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An atmospheric electrical discharge. Its high-temperature chemistry can produce nitrogen oxides, influencing ozone and radical chemistry away from surface emission sources. The ozone response depends on subsequent transport and the local chemical regime.

A regime in which additional available nitrogen oxides tends to increase net photochemical ozone production when other factors are fixed. Reducing nitrogen oxides generally suppresses that production, but local ozone concentration also depends on transport, deposition and NO titration.

A regime in which additional nitrogen oxides promote net tropospheric ozone production because NO availability limits peroxy-radical oxidation of NO. Reducing NOx then generally lowers ozone production, subject to transport and chemical history. At much larger NOx burdens, OH removal by nitric-acid formation can instead produce a VOC-limited ozone-production regime.

A contour plot of a specified ozone outcome against the loadings of nitrogen oxides and volatile organic compounds. The ridge separating opposite NOx sensitivities distinguishes an NOx-limited ozone-production regime from a VOC-limited ozone-production regime. The outcome may be peak concentration, a time-integrated measure or a model production rate; axes and averaging conditions must be stated. A schematic diagram is not a calibrated emissions-response prediction.

Peroxyacetyl nitrate forms reversibly from a peroxyacetyl radical and NO${}_2$. It stores reactive nitrogen while intact, then thermal dissociation can release it again. It can therefore transport nitrogen through colder free-tropospheric air and recycle it more rapidly in a warmer planetary boundary layer. Deposition or further chemistry can compete, but reversible storage is its reservoir role.

Dinitrogen pentoxide is likewise a reservoir because $\mathrm{N}{\mathrm{O}}_2+\mathrm{N}{\mathrm{O}}_3+\mathrm{M}\rightleftharpoons {\mathrm{N}}_2{\mathrm{O}}_5+\mathrm{M}$ can return reactive species. Conversion to nitric acid makes it a sink from the active nitrogen oxides family, especially through heterogeneous uptake on wet aerosol or cloud particles, followed by deposition or washout. Water vapour alone and wet-particle uptake are different mechanisms; they should not be assigned the same rate constant. Temperature, particle surface area and removal are often very different between the boundary layer and the free troposphere, so a reservoir can be transported in one environment and chemically removed in another.

Methane consumes hydroxyl radicals during its initial oxidation and therefore influences atmospheric oxidative capacity and the removal of other trace gases. Subsequent oxidation generates peroxy radicals. In the presence of sufficient nitrogen monoxide, these convert nitrogen monoxide to nitrogen dioxide without consuming ozone, allowing net photochemical ozone production. At low nitrogen monoxide, peroxide formation and ozone-destroying radical reactions compete. Methane oxidation ultimately produces carbon monoxide and carbon dioxide; the sign and rate of the associated ozone response depend on the radical/nitrogen oxides regime, not methane alone.

A humidity assumption is necessary: warming does not by itself specify how much water is available. At fixed relative humidity, the Clausius-Clapeyron equation gives the first-order vapor pressure responseIf water abundance is approximated as proportional to $e_s$, the usual first-order estimate is a $26.7\%$ increase in hydroxyl radical production. Integrating the constant-latent-heat expression instead of linearizing gives $e_s(289)/e_s(285)=\exp [(45050/8.314)(1/285-1/289)]=1.3010$.

There is a small but explicit concentration correction when the square brackets are literal molecules per volume. Since $[{\mathrm{H}}_2\mathrm{O}]=\varphi e_s/(k_{B}T)$ and $P_{\mathrm{OH}}=2k[\mathrm{O}({}^{1}D)][{\mathrm{H}}_2\mathrm{O}]$,The corresponding finite-change concentration result is $(285/289)(1.3010)-1=0.2830$. The $25.3\%$ linear concentration estimate and the $26.7\%$ vapor pressure approximation express different stated approximations; neither follows without a humidity condition.

Extra water increases the fraction of oxygen singlet D atoms reacting to form hydroxyl radicals rather than being collisionally quenched and recycled into ozone. In clean, low-nitrogen oxides air this tends to increase ozone loss. In nitrogen oxides-rich air, increased hydroxyl radical can accelerate precursor oxidation and ozone production, while radical termination also changes. Greater hydroxyl radical production therefore does not imply a universally positive or negative ozone response.

Lower carbon monoxide and VOC emissions generally reduce the supply of peroxy radicals from their oxidation and therefore reduce photochemical ozone production where nitrogen monoxide is available. Lower nitrogen oxides abundance suppresses the nitrogen monoxide-to-nitrogen dioxide pathway in NOx-limited ozone production regions. These are reasons for a regional/background ozone decrease when fossil-fuel precursor emissions decline.

The response can have the opposite local sign in a nitrogen monoxide-rich urban core: reducing nitrogen monoxide weakens ozone titration, and reducing nitrogen oxides in a VOC-limited ozone production regime can initially lessen radical termination and raise ozone. Less carbon monoxide can also leave more hydroxyl radical available to remove methane, partly opposing the hydroxyl radical demand from leaked hydrogen. Changes in precursor transport, chemistry and sunlight matter. Reduced emissions usually reduce net regional ozone production, but do not guarantee a lower ozone concentration at every monitoring site.

At fixed precursor emissions, more tropospheric water vapor in a warmer climate promotes the ozone-photolysis/O(${}^{1}D$)-water loss route, especially in clean tropical lower air. Warmer conditions also accelerate some precursor chemistry; lightning nitrogen oxides, stagnation and stronger stratosphere-to-troposphere ozone transport can increase ozone elsewhere. Thus near-surface regional ozone and the integrated tropospheric burden can respond differently. There is no universal decrease of the global burden from the humidity effect alone; competing chemistry and transport must be evaluated consistently.

Photodissociation of nitrogen dioxide forms an oxygen atom, followed by third-body recombination:If NO is converted back to nitrogen dioxide only by $\mathrm{NO}+{\mathrm{O}}_3\to \mathrm{N}{\mathrm{O}}_2+{\mathrm{O}}_2$, this is a null cycle with no net ozone production. Oxidation of a volatile organic compound by OH generates organic peroxy radicals, and often hydroperoxyl radicals, which instead oxidize NO without consuming ozone:The newly formed nitrogen dioxide can then photolyse, giving net ozone production. These chain reactions require sunlight, radical production and both precursor families.

In an NOx-limited ozone-production regime, NO is too scarce for efficient peroxy-radical conversion; competing peroxide-forming radical termination becomes important. Reducing nitrogen-oxide emissions then reduces ozone production. In a VOC-limited ozone-production regime, plentiful nitrogen oxides remove OH through $\mathrm{OH}+\mathrm{N}{\mathrm{O}}_2+\mathrm{M}\to \mathrm{HN}{\mathrm{O}}_3+\mathrm{M}$, so organic oxidation and radical supply are limiting. Reducing volatile-organic emissions is effective there, while a modest nitrogen-oxide reduction can initially increase ozone by reducing radical termination and NO titration. Sufficiently large nitrogen-oxide reductions can move the system into the limited regime and eventually lower ozone. The appropriate control therefore depends on the chemical regime and the regional transport history, not merely the largest local emission. The ozone isopleth diagram below uses an explicitly schematic response function to show derivative signs, not a fitted concentration or emissions-control prediction.^[ref]

A suitable historical answer gives both the chemical transformation and the means of isolating a reasonably pure product. The following are preparative principles, with modern chemical formulas; they are not instructions for self-preparation of medicines. In particular, hydrogen cyanide, mercury compounds and arsenic compounds present serious toxicity hazards. Hydrogen cyanide is extremely volatile, can be absorbed by several routes, and its odour is an unreliable warning.^[ref]

For hydrocyanic acid, a customary early nineteenth-century route is the controlled distillation of potassium ferrocyanide with dilute sulfuric acid and water. Hydrogen cyanide passes into the distillate; nonvolatile salts remain behind. A generating flask, efficient laboratory condenser and chilled receiver are needed because hydrogen cyanide is readily lost. The contemporary account of hydrocyanic acid preparation in Pharmacopoeia Castrensis Ruthenica in The British and Foreign Medical Review (1842) specifically describes the ferrocyanide–sulfuric acid method and its losses on imperfect condensation. The concentration must be determined analytically before any defined dilution: a nominal starting mixture does not establish the final strength.^[ref] At the level of the released cyanide, the essential acid-base reaction is $\mathrm{C}{\mathrm{N}}^{-}+{\mathrm{H}}^{+}\to \mathrm{HCN}$. This equation alone is not a complete equation for decomposing the stable ferrocyanide coordination complex. Concentrated sulfuric acid can lead to different decomposition chemistry, so the specified dilution is meaningful.

For formic acid, gently heat oxalic acid with glycerol and some water, then separate the aqueous formic acid by distillation. The useful net transformation isThe glycerol participates through intermediate esters and is regenerated rather than consumed in this net equation. Marcellin Berthelot's 1856 account describes the carbon dioxide evolution, subsequent water addition and distillation, and reuse of the residual glycerol. Overheating is disadvantageous because the formic acid can decompose further to carbon monoxide and water. The method therefore separates a volatile acid from a much less volatile carrier; it is not merely the dry destructive heating of oxalic acid.^[ref]

For oxalic acid, oxidize a carbohydrate, such as sucrose or starch, with nitric acid; concentrate the resulting chemical solution appropriately and cool it to obtain crystals, followed by recrystallization. Nitrogen oxides evolve, and excessive oxidation can carry carbon onward to carbon dioxide, so an idealized single-product equation would conceal losses. The use of starch and nitric acid for oxalic acid production is explicitly documented in an 1850 account of nitric oxide recovery.^[ref] An alternative already established before 1864 is alkaline treatment of sawdust, producing an oxalate which can be purified through sparingly soluble calcium oxalate.^[ref] Decomposing the isolated calcium oxalate with sulfuric acid separates much of the calcium as calcium sulfate:Filter off the calcium sulfate and crystallize the oxalic acid from the filtrate. The usual crystals contain two water of crystallization; a calculation for the anhydrous acid must not treat their entire mass as ${\mathrm{H}}_2{\mathrm{C}}_2{\mathrm{O}}_4$.

For calomel, intimately combine mercury(II) sulfate, metallic Mercury and dry sodium chloride, then collect the volatile mercury(I) chloride in a cooler receiving chamber. The net change isThe extra metallic Mercury lowers the average oxidation state from +2 to +1. The product is purified by appropriate sublimation and washing to remove soluble contamination, particularly mercury(II) chloride. Both the chemical composition and the fine subdivision of the product matter: condensation through a chamber can give a powder rather than an adherent crystalline crust. This physical part of calomel preparation is described in the 1843 primary account in Polytechnisches Journal volume 87 (1843); the account's historical therapeutic claims are not adopted here.^[ref]

For corrosive sublimate, heat dry mercury(II) sulfate with sodium chloride and collect the mercury(II) chloride in the cooler part of a suitable sublimation apparatus:Here no metallic Mercury is added. Nonvolatile sodium sulfate stays behind, while mercury(II) chloride is transferred and condensed. Further sublimation or recrystallization separates impurities. ^[ref] The crucial distinction from calomel preparation is whether Mercury remains in the +2 oxidation state or is converted to +1.

For the solid called arsenious acid, roast an arsenic-bearing ore in air and collect the volatile arsenic trioxide in cooler chambers; purify by resublimation. The essential oxidation is represented byActual ores also contain other chemical elements, so their roasting equations and gaseous by-products depend on composition. A contemporary 1860 report describes roasting arseniferous ores, collecting the volatile arsenic trioxide in chambers and refining it by another sublimation. It also demonstrates that apparently good material could contain antimony(III) oxide: physical appearance alone does not establish purity.^[ref] The product intended here is dry $\mathrm{A}{\mathrm{s}}_2{\mathrm{O}}_3$; dissolving it does not make the bottled solid a pure crystalline hydrate ${\mathrm{H}}_3\mathrm{As}{\mathrm{O}}_3$.

For tartar emetic, react antimony(III) oxide with potassium hydrogen tartrate in water. A small amount of water first gives intimate contact in a paste; subsequent heating with more water completes formation of the soluble potassium antimony tartrate. Filter away unreacted insoluble matter and cool the clear filtrate to crystallize the product. This sequence is directly described in a January 1864 discussion of the newly issued British Pharmacopoeia, rather than inferred from its later editions.^[ref] In the traditional antimonyl notation the underlying reaction, before assigning crystal water of crystallization, isThis notation expresses the overall composition; tartrate coordination complexes and water of crystallization need the qualification given in 4. The historical pharmaceutical use of tartar emetic does not make it a safe contemporary emetic.^[ref]

For potassium chlorate, pass chlorine into hot concentrated potassium hydroxide. The disproportionation gives potassium chloride together with potassium chlorate:Concentration and cooling permit fractional crystallization of the much less soluble cold-water potassium chlorate, leaving much of the potassium chloride in the mother liquor; recrystallization improves purity. Cold dilute alkali instead favours hypochlorite, so the heating condition is essential. An economical contemporary manufacturing alternative uses hot calcium hydroxide and potassium chloride; the overall accounting isThe 1850 account in Polytechnisches Journal volume 118 (1850) describes a steam-heated stirred vessel and recovery by cooling, rather than an electrical manufacturing process introduced retrospectively.^[ref]

A regime in which precursor/radical supply and NOx-dependent radical termination make net ozone production more responsive to additional reactive organic compounds than to additional nitrogen oxides. Reducing nitrogen oxides can initially raise ozone in some such conditions. Diagnose the chemical regime before extrapolating an emission reduction to every location.

A regime with abundant nitrogen oxides in which volatile-organic oxidation and radical supply limit net tropospheric ozone production. Reducing volatile organic compounds lowers production. A small NOx reduction can initially increase ozone by weakening OH termination and NO titration; sufficiently large reductions can cross into an NOx-limited ozone-production regime. Local concentration and regional emission sensitivities need not coincide.