reactions

As an acid, Sulphuric acid reacts with most bases to give the corresponding sulfate. For example, copper sulfate. This blue salt of copper, commonly used for electroplating and as a fungicide, is prepared by the reaction of copper oxide (CuO) with Sulphuric acid:

CuO + H₂SO₄ CuSO₄+ H₂O

Sulphuric acid can also be used to displace weaker acids from their salts. Reaction with sodium acetate, for example, displaces acetic acid:

H₂SO₄+ CH₃COONa NaHSO₄ + CH₃COOH

Similarly, reacting Sulphuric acid with potassium nitrate can be used to produce nitric acid and a precipitate of potassium bisulfate. When combined with nitric acid, Sulphuric acid acts both as an acid and a dehydrating agent, forming the nitronium ion NO₂+, which is important in nitration reactions involving electrophilic aromatic substitution. This type of reaction, where protonation occurs on an oxygen atom, is important in many organic chemistry reactions, such as Fischer esterification and dehydration of alcohols.

Sulphuric acid reacts with most metals via a single displcement reaction to produce hydrogen gas and the metal sulfate. Dilute H₂SO₄attacks iron, aluminium, zinc, manganese, magnesium and nickel, but reactions with tin and copper require the acid to be hot and concentrated. Lead and tungsten, however, are resistant to Sulphuric acid. The reaction with iron (shown) is typical for most of these metals, but the reaction with tin is unusual in that it produces sulfur dioxide rather than hydrogen.

Fe(s) + H₂SO₄(aq) H₂(g) + FeSO₄(aq)
Sn(s) + 2 H₂SO₄(aq) SnSO₄(aq) + 2 H₂O(l) + SO₂(g)

Sulphuric acid undergoes electrophilic aromatic substitution with aromatic compounds to give the corresponding sulfonic acids:

Sulphuric acid production in 2000

Production trend in some countries

Sulphuric acid is a very important commodity chemical, and indeed, a nation's Sulphuric acid production is a good indicator of its industrial strength. The major use (60% of total production worldwide) for Sulphuric acid is in the "wet method" for the production of phosphoric acid, used for manufacture of phosphate fertilizers as well as trisodium phosphate for detergents. In this method, phosphate rock is used, and more than 100 million tonnes are processed annually. This raw material is shown below as fluorapatite, though the exact composition may vary. This is treated with 93% Sulphuric acid to produce calcium sulfate, hydrogen fluoride (HF) and phosphoric acid. The HF is removed as hydrofluoric acid. The overall process can be represented as:

Ca5F(PO₄)³ + 5 H₂SO₄ + 10 H₂O 5 CaSO₄•2 H₂O + HF + 3 H₃PO₄.

Sulphuric acid is used in large quantities by the iron and steelmaking industry to remove oxidation, rust and scale from rolled sheet and billets prior to sale to the automobile and white-goods industry. Used acid is often recycled using a Spent Acid Regeneration (SAR) plant. These plants combust spent acid with natural gas, refinery gas, fuel oil or other fuel sources. This combustion process produces gaseous sulfur dioxide (SO₂) and sulfur trioxide (SO₃) which are then used to manufacture "new" Sulphuric acid. SAR plants are common additions to metal smelting plants, oil refineries, and other industries where Sulphuric acid is consumed in bulk, as operating a SAR plant is much cheaper than the recurring costs of spent acid disposal and new acid purchases.

Al₂O₃+ 3 H₂SO₄ Al₂(SO₄)³ + 3 H₂O

Sulphuric acid is used for a variety of other purposes in the chemical industry. For example, it is the usual acid catalyst for the conversion of cyclohexanoneoxime to caprolactam, used for making nylon. It is used for making hydrochloric acid from salt via the Mannheim process. Much H₂SO₄ is used in petroleum refining, for example as a catalyst for the reaction of isobutane with isobutylene to give isooctane, a compound that raises the octane rating of gasoline (petrol). Sulphuric acid is also important in the manufacture of dyestuffs solutions and is the "acid" in lead-acid (car) batteries.

Sulphuric acid is also used as a general dehydrating agent in its concentrated form. See Reaction with water.

Sulfur-iodine cycle

The sulfur-iodine cycle is a series of thermo-chemical processes used to obtain hydrogen. It consists of three chemical reactions whose net reactant is water and whose net products are hydrogen and oxygen.