What you need to know about the spirit of niter
The pure compound is colorless, but older samples tend to acquire a yellow cast due to decomposition into oxides of nitrogen and water. Most commercially available nitric acid has a concentration of 68% in water. When the solution contains more than 86% HNO3, it is referred to as fuming nitric acid. Depending on the amount of nitrogen dioxide present, fuming nitric acid is further characterized as white fuming nitric acid or red fuming nitric acid, at concentrations above 95%.
Nitric acid is the primary reagent used for nitration – the addition of a nitro group, typically to an organic molecule. While some resulting nitro compounds are shock- and thermally-sensitive explosives, a few are stable enough to be used in munitions and demolition, while others are still more stable and used as pigments in inks and dyes. Nitric acid is also commonly used as a strong oxidizing agent.
Physical and chemical properties
Commercially
available nitric acid is an azeotrope with water at a concentration of 68%
HNO3, which is the ordinary concentrated nitric acid of commerce.
This solution
has a boiling temperature of 120.5 °C at 1 atm. Two solid hydrates are known;
the monohydrate (HNO3·H2O) and the trihydrate (HNO3·3H2O).
Nitric acid
70%
Nitric acid
of commercial interest usually consists of the maximum boiling azeotrope of
nitric acid and water, which is approximately 68% HNO3, (approx. 15 molar).
This is considered concentrated or technical grade, while reagent grades are
specified at 70% HNO3. The density of concentrated nitric acid is 1.42
g/cm3[inconsistent]. An older density scale is occasionally seen, with
concentrated nitric acid specified as 42° Baumé.
Contamination
with nitrogen dioxide
Fuming
nitric acid contaminated with yellow nitrogen dioxide.
Nitric acid
is subject to thermal or light decomposition and for this reason it was often
stored in brown glass bottles: 4 HNO3 → 2 H2O + 4 NO2 + O2.
This reaction may
give rise to some non-negligible variations in the vapor pressure above the
liquid because the nitrogen oxides produced dissolve partly or completely in
the acid.
The nitrogen
dioxide (NO2) remains dissolved in the nitric acid coloring it yellow or even
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 vapors, leading to the common name "red fuming acid" or
"fuming nitric acid" – the most concentrated form of nitric acid at
Standard Temperature and Pressure (STP). Nitrogen oxides (NOx) are soluble in
nitric acid.
Fuming
nitric acid
A commercial
grade of fuming nitric acid contains 90% HNO3 and has a density of 1.50 g/cm3.
This grade is often used in the explosives industry. It is not as volatile nor
as corrosive as the anhydrous acid and has the approximate concentration of
21.4 molar.
Red fuming
nitric acid, or RFNA, contains substantial quantities of dissolved nitrogen
dioxide (NO2) leaving the solution with a reddish-brown color. Due to the
dissolved nitrogen dioxide, the density of red fuming nitric acid is lower at
1.490 g/cm3.
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.
Anhydrous
nitric acid
White fuming
nitric acid, pure nitric acid or WFNA, is very close to anhydrous nitric acid.
It is available as 99.9% nitric acid by assay. One specification for white fuming
nitric acid is that it has a maximum of 2% water and a maximum of 0.5%
dissolved NO2. Anhydrous nitric acid has a density of 1.513 g/cm3 and has the
approximate concentration of 24 molar. Anhydrous nitric acid is a colorless
mobile liquid with a density of 1.512 g/cm3, which solidifies at −42 °C to form
white crystals. As it decomposes to NO2 and water, it obtains a yellow tint. It
boils at 83 °C. It is usually stored in a glass shatterproof amber bottle with
twice the volume of head space to allow for pressure build up. When received,
the pressure must be released and repeated monthly until finished.
Structure
and bonding
Two major
resonance representations of HNO3
The molecule
is planar. Two of the N–O bonds are equivalent and relatively short (this can
be explained by theories of resonance; the canonical forms show double-bond
character in these two bonds, causing them to be shorter than typical N–O
bonds), and the third N–O bond is elongated because the O atom is also attached
to a proton.
Reactions.
Nitric acid
is normally considered to be a strong acid at ambient temperatures. There is
some disagreement over the value of the acid dissociation constant, though the
pKa value is usually reported as less than −1. This means that the nitric acid
in diluted solution is fully dissociated except in extremely acidic solutions.
The pKa value rises to 1 at a temperature of 250 °C.
Nitric acid
can act as a base with respect to an acid such as sulfuric acid:
HNO3 + 2
H2SO4 ⇌ NO2+ + H3O+ + 2HSO4−; Equillibrium constant: K ~ 22
The
nitronium ion, NO2+, is the active reagent in aromatic nitration reactions.
Since nitric acid has both acidic and basic properties, it can undergo an
autoprotolysis reaction, similar to the self-ionization of water:
2 HNO3 ⇌ NO2+ + NO3− + H2O
Reactions
with metals
Nitric acid
reacts with most metals, but the details depend on the concentration of the
acid and the nature of the metal. Dilute nitric acid behaves as a typical acid
in its reaction with most metals. Magnesium, manganese and zinc liberate H2:
Mg + 2 HNO3
→ Mg(NO3)2 + H2 (Magnesium nitrate)
Mn + 2 HNO3
→ Mn(NO3)2 + H2 (Manganese nitrate)
Zn + 2 HNO3
→ Zn(NO3)2 + H2 (Zinc nitrate)
Nitric acid
can oxidize non-active metals such as copper and silver. With these non-active
or less electropositive metals the products depend on temperature and the acid
concentration. For example, copper reacts with dilute nitric acid at ambient
temperatures with a 3:8 stoichiometry:
3 Cu + 8
HNO3 → 3 Cu2+ + 2 NO + 4 H2O + 6 NO3−
The nitric
oxide produced may react with atmospheric oxygen to give nitrogen dioxide. With
more concentrated nitric acid, nitrogen dioxide is produced directly in a
reaction with 1:4 stoichiometry:
Cu + 4 H+ +
2 NO3− → Cu2+ + 2 NO2 + 2 H2O
Upon
reaction with nitric acid, most metals give the corresponding nitrates. Some
metalloids and metals give the oxides; for instance, Sn, As, Sb, and Ti are
oxidized into SnO2, As2O5, Sb2O5, and TiO2 respectively.
Some
precious metals, such as pure gold and platinum-group metals do not react with
nitric acid, though pure gold does react with aqua regia, a mixture of
concentrated nitric acid and hydrochloric acid. However, some less noble metals
(Ag, Cu, ...) present in some gold alloys relatively poor in gold such as
colored gold can be easily oxidized and dissolved by nitric acid, leading to
colour changes of the gold-alloy surface. Nitric acid is used as a cheap means
in jewelry shops to quickly spot low-gold alloys (< 14 carats) and to
rapidly assess the gold purity.
Being a
powerful oxidizing agent, nitric acid reacts violently with many non-metallic
compounds, and the reactions may be explosive. Depending on the acid
concentration, temperature and the reducing agent involved, the end products
can be variable. Reaction takes place with all metals except the noble metals
series and certain alloys. As a general rule, oxidizing reactions occur
primarily with the concentrated acid, favoring the formation of nitrogen
dioxide (NO2). However, the powerful oxidizing properties of nitric acid are
thermodynamic in nature, but sometimes its oxidation reactions are rather
kinetically non-favored. The presence of small amounts of nitrous acid (HNO2)
greatly enhance the rate of reaction.
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 bulk of
the metal from further oxidation. The formation of this protective layer is
called passivation. Typical passivation concentrations range from 20% to 50% by
volume (see ASTM A967-05). Metals that are passivated by concentrated nitric
acid are iron, cobalt, chromium, nickel, and aluminium.
Reactions
with non-metals[edit]
Being a
powerful oxidizing acid, nitric acid reacts violently with many organic
materials and the reactions may be explosive. The hydroxyl group will typically
strip a hydrogen from the organic molecule to form water, and the remaining
nitro group takes the hydrogen's place. Nitration of organic compounds with
nitric acid is the primary method of synthesis of many common explosives, such
as nitroglycerin and trinitrotoluene (TNT). As very many less stable byproducts
are possible, these reactions must be carefully thermally controlled, and the
byproducts removed to isolate the desired product.
Reaction
with non-metallic elements, with the exceptions of nitrogen, oxygen, noble
gases, silicon, and halogens other than iodine, usually oxidizes them to their
highest oxidation states as acids with the formation of nitrogen dioxide for
concentrated acid and nitric oxide for dilute acid.
C + 4 HNO3 →
CO2 + 4 NO2 + 2 H2O
or
3 C + 4 HNO3
→ 3 CO2 + 4 NO + 2 H2O
Concentrated
nitric acid oxidizes I2, P4, and S8 into HIO3, H3PO4, and H2SO4,
respectively.[8]
Xanthoproteic
test
Nitric acid
reacts with proteins to form yellow nitrated products. This reaction is known
as the xanthoproteic reaction. This test is carried out by adding concentrated
nitric acid to the substance being tested, and then heating the mixture. If
proteins that contain amino acids with aromatic rings are present, the mixture
turns yellow. Upon adding a base such as ammonia, the color turns orange. These
color changes are caused by nitrated aromatic rings in the protein.
Xanthoproteic acid is formed when the acid contacts epithelial cells.
Respective local skin color changes are indicative of inadequate safety
precautions when handling nitric acid.
Production
Nitric acid
is made by reaction of nitrogen dioxide (NO2) with water.
3 NO2 + H2O
→ 2 HNO3 + NO
Normally,
the nitric oxide produced by the reaction is reoxidized by the oxygen in air to
produce additional nitrogen dioxide.
Bubbling
nitrogen dioxide through hydrogen peroxide can help to improve acid yield.
2 NO2 + H2O2
→ 2 HNO3
Commercial
grade nitric acid solutions are usually between 52% and 68% nitric acid.
Production of nitric acid is via the Ostwald process, named after German
chemist Wilhelm Ostwald. In this process, anhydrous ammonia is oxidized to
nitric oxide, in the presence of platinum or rhodium gauze catalyst at a high
temperature of about 500 K and a pressure of 9 bar.
4 NH3 (g) +
5 O2 (g) → 4 NO (g) + 6 H2O (g) (ΔH = −905.2 kJ)
Nitric oxide
is then reacted with oxygen in air to form nitrogen dioxide.
2 NO (g) +
O2 (g) → 2 NO2 (g) (ΔH = −114 kJ/mol)
This is
subsequently absorbed in water to form nitric acid and nitric oxide.
3 NO2 (g) +
H2O (l) → 2 HNO3 (aq) + NO (g) (ΔH = −117 kJ/mol)
The nitric
oxide is cycled back for reoxidation. Alternatively, if the last step is
carried out in air:
4 NO2 (g) +
O2 (g) + 2 H2O (l) → 4 HNO3 (aq)
The aqueous
HNO3 obtained can be concentrated by distillation up to about 68% by mass.
Further concentration to 98% can be achieved by dehydration with concentrated
H2SO4. By using ammonia derived from the Haber process, the final product can
be produced from nitrogen, hydrogen, and oxygen which are derived from air and
natural gas as the sole feedstocks.
Prior to the
introduction of the Haber process for the production of ammonia in 1913, nitric
acid was produced using the Birkeland–Eyde process, also known as the arc
process. This process is based upon the oxidation of atmospheric nitrogen by
atmospheric oxygen to nitric oxide at very high temperatures. An electric arc
was used to provide the high temperatures, and yields of up to 4% nitric oxide
were obtained. The nitric oxide was cooled and oxidized by the remaining
atmospheric oxygen to nitrogen dioxide, and this was subsequently absorbed in
dilute nitric acid. The process was very energy intensive and was rapidly
displaced by the Ostwald process once cheap ammonia became available.
Laboratory synthesis
In
laboratory, nitric acid can be made by thermal decomposition of copper (II)
nitrate, producing nitrogen dioxide and oxygen gases, which are then passed
through water to give nitric acid.
2 Cu(NO3)2 →
2 CuO (s) + 4 NO2 (g) + O2 (g)
An alternate
route is by reaction of approximately equal masses of any nitrate salt such as
sodium nitrate with 96% sulfuric acid (H2SO4), and distilling this mixture at
nitric acid's boiling point of 83 °C. A nonvolatile residue of the metal
sulfate remains in the distillation vessel. The red fuming nitric acid obtained
may be converted to the white nitric acid.
2 NaNO3 +
H2SO4 → 2 HNO3 + Na2SO4
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.
Dilute
nitric acid may be concentrated by distillation up to 68% acid, which is a
maximum boiling azeotrope containing 32% water. In the laboratory, further
concentration involves distillation with either sulfuric acid or magnesium
nitrate which act as dehydrating agents. Such distillations must be done with
all-glass apparatus at reduced pressure, to prevent decomposition of the acid.
Industrially, highly concentrated nitric acid is produced by dissolving
additional nitrogen dioxide in 68% nitric acid in an absorption tower.[12]
Dissolved nitrogen oxides are either stripped in the case of white fuming
nitric acid, or remain in solution to form red fuming nitric acid. More
recently, electrochemical means have been developed to produce anhydrous acid
from concentrated nitric acid feedstock.
Nitric acid in a laboratory.
The main
industrial use of nitric acid is for the production of fertilizers. Nitric acid
is neutralized with ammonia to give ammonium nitrate. This application consumes
75–80% of the 26M tons produced annually (1987). The other main applications
are for the production of explosives, nylon precursors, and specialty organic
compounds.
Precursor to organic nitrogen
compounds
In organic
synthesis, industrial and otherwise, the nitro group is a versatile functional
group. Most derivatives of aniline are prepared via nitration of aromatic
compounds followed by reduction. Nitrations entail combining nitric and
sulfuric acids to generate the nitronium ion, which electrophilically reacts
with aromatic compounds such as benzene. Many explosives, such as TNT, are
prepared this way.
Use as an oxidant
The
precursor to nylon, adipic acid, is produced on a large scale by oxidation of
cyclohexanone and cyclohexanol with nitric acid.
Rocket propellant
Nitric acid
has been used in various forms as the oxidizer in liquid-fueled rockets. These
forms include red fuming nitric acid, white fuming nitric acid, mixtures with
sulfuric acid, and these forms with HF inhibitor.[15] IRFNA (inhibited red
fuming nitric acid) was one of 3 liquid fuel components for the BOMARC
missile.[16]
Niche uses
Analytical reagent
In elemental
analysis by ICP-MS, ICP-AES, GFAA, and Flame AA, dilute nitric acid (0.5 to
5.0%) is used as a matrix compound for determining metal traces in
solutions.[17] Ultrapure trace metal grade acid is required for such
determination, because small amounts of metal ions could affect the result of
the analysis.
It is also
typically used in the digestion process of turbid water samples, sludge
samples, solid samples as well as other types of unique samples which require
elemental analysis via ICP-MS, ICP-OES, ICP-AES, GFAA and flame atomic
absorption spectroscopy. Typically these digestions use a 50% solution of the
purchased HNO
3 mixed with Type 1 DI Water.
In
electrochemistry, nitric acid is used as a chemical doping agent for organic
semiconductors, and in purification processes for raw carbon nanotubes.
Woodworking
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).
Etchant and cleaning agent
The
corrosive effects of nitric acid are exploited for a number of specialty
applications, such as pickling stainless steel or cleaning silicon wafers in
electronics.
A solution
of nitric acid, water and alcohol, Nital, is used for etching of metals to
reveal the microstructure. ISO 14104 is one of the standards detailing this
well known procedure.
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).
The phosphoric acid content helps to passivate ferrous alloys against corrosion
by the dilute nitric acid.[citation needed]
Nitric acid
can be used as a spot test for alkaloids like LSD, giving a variety of colours
depending on the alkaloid.
Safety
Second
degree burn caused by nitric acid
Nitric acid
is a corrosive acid and a powerful oxidizing agent. The major hazard posed by
it is chemical burns as it carries out acid hydrolysis with proteins (amide)
and fats (ester) which consequently decomposes living tissue (e.g. skin and
flesh). Concentrated nitric acid stains human skin yellow due to its reaction
with the keratin. These yellow stains turn orange when neutralized. Systemic
effects are unlikely, however, and the substance is not considered a carcinogen
or mutagen.
The standard
first aid treatment for acid spills on the skin is, as for other corrosive
agents, irrigation with large quantities of water. Washing is continued for at
least ten to fifteen minutes to cool the tissue surrounding the acid burn and
to prevent secondary damage. Contaminated clothing is removed immediately and
the underlying skin washed thoroughly.
Being a
strong oxidizing agent, reactions of nitric acid with compounds such as
cyanides, carbides, metallic powders can be explosive and those with many
organic compounds, such as turpentine, are violent and hypergolic (i.e.
self-igniting). Hence, it should be stored away from bases and organics.
No comments