Chapter in a nutshell: Nitric acid (HNO₃, aqua fortis — "strong water") is a colourless, strongly acidic, monobasic liquid prepared in the laboratory by distilling a nitrate (KNO₃ or NaNO₃) with concentrated sulphuric acid below 200 °C. It is a typical strong acid (turns blue litmus red, neutralises bases, reacts with carbonates/sulphites) but is exceptional in that it is also a powerful oxidising agent, so it generally does not liberate hydrogen with metals — instead giving oxides of nitrogen (NO with dilute, NO₂ with concentrated). Its salts are nitrates, identified by the brown ring test.
1. Introduction, Names and Structure
Nitric acid is an important mineral acid, formerly called aqua fortis (Latin, "strong water") because it reacts with nearly all metals and even dissolves silver, which other common acids cannot.
| Property | Value / Fact |
|---|---|
| Molecular formula | HNO₃ |
| Relative molecular mass | 63 u (1 + 14 + 16×3) |
| Common name | Aqua fortis / spirit of nitre |
| Salts | Nitrates (e.g., NaNO₃, KNO₃) |
| Basicity | Monobasic (1 replaceable H⁺) |
Discovery / history:
- Glauber (1658) first obtained nitric acid by distilling nitre (KNO₃) with sulphuric acid.
- Lavoisier (1776) proved nitric acid contains oxygen.
- Cavendish (1784) proved it also contains hydrogen and nitrogen.
2. Occurrence and Formation in the Atmosphere
- Free state: found in traces in rain water (formed after lightning).
- Combined state: found as metallic nitrates — Chile saltpetre NaNO₃, nitre KNO₃, and calcium nitrate Ca(NO₃)₂.
During a thunderstorm the high temperature of a lightning bolt makes atmospheric N₂ and O₂ combine, and the acid is washed down as acid rain:
- Nitrogen + oxygen → nitric oxide (lightning discharge): $\mathrm{N_2 + O_2 \xrightarrow{\;lightning\;} 2NO}$
- Nitric oxide oxidised to nitrogen dioxide: $\mathrm{2NO + O_2 \longrightarrow 2NO_2}$
- NO₂ dissolves in rain water with O₂: $\mathrm{4NO_2 + 2H_2O + O_2 \longrightarrow 4HNO_3}$
The conversion of free atmospheric nitrogen into useful nitrogenous compounds in the soil is called fixation of atmospheric nitrogen.
3. Laboratory Preparation of Nitric Acid
Principle: A less volatile acid (conc. H₂SO₄) displaces a more volatile acid (HNO₃) from its salt on heating.
| Item | Detail |
|---|---|
| Reactants | A nitrate — potassium nitrate KNO₃ (nitre) or sodium nitrate NaNO₃ (Chile saltpetre) + concentrated H₂SO₄ |
| Ratio | Equal parts by weight |
| Temperature | Gently heated to 180–200 °C (below 200 °C) |
| Apparatus | All-glass retort (no rubber/cork) |
| Collection | Vapours condensed by a receiver chilled with cold water |
Reactions (below 200 °C — bisulphate formed):
$$\mathrm{KNO_3 + H_2SO_4(conc.) \xrightarrow{<200^{\circ}C} KHSO_4 + HNO_3}$$ $$\mathrm{NaNO_3 + H_2SO_4(conc.) \xrightarrow{<200^{\circ}C} NaHSO_4 + HNO_3}$$
Above 200 °C (normal sulphate forms — not preferred):
$$\mathrm{2KNO_3 + H_2SO_4(conc.) \xrightarrow{>200^{\circ}C} K_2SO_4 + 2HNO_3}$$ $$\mathrm{2NaNO_3 + H_2SO_4(conc.) \xrightarrow{>200^{\circ}C} Na_2SO_4 + 2HNO_3}$$ $$\mathrm{NaNO_3 + NaHSO_4 \xrightarrow{>200^{\circ}C} Na_2SO_4 + HNO_3}$$
Collection / colour: Vapours condense to a light yellow liquid. Pure acid is colourless; the laboratory sample is yellow because thermal decomposition of some HNO₃ liberates reddish-brown NO₂ which dissolves in the acid.
Important precautions & reasons (frequent exam questions)
| Precaution | Reason |
|---|---|
| All-glass apparatus (no cork/rubber) | HNO₃ vapours attack rubber and cork |
| Conc. H₂SO₄, not conc. HCl | HCl is volatile → its vapours distil over and contaminate the product |
| Temperature below 200 °C | Above 200 °C: hard sulphate crust sticks to retort, HNO₃ decomposes, fuel wasted |
Note: 68 % nitric acid cannot be concentrated further by simple distillation — it forms a constant boiling mixture (azeotrope) boiling at 121 °C.
4. Physical Properties
| # | Property | Description |
|---|---|---|
| 1 | Colour | Pure: colourless liquid (98%); commercial (68%): yellowish-brown |
| 2 | Odour | Suffocating smell |
| 3 | Taste | Acidic (sour) |
| 4 | Nature | Hygroscopic, fumes in moist air; bottle must be stoppered |
| 5 | Density | 1.54 g/cm³ (98%); 1.42 g/cm³ (68%) |
| 6 | B.p. / M.p. | Boils at 86 °C, melts at –42 °C |
| 7 | Solubility | Soluble in water in all proportions |
| 8 | Physiological | Corrosive; stains skin yellow by forming xanthoproteic acid with skin protein |
5. Chemical Properties — A. Stability & B. Acidic Character
(1) Instability (decomposition)
Pure nitric acid is unstable to heat and sunlight; it decomposes even at room temperature, giving the yellow NO₂ colour. Hence 100 % acid is rarely used, and it is stored in coloured bottles.$$\mathrm{4HNO_3 \xrightarrow{\;\Delta / sunlight\;} 4NO_2 + 2H_2O + O_2}$$
(2) Acidic properties — a strong monobasic acid
HNO₃ ionises almost completely in water:$$\mathrm{HNO_3(aq) \rightleftharpoons H^+ + NO_3^-}$$
(i) Action on indicators
| Indicator | Colour change |
|---|---|
| Blue litmus | turns red |
| Methyl orange | turns pink/red |
| Phenolphthalein | remains colourless |
$$\mathrm{CuO + 2HNO_3 \longrightarrow Cu(NO_3)_2 + H_2O}$$ $$\mathrm{ZnO + 2HNO_3 \longrightarrow Zn(NO_3)_2 + H_2O}$$ $$\mathrm{NaOH + HNO_3 \longrightarrow NaNO_3 + H_2O}$$ $$\mathrm{Mg(OH)_2 + 2HNO_3 \longrightarrow Mg(NO_3)_2 + 2H_2O}$$ $$\mathrm{Fe(OH)_3 + 3HNO_3 \longrightarrow Fe(NO_3)_3 + 3H_2O}$$
(iii) Reaction with carbonates & bicarbonates → salt + water + CO₂
$$\mathrm{CaCO_3 + 2HNO_3 \longrightarrow Ca(NO_3)_2 + H_2O + CO_2\uparrow}$$ $$\mathrm{Na_2CO_3 + 2HNO_3 \longrightarrow 2NaNO_3 + H_2O + CO_2\uparrow}$$ $$\mathrm{NaHCO_3 + HNO_3 \longrightarrow NaNO_3 + H_2O + CO_2\uparrow}$$ $$\mathrm{CuCO_3 + 2HNO_3 \longrightarrow Cu(NO_3)_2 + H_2O + CO_2\uparrow}$$
(iv) Reaction with sulphites & bisulphites → nitrate + water + SO₂
$$\mathrm{K_2SO_3 + 2HNO_3 \longrightarrow 2KNO_3 + H_2O + SO_2\uparrow}$$ $$\mathrm{Ca(HSO_3)_2 + 2HNO_3 \longrightarrow Ca(NO_3)_2 + 2H_2O + 2SO_2\uparrow}$$
6. Chemical Properties — C. Oxidising Properties
Nitric acid is a powerful oxidising agent owing to the nascent oxygen [O] released on decomposition. It is itself reduced to NO, NO₂, N₂O, etc., depending on concentration and temperature.
$$\mathrm{2HNO_3(conc.) \longrightarrow 2NO_2 + H_2O + [O]}$$ $$\mathrm{2HNO_3(dil.) \longrightarrow 2NO + H_2O + 3[O]}$$
(i) Action on non-metals (conc. acid → CO₂/oxide + NO₂ + H₂O)
$$\mathrm{C + 4HNO_3(conc.) \longrightarrow CO_2 + 4NO_2 + 2H_2O}$$ $$\mathrm{S + 6HNO_3(conc.) \longrightarrow H_2SO_4 + 6NO_2 + 2H_2O}$$ $$\mathrm{P_4 + 20HNO_3(conc.) \longrightarrow 4H_3PO_4 + 20NO_2 + 4H_2O}$$
(Phosphorus may also be written as: P + 5HNO₃(conc.) → H₃PO₄ + 5NO₂ + H₂O.)
(ii) Oxidation of compounds
$$\mathrm{H_2S + 2HNO_3(conc.) \longrightarrow 2H_2O + 2NO_2 + S}$$ $$\mathrm{SO_2 + 2HNO_3 \longrightarrow H_2SO_4 + 2NO_2\uparrow}$$
7. Action on Metals
As an oxidising agent, nitric acid generally does NOT liberate hydrogen with metals; it gives oxides of nitrogen instead, the product depending on concentration & temperature. It reacts with all metals except gold and platinum.
| Condition | Gas evolved | Example |
|---|---|---|
| Cold dilute HNO₃ | Nitric oxide (NO) | Cu, Zn, Fe |
| Conc. (or hot dilute) | Nitrogen dioxide (NO₂) | Cu, Zn, Fe |
| Very dilute (~1%) with Mg / Mn | Hydrogen (H₂) | Mg, Mn only |
$$\mathrm{3Cu + 8HNO_3(dil.) \longrightarrow 3Cu(NO_3)_2 + 4H_2O + 2NO}$$ $$\mathrm{3Zn + 8HNO_3(dil.) \longrightarrow 3Zn(NO_3)_2 + 4H_2O + 2NO}$$ $$\mathrm{3Fe + 8HNO_3(dil.) \longrightarrow 3Fe(NO_3)_2 + 4H_2O + 2NO}$$ $$\mathrm{3Ag + 4HNO_3(dil.) \longrightarrow 3AgNO_3 + 2H_2O + NO}$$
(b) Concentrated acid → nitrate + water + NO₂
$$\mathrm{Cu + 4HNO_3(conc.) \longrightarrow Cu(NO_3)_2 + 2H_2O + 2NO_2}$$ $$\mathrm{Zn + 4HNO_3(conc.) \longrightarrow Zn(NO_3)_2 + 2H_2O + 2NO_2}$$ $$\mathrm{Fe + 6HNO_3(conc.) \longrightarrow Fe(NO_3)_3 + 3H_2O + 3NO_2\uparrow}$$ $$\mathrm{Ag + 2HNO_3(conc.) \longrightarrow AgNO_3 + H_2O + NO_2}$$
Copper dissolves in conc. nitric acid to give a blue solution and dense reddish-brown NO₂ of pungent smell, which turns freshly-prepared acidified ferrous sulphate solution brown-black.
(c) Passivity: Iron, aluminium, cobalt and nickel become passive (inert) with cold conc. nitric acid due to a thin insoluble protective oxide layer. Passivity is removed by rubbing with sand paper or using a strong reducing agent.
(d) Very dilute (~1%) acid + Mg / Mn → hydrogen (the only proof that nitric acid contains hydrogen):
$$\mathrm{Mg + 2HNO_3(v.dil.) \longrightarrow Mg(NO_3)_2 + H_2}$$ $$\mathrm{Mn + 2HNO_3(v.dil.) \longrightarrow Mn(NO_3)_2 + H_2}$$
8. Aqua Regia
Aqua regia ("royal water") = 1 part conc. HNO₃ + 3 parts conc. HCl (by volume). It dissolves even gold and platinum because it produces nascent chlorine [Cl] that attacks them, forming soluble chlorides.
$$\mathrm{HNO_3 + 3HCl \longrightarrow NOCl + 2H_2O + 2[Cl]}$$ $$\mathrm{Au + 3[Cl] \longrightarrow AuCl_3\ (gold(III)\ chloride,\ soluble)}$$ $$\mathrm{Pt + 4[Cl] \longrightarrow PtCl_4\ (platinum(IV)\ chloride,\ soluble)}$$
Here HNO₃ oxidises HCl to chlorine, and NOCl = nitrosyl chloride.
9. Uses of Nitric Acid
- Etching designs on copper and brassware (acts as solvent for most metals except noble metals).
- Purifying gold (dissolves Cu, Ag, Zn, Pb impurities).
- Rocket fuel oxidant.
- Manufacture of fertilizers: calcium nitrate Ca(NO₃)₂, ammonium nitrate NH₄NO₃, nitro-chalk [NH₄NO₃ + CaCO₃], basic calcium nitrate.
- Manufacture of explosives — T.N.T., nitroglycerine, dynamite.
- Synthetic fibres — artificial silk, nylon, celluloid, plastics, photographic film.
- Making nitrates of K, NH₄⁺, Ag, and dyes, drugs, perfumes.
10. Tests for Nitric Acid and Nitrates
- On heating, conc. nitric acid gives brown fumes of NO₂:
- Nitrates (other than Na, K, NH₄⁺) on heating give reddish-brown NO₂:
- Adding copper to HNO₃ / acidified nitrate gives dense reddish-brown NO₂ fumes.
Brown Ring Test (confirmatory test for nitrate ion NO₃⁻)
Procedure: To the nitrate solution add freshly prepared saturated iron(II) sulphate (FeSO₄) solution; then add conc. H₂SO₄ carefully down the side to form a layer at the bottom; cool. A brown ring appears at the junction.$$\mathrm{6FeSO_4 + 3H_2SO_4 + 2HNO_3 \longrightarrow 3Fe_2(SO_4)_3 + 4H_2O + 2NO}$$ $$\mathrm{FeSO_4 + NO \longrightarrow FeSO_4\!\cdot\!NO\ (nitroso\ ferrous\ sulphate,\ brown)}$$
- Freshly prepared FeSO₄ is used because a stale solution is oxidised to ferric sulphate, which gives no brown ring.
- The ring of nitroso ferrous sulphate forms at the junction (heavier conc. H₂SO₄ settles below); on shaking/heating it decomposes and the ring vanishes.
11. Effects of Heat on Nitrates
| Type of nitrate | Products on heating | Equation |
|---|---|---|
| Na, K (alkali metals) | nitrite + O₂ only | 2NaNO₃ → 2NaNO₂ + O₂↑ ; 2KNO₃ → 2KNO₂ + O₂↑ |
| Other metals (Ca, Zn, Cu, Pb, Mg…) | oxide + NO₂ + O₂ | see below |
| Ag, Hg | metal + NO₂ + O₂ | see below |
| Ammonium nitrate | N₂O (laughing gas) + H₂O, no residue | NH₄NO₃ → N₂O↑ + 2H₂O |
- Na/K nitrate gives only oxygen (no NO₂) and leaves a residue (nitrite); a glowing splinter relights near the melt.
- Ammonium nitrate leaves no residue and gives laughing gas N₂O.
- Zinc nitrate residue (ZnO) is yellow when hot, white when cold; lead nitrate decrepitates and its PbO residue fuses with glass.
Worked Examples
1. Molar mass / basicity. Compute the relative molecular mass of HNO₃ and state its basicity. Solution: M = 1 + 14 + (16×3) = 63 u. It has one replaceable H⁺, so it is monobasic.
2. Mass of acid to neutralise NaOH. What mass of HNO₃ neutralises 4 g of NaOH? (NaOH = 40) Solution: NaOH + HNO₃ → NaNO₃ + H₂O. Moles NaOH = 4/40 = 0.1 mol → 0.1 mol HNO₃ = 0.1 × 63 = 6.3 g.
3. Volume of NO₂ from conc. acid + copper. 0.5 mol Cu reacts with conc. HNO₃. Find volume of NO₂ at STP. Solution: Cu + 4HNO₃ → Cu(NO₃)₂ + 2H₂O + 2NO₂. 1 mol Cu → 2 mol NO₂, so 0.5 mol → 1 mol NO₂ = 22.4 L at STP.
4. NO from dilute acid + copper. 0.6 mol Cu reacts with dilute HNO₃. Find moles of NO. Solution: 3Cu + 8HNO₃ → 3Cu(NO₃)₂ + 4H₂O + 2NO. 3 mol Cu → 2 mol NO, so 0.6 mol Cu → (2/3)(0.6) = 0.4 mol NO.
5. Acid needed for the brown ring. How many moles of HNO₃ react with 6 mol FeSO₄ in the brown ring test? Solution: 6FeSO₄ + 3H₂SO₄ + 2HNO₃ → 3Fe₂(SO₄)₃ + 4H₂O + 2NO. 6 mol FeSO₄ needs 2 mol HNO₃.
6. Identifying a gas. A blue crystalline solid is heated and gives a reddish-brown gas that relights a glowing splint, leaving a black residue. Identify the solid and gases. Solution: The solid is copper nitrate Cu(NO₃)₂; 2Cu(NO₃)₂ → 2CuO (black) + 4NO₂ (reddish-brown) + O₂ (relights splint).
7. Sulphur with conc. acid. Write the balanced equation when sulphur is oxidised by conc. HNO₃ and name the acid formed. Solution: S + 6HNO₃(conc.) → H₂SO₄ + 6NO₂ + 2H₂O. Sulphuric acid is formed.
Key Terms — Quick Glossary
| Term | Meaning |
|---|---|
| Aqua fortis | Old name of nitric acid ("strong water") |
| Aqua regia | 3 HCl : 1 HNO₃ (conc.); dissolves Au & Pt |
| Nitre / KNO₃ | Potassium nitrate |
| Chile saltpetre / NaNO₃ | Sodium nitrate |
| Monobasic acid | Acid with one replaceable H⁺ |
| Nascent oxygen [O] | Atomic oxygen freed on decomposition; cause of oxidising power |
| Passivity | Inertness of Fe/Al/Co/Ni in conc. HNO₃ due to oxide film |
| Xanthoproteic acid | Yellow compound formed when HNO₃ stains skin protein |
| Brown ring | Nitroso ferrous sulphate FeSO₄·NO; confirms NO₃⁻ |
| Nitrosyl chloride | NOCl, formed in aqua regia |
| Constant boiling mixture | 68% HNO₃ azeotrope boiling at 121 °C |
| Fixation of nitrogen | Conversion of free N₂ to soil nitrogenous compounds |
Common Mistakes to Avoid
- Don't say nitric acid liberates H₂ with metals. It is an oxidising agent → gives NO (dilute) or NO₂ (conc.). The only exception is very dilute (~1%) acid with Mg/Mn, which gives H₂.
- Don't use conc. HCl in place of conc. H₂SO₄ for lab preparation — HCl is volatile and contaminates the product.
- Don't heat above 200 °C in the preparation (hard crust + decomposition + fuel waste).
- Brown ring test uses FeSO₄ (iron(II) sulphate), not Fe₂(SO₄)₃, and must be freshly prepared; the acid added is conc. H₂SO₄, not HNO₃.
- Na/K nitrates give only O₂ (no NO₂) and leave a residue (nitrite); NH₄NO₃ leaves no residue and gives N₂O. Don't mix these up.
- Pure HNO₃ is colourless; the yellow colour is dissolved NO₂, not an inherent property.
- Gold and platinum do NOT react with nitric acid alone — only with aqua regia.
Likely Exam Questions (with crisp answers)
- Q: Why is nitric acid called aqua fortis? A: Latin for "strong water"; it reacts with nearly all metals and even dissolves silver.
- Q: Why is all-glass apparatus used in its preparation? A: HNO₃ vapours attack rubber and cork.
- Q: Why is conc. H₂SO₄ used, not conc. HCl? A: HCl is volatile; its vapours would distil over and contaminate the nitric acid.
- Q: Why keep temperature below 200 °C? A: Above 200 °C a hard sulphate crust sticks to the retort, the acid may decompose, and fuel is wasted.
- Q: Why is laboratory nitric acid yellow though pure acid is colourless? A: Dissolved NO₂ from partial thermal decomposition gives the yellow colour.
- Q: How is the yellow colour removed? A: Bubble dry air/CO₂ (drives out NO₂) or add excess water (NO₂ dissolves back).
- Q: Why is nitric acid stored in coloured bottles? A: It decomposes in sunlight (4HNO₃ → 4NO₂ + 2H₂O + O₂).
- Q: Why does nitric acid generally not give hydrogen with metals? A: It is a strong oxidising agent and oxidises any H₂ formed to water; it gives oxides of nitrogen.
- Q: Name two metals that give H₂ with nitric acid and the condition. A: Magnesium and manganese with very dilute (~1%) acid.
- Q: What is passivity? How is it removed? A: Fe/Al/Co/Ni become inert in cold conc. HNO₃ due to an oxide film; removed by rubbing with sandpaper or a strong reducing agent.
- Q: Gas when copper reacts with (a) dilute (b) conc. nitric acid? A: (a) NO (nitric oxide), (b) NO₂ (nitrogen dioxide).
- Q: What is aqua regia and its ratio? A: 1 part conc. HNO₃ + 3 parts conc. HCl; dissolves gold and platinum.
- Q: Which gas/compound is responsible for the brown ring? A: Nitroso ferrous sulphate, FeSO₄·NO.
- Q: Which nitrate on heating gives only oxygen? A: Sodium or potassium nitrate (gives nitrite + O₂).
- Q: Which nitrate leaves no residue on heating, and what gas forms? A: Ammonium nitrate; gives N₂O (laughing gas) and water.
- Q: Name the catalyst in Ostwald's process / why is HNO₃ molecular mass 63? A: Platinum; M = 1+14+48 = 63 u.
- Q: What is formed when nitric acid stains the skin yellow? A: Xanthoproteic acid.
Chemical Equations & Formulas (quick reference)
Key formulas: HNO₃ (63 u, monobasic), KNO₃ (nitre), NaNO₃ (Chile saltpetre), Ca(NO₃)₂, NH₄NO₃, NOCl (nitrosyl chloride), FeSO₄·NO (brown ring).
Atmospheric formation: N₂ + O₂ → 2NO ; 2NO + O₂ → 2NO₂ ; 4NO₂ + 2H₂O + O₂ → 4HNO₃
Laboratory preparation
- KNO₃ + H₂SO₄(conc.) →(<200°C) KHSO₄ + HNO₃
- NaNO₃ + H₂SO₄(conc.) →(<200°C) NaHSO₄ + HNO₃
- 2KNO₃ + H₂SO₄ →(>200°C) K₂SO₄ + 2HNO₃
Decomposition: 4HNO₃ →(Δ/sunlight) 4NO₂ + 2H₂O + O₂
Ionisation & neutralisation
- HNO₃ ⇌ H⁺ + NO₃⁻
- CuO + 2HNO₃ → Cu(NO₃)₂ + H₂O
- NaOH + HNO₃ → NaNO₃ + H₂O
- Fe(OH)₃ + 3HNO₃ → Fe(NO₃)₃ + 3H₂O
Carbonates / sulphites
- CaCO₃ + 2HNO₃ → Ca(NO₃)₂ + H₂O + CO₂
- NaHCO₃ + HNO₃ → NaNO₃ + H₂O + CO₂
- K₂SO₃ + 2HNO₃ → 2KNO₃ + H₂O + SO₂
Oxidation of non-metals/compounds
- C + 4HNO₃ → CO₂ + 4NO₂ + 2H₂O
- S + 6HNO₃ → H₂SO₄ + 6NO₂ + 2H₂O
- P₄ + 20HNO₃ → 4H₃PO₄ + 20NO₂ + 4H₂O
- H₂S + 2HNO₃ → 2H₂O + 2NO₂ + S ; SO₂ + 2HNO₃ → H₂SO₄ + 2NO₂
Action on metals
- 3Cu + 8HNO₃(dil.) → 3Cu(NO₃)₂ + 4H₂O + 2NO
- Cu + 4HNO₃(conc.) → Cu(NO₃)₂ + 2H₂O + 2NO₂
- 3Zn + 8HNO₃(dil.) → 3Zn(NO₃)₂ + 4H₂O + 2NO
- Zn + 4HNO₃(conc.) → Zn(NO₃)₂ + 2H₂O + 2NO₂
- 3Fe + 8HNO₃(dil.) → 3Fe(NO₃)₂ + 4H₂O + 2NO
- Fe + 6HNO₃(conc.) → Fe(NO₃)₃ + 3H₂O + 3NO₂
- 3Ag + 4HNO₃(dil.) → 3AgNO₃ + 2H₂O + NO
- Mg + 2HNO₃(v.dil.) → Mg(NO₃)₂ + H₂
- Mn + 2HNO₃(v.dil.) → Mn(NO₃)₂ + H₂
Aqua regia
- HNO₃ + 3HCl → NOCl + 2H₂O + 2[Cl]
- Au + 3[Cl] → AuCl₃ ; Pt + 4[Cl] → PtCl₄
Brown ring test
- 6FeSO₄ + 3H₂SO₄ + 2HNO₃ → 3Fe₂(SO₄)₃ + 4H₂O + 2NO
- FeSO₄ + NO → FeSO₄·NO
Heat on nitrates
- 2NaNO₃ → 2NaNO₂ + O₂ ; 2KNO₃ → 2KNO₂ + O₂
- 2Cu(NO₃)₂ → 2CuO + 4NO₂ + O₂ ; 2Pb(NO₃)₂ → 2PbO + 4NO₂ + O₂
- 2Ca(NO₃)₂ → 2CaO + 4NO₂ + O₂ ; 2Zn(NO₃)₂ → 2ZnO + 4NO₂ + O₂
- 2AgNO₃ → 2Ag + 2NO₂ + O₂ ; Hg(NO₃)₂ → Hg + 2NO₂ + O₂
- NH₄NO₃ → N₂O + 2H₂O