Nitric acid is most commonly prepared in the laboratory by distilling a mixture of a metal nitrate salt and concentrated sulfuric acid, collecting the vapors that condense into the familiar colorless-to-pale-yellow fuming liquid. The reaction itself is straightforward, but the practical details of glassware, temperature control, and above all safe handling of toxic fumes make the difference between a successful preparation and a dangerous one. Several routes exist depending on the quantity and purity required, from small-scale retort distillations to the industrial Ostwald process that supplies the world’s commercial nitric acid.
The Salt-and-Acid Distillation Method
The oldest and simplest laboratory route starts with a nitrate salt, usually potassium nitrate or sodium nitrate, and concentrated sulfuric acid. When the two are heated together, sulfuric acid donates a proton to the nitrate ion, liberating nitric acid as a vapor. Potassium nitrate is often preferred because it is widely available, stable, and leaves behind a clean bisulfate residue. Sodium nitrate works just as well chemically but tends to be slightly more hygroscopic, which can complicate storage of the starting material.
The setup for this reaction is a simple distillation. A round-bottom or retort flask holds the mixed reagents. Heat is applied gently with a heating mantle or sand bath, and the nitric acid vapor travels through a condenser cooled by running water, then drips into a collection flask. The collection flask should be kept in an ice bath to minimize losses from the vapor pressure of the product. Everything downstream of the reaction flask needs to be connected with ground-glass joints or high-quality tubing that can tolerate strong acid fumes, and the entire system must be run inside a working fume hood.
Temperature matters more than most beginners expect. The reaction begins to release nitric acid vapor at moderate heat, but if you push the temperature too high, the product starts to decompose. Research on the thermal decomposition of nitric acid vapor shows that at around 400 °C, the molecule breaks apart into nitrogen dioxide and water, with nitrogen dioxide then further breaking down to yield oxygen.1Journal of Physical Chemistry. Kinetics of the thermal decomposition of nitric acid vapor. II. Mechanism That decomposition is why the distillation should be run at the lowest temperature that gives a steady rate of vapor production. Overheating turns your product brown with dissolved nitrogen dioxide and lowers both purity and yield.
The stoichiometry calls for roughly equal molar amounts of the nitrate salt and sulfuric acid. In practice, a slight excess of sulfuric acid helps drive the reaction to completion. The crude product that collects in the receiving flask is typically a dilute nitric acid, somewhere around 60–68% concentration depending on how carefully you control conditions. Pure, anhydrous nitric acid is a much harder target and requires additional dehydration steps that go beyond the basic preparation.
The Ostwald Process at Industrial Scale
Nearly all commercial nitric acid is made by the Ostwald process, which starts not with a nitrate salt but with ammonia gas. Ammonia is catalytically oxidized over a platinum-rhodium gauze at high temperature to produce nitric oxide, which is then further oxidized in air to nitrogen dioxide, and finally absorbed in water to yield nitric acid. The process runs continuously at enormous scale and produces acid in concentrations ranging from about 50% to 68% in a single pass.
The platinum gauze catalyst is the heart of the process. It operates under harsh conditions, with temperatures around 800–900 °C and corrosive gases flowing across its surface. Over time, the gauze degrades, losing platinum to volatilization and mechanical erosion. This degradation is a well-known industrial problem, and computational modeling of the ammonia oxidation process on platinum gauze continues to be an active area of research aimed at extending catalyst life and maintaining selectivity.2Chemical Engineering Research and Design. Study of the ammonia oxidation process on platinum gauze and catalyst degradation phenomenon – CFD simulation with surface reaction kinetics and catalyst entrained particles motion and deposition tracking
For a bench chemist, the Ostwald process is not practical. It requires a source of compressed ammonia, a platinum catalyst, and engineering controls for the extremely high temperatures and gas volumes involved. But understanding it gives useful context: the commercial acid you buy in a bottle was almost certainly made this way, and its concentration and impurity profile reflect the absorption step at the end of that process.
A Historical Alternative Using Electrical Discharge
Before the Ostwald process dominated, one of the first industrial nitrogen fixation methods was the Birkeland-Eyde process, which used an electric arc to force atmospheric nitrogen and oxygen to combine directly into nitric oxide.3PubMed Central. From the Birkeland-Eyde process towards energy-efficient plasma-based NOX synthesis: a techno-economic analysis That nitric oxide was then processed similarly to the Ostwald route: further oxidation to nitrogen dioxide, followed by absorption in water. The process worked but consumed staggering amounts of electricity, making it viable only in places with cheap hydroelectric power, like early twentieth-century Norway.
Interest in plasma-based nitrogen fixation has revived somewhat in recent years, driven by the idea that renewable electricity could make the process greener than the Ostwald route, which depends on ammonia derived from natural gas. For the laboratory chemist, the Birkeland-Eyde approach is a curiosity rather than a practical option, but it illustrates that the fundamental chemistry of nitric acid synthesis, combining nitrogen and oxygen and dissolving the resulting oxides in water, can be achieved by very different energy sources.
Concentration Grades and What They Mean
Nitric acid comes in several standard concentrations, and the terminology can be confusing. “Concentrated” nitric acid typically means 68% by weight in water, which is the maximum-boiling azeotrope. You cannot get above that concentration by simple distillation of a water-nitric acid mixture because the azeotrope distills as a single composition. Going higher requires breaking the azeotrope, which is usually done by distilling the acid in the presence of a dehydrating agent such as concentrated sulfuric acid.
“Fuming” nitric acid refers to grades above the azeotrope, generally above 86%. White fuming nitric acid is nearly pure and appears colorless to very pale yellow, while red fuming nitric acid contains dissolved nitrogen dioxide that gives it a distinctive reddish-brown color and makes it an even more aggressive oxidizer. Red fuming nitric acid has historically been used as a rocket propellant oxidizer, which gives some indication of how reactive it is.
For most laboratory reactions, the standard 68% concentrate is what you want. Dilute solutions in the 10–30% range are used for metal dissolution and cleaning glassware. If your preparation method yielded a dilute product and you need higher concentration, careful distillation with a dehydrating agent is the next step, but this is significantly more hazardous than the initial preparation and should only be attempted with proper training and equipment.
Why Nitrogen Dioxide Fumes Are So Dangerous
The single biggest hazard when making or handling nitric acid is nitrogen dioxide gas. It is the brown, acrid-smelling gas that appears whenever nitric acid decomposes, overheats, or reacts with metals and organic materials. Nitrogen dioxide is insidious because it can cause serious lung damage even at concentrations that do not feel immediately unbearable.
A classic medical study of workers exposed to nitrogen dioxide in agricultural silos found that even brief exposure caused respiratory symptoms ranging from coughing and chest tightness to severe pulmonary edema. In two of four documented cases, patients died on the 27th and 30th days after exposure, following a deceptive pattern where initial symptoms seemed to stabilize before suddenly worsening.4JAMA. “SILO-FILLER’S DISEASE”—A SYNDROME CAUSED BY NITROGEN DIOXIDE The delayed worsening is caused by a condition called bronchiolitis fibrosa obliterans, where scar tissue progressively blocks the small airways over days to weeks after the initial chemical burn.
This delayed toxicity pattern is what makes nitrogen dioxide uniquely treacherous. You could inhale a dangerous dose during a nitric acid preparation, feel only mild irritation, and assume you are fine, only to develop life-threatening lung damage days later. A properly functioning fume hood is not optional. If you do not have one, do not attempt to make or concentrate nitric acid. Period.
Reactive Hazards and What Never to Mix
Nitric acid is a powerful oxidizer, and its reactions with organic materials can be explosively violent. This is not a theoretical concern. A documented industrial accident involved the inadvertent mixing of 62% nitric acid with isopropanol, which reacted to produce isopropyl nitrate, a compound used as a rocket propellant. The resulting explosion was devastating, and a review of the incident fifteen years later found that basic accident data had been corrupted and there was no evidence that broad safety lessons had been learned across the industry.5Safety Science. Violent explosion after inadvertent mixing of nitric acid and isopropanol – Review 15 years later finds basic accident data corrupted, no evidence of broad learning
The danger scales with concentration. Calorimetric studies of nitric acid mixed with n-hexane, a common hydrocarbon solvent, showed that as nitric acid concentration increased from about 3.7 to 5.4 moles per liter, the system’s thermal stability collapsed. The heat released exceeded the threshold for chemical explosion, the maximum rate of temperature rise jumped by at least sevenfold, and the maximum pressure rise rate increased at least elevenfold. Actual pressures inside test reactors reached far beyond rated limits, resulting in rupture.6Chinese Journal of Chemical Engineering. Investigation of a violent explosion for nitric acid–n-hexane system using calorimetric method
The practical takeaway is a short list of rules that should be treated as absolute:
- No organic solvents: Alcohols, ketones, ethers, and hydrocarbons must never contact nitric acid. Even residual traces on glassware can be enough to cause a violent reaction at higher acid concentrations.
- No metal powders or turnings: Finely divided metals react vigorously and can ignite, especially copper, zinc, and iron. Dissolving metals in nitric acid is a legitimate procedure but must be done slowly, in dilute acid, and with careful temperature control.
- No reducing agents: Sulfides, phosphorus, and charcoal are all incompatible. Hydrogen sulfide in particular reacts violently.
- No contact with cellulose or sawdust: These organic materials are easily nitrated and can become explosive compounds. Laboratory spills of nitric acid onto wooden benchtops have started fires.
Equipment and Material Compatibility
Nitric acid attacks many common materials, so your choice of glassware, tubing, and storage containers matters. Borosilicate glass is the standard for laboratory work and holds up well against all concentrations. Ordinary soda-lime glass is acceptable for dilute solutions but can be etched by concentrated acid over time.
For non-glass components like stoppers, tubing, and gaskets, fluoropolymers such as PTFE and FEP are the safest choices. Standard rubber and most common plastics are attacked by concentrated nitric acid. Silicone tubing, sometimes used in peristaltic pumps, degrades rapidly. Even some metals that seem corrosion-resistant have limits. Stainless steel resists dilute nitric acid reasonably well but can fail with concentrated or hot acid. Titanium performs much better; research has shown that thermally oxidized titanium samples exhibit markedly improved corrosion resistance in boiling nitric acid compared to untreated titanium or titanium alloys.7Surface and Coatings Technology. Thermal oxidation of titanium to improve corrosion resistance in boiling nitric acid medium
For storage, nitric acid should be kept in borosilicate glass bottles with PTFE-lined caps, in a dedicated acid cabinet away from organic chemicals and flammable materials. Concentrated nitric acid slowly decomposes in storage, especially if exposed to light or heat, producing dissolved nitrogen dioxide that turns the acid yellow and then brown. Keeping bottles in a cool, dark location extends their shelf life. Amber glass helps, but light exclusion alone is not enough if the storage area is warm.
Dealing with Off-Gases During Preparation
Even with a well-controlled distillation, some nitrogen dioxide will escape into the fume hood exhaust. In a research or teaching laboratory with a properly vented hood, this is usually an acceptable outcome because the quantities are small and the exhaust is diluted before discharge. But if you are working at a slightly larger scale, or if local regulations require it, you may need to scrub the exhaust gas.
The simplest scrubber for nitrogen dioxide is a caustic wash. Passing the off-gas through a solution of sodium hydroxide converts nitrogen dioxide into sodium nitrate and sodium nitrite. For more aggressive removal, sodium sulfide solutions have been shown to outperform sodium sulfite because of sulfide’s stronger reducing ability. Field tests found that for influent nitrogen dioxide concentrations of several tenths of parts per million, removal efficiency depended more on the sulfide dose than on the rate of gas-liquid contact.8Journal of Chemical Technology & Biotechnology. Reduction of nitrogen dioxide from etching vent gases by scrubbing with caustic sodium sulfide solution In a lab context, a simple bubbler or gas-washing bottle containing dilute sodium hydroxide solution at the end of your distillation train will catch most of the nitrogen dioxide that makes it through the condenser.
A good practice is to connect the collection flask’s vent to a length of tubing that dips into the scrubbing solution. This way, any gas that escapes the condenser encounters the caustic solution before reaching the fume hood’s airstream. The scrubbing solution should be replaced after each preparation because its capacity to absorb nitrogen dioxide is finite and drops off as it becomes saturated with dissolved nitrate and nitrite salts.
Aqua Regia and Other Nitric Acid Mixtures
One of the most well-known uses of nitric acid in the laboratory is as a component of aqua regia, the mixture of concentrated nitric acid and concentrated hydrochloric acid (traditionally in a 1:3 ratio by volume) that dissolves gold and platinum. The nitric acid acts as the oxidizer, converting the metal to ions, while the hydrochloric acid provides chloride ions that stabilize the dissolved metal as a soluble chloro complex. Neither acid alone dissolves gold at any meaningful rate.
Aqua regia must be prepared fresh before use. The mixture reacts with itself, producing highly reactive chlorine and nitrosyl chloride gases that are the actual dissolving agents. Left standing, it loses its potency as these species escape. Never store aqua regia in a sealed container; the gases it generates will build pressure and eventually shatter the bottle. Mix it in an open vessel inside a fume hood, use it immediately, and neutralize the waste when you are done.
Other useful nitric acid mixtures include dilute nitric acid for cleaning metal surfaces (it dissolves oxide layers and leaves a bright, passivated surface on stainless steel), and mixtures of nitric and hydrofluoric acid used to etch silicon wafers in semiconductor processing. Each of these applications exploits nitric acid’s oxidizing power in combination with another reagent that provides a specific complementary function.
Nitric Acid’s Role in the Atmosphere
Nitric acid is not just a laboratory and industrial chemical. It forms naturally in the atmosphere whenever nitrogen dioxide, produced by combustion, lightning, or biological processes, reacts with hydroxyl radicals and water vapor. For decades, atmospheric scientists treated the deposition of gaseous nitric acid onto surfaces as a permanent removal of nitrogen oxides from the air, effectively a dead end in the chain of reactions that produces ground-level ozone.
More recent modeling work has challenged that assumption. Simulations incorporating the reaction of gaseous nitric oxide with deposited nitric acid on surfaces suggest that this “renoxification” pathway can recycle nitrogen oxides back into the atmosphere, where they participate again in ozone formation. In some modeled regions, ozone predictions increased by as much as roughly 30 parts per billion, or about 20%, compared to simulations that treated nitric acid deposition as irreversible.9Atmospheric Environment. Modeling surface-mediated renoxification of the atmosphere via reaction of gaseous nitric oxide with deposited nitric acid This finding matters for air quality forecasting and helps explain why some urban areas have higher ozone levels than their nitrogen oxide emissions alone would predict.
For the lab chemist, the atmospheric chemistry of nitric acid is a reminder that this compound is highly reactive even at trace concentrations and on surfaces you might not expect. Residual acid on glassware, bench surfaces, and even skin can participate in reactions long after you think the preparation is finished. Thorough cleanup after every preparation session is not just good housekeeping; it prevents slow corrosion of equipment and unexpected reactions with materials left nearby.