Hydrochloric Acid

Part of Acids and Alkalis

Production, handling, and uses of hydrochloric acid (muriatic acid) — a strong acid accessible via several routes with simple equipment.

Why This Matters

Hydrochloric acid (HCl) is one of the most versatile strong acids available to an early industrial civilization. It dissolves metals and metal oxides rapidly, making it ideal for cleaning and preparing metal surfaces. Mixed with nitric acid, it becomes aqua regia — the only acid that dissolves gold and platinum. It is used in tanning, textile processing, food production (making gelatin, processing corn), and as a precursor for producing metal chloride salts.

Unlike sulfuric acid, which requires the lead-chamber process and specialized roasting equipment, hydrochloric acid can be produced by several relatively accessible routes. The earliest industrial chemistry programs typically developed HCl before sulfuric acid precisely because the production equipment is simpler.

The acid is corrosive and produces irritating fumes, but these hazards are manageable with basic protective measures and good ventilation. For a rebuilding community that has established basic metalworking, pottery, and salt production, HCl is within reach.

Production Routes

Route 1: Salt and Sulfuric Acid (Preferred when H₂SO₄ available)

This is the classical industrial production method:

NaCl + H₂SO₄ → NaHSO₄ + HCl (at moderate temperature) NaCl + NaHSO₄ → Na₂SO₄ + HCl (at high temperature)

Both reactions proceed at temperatures achievable in a pottery kiln or hearth. The HCl gas produced must be captured by dissolving it in water.

Procedure:

1. Mix common salt (sodium chloride) with concentrated sulfuric acid in a ceramic retort — one part salt to roughly 1.2 parts acid by weight
2. Heat gently at first. HCl gas begins to evolve around 150°C.
3. Lead the gas through a clay tube into a vessel of cold water, or bubble it through water in a sealed container
4. HCl dissolves readily in cold water to form hydrochloric acid solution
5. At higher temperature (above 300°C), a second reaction converts the sodium bisulfate to sodium sulfate and more HCl — continue heating for full conversion

The byproduct sodium sulfate (Glauber’s salt) is useful as a laxative and in glassmaking.

HCl gas hazards

Hydrogen chloride gas is severely irritating to the respiratory tract, eyes, and skin. It reacts with moisture in mucous membranes to form hydrochloric acid. All operations involving HCl gas must take place outdoors or with strong downwind ventilation. A cloth soaked in baking soda solution held over the face provides partial protection for short exposures.

Route 2: Burning Salt in Acid Flame

A simpler, lower-yield method: Salt will release traces of HCl when strongly heated with acidic materials. Historical methods included burning seaweed ash (salt-rich) with sulfur-containing minerals. This produces very dilute HCl mixed with many other products — useful only as a last resort.

Route 3: Salt and Acid Clay (Volcanic Regions)

In regions with volcanic sulfur deposits or acidic hot springs, naturally acidic mineral waters can be reacted with salt to produce dilute HCl. The naturally occurring sulfuric acid does the conversion. This is a geological windfall rather than a designed process.

Route 4: Combustion of Organic Chlorides

Burning PVC plastic (if salvaged), certain salt-treated wood, or kelp/seaweed produces HCl gas as a combustion product. This is uncontrolled, produces a mixture of gases, and should only be considered if the above routes are completely inaccessible. The gas collection setup remains the same.

Apparatus for HCl Production

A practical retort system for small-scale production:

Components needed:

• A ceramic retort or clay-sealed flask (reaction vessel)
• A clay or lead pipe for gas transport (iron is attacked by HCl gas — use clay)
• A collection vessel — a wide-necked ceramic pot or barrel filled with cold water
• The gas inlet tube should extend below the water surface to ensure dissolution

Key design features:

• Keep the collection vessel cold (surround with cool water or ice if available)
• Do not create an airtight system — pressure buildup will shatter vessels; a water-trap outlet serves as a pressure relief
• The water in the collection vessel will become increasingly concentrated HCl as it absorbs gas

At 20°C, water can dissolve approximately 40% HCl by weight before saturation (this is “concentrated” HCl, also called fuming hydrochloric acid). Practical solutions of 20–30% are achievable without elaborate equipment.

Concentration and Handling

Standard production will yield dilute HCl (5–15%). For most applications this is adequate. To concentrate:

1. Distill the dilute solution — HCl and water form a constant-boiling mixture at ~20% concentration at atmospheric pressure, called an azeotrope. You cannot concentrate beyond this point by normal distillation.
2. To get above 20%, cool the collection vessel further (ice water) and use higher gas flow rates — the gas concentration at the inlet determines the solution concentration.

Storage: Glass or clay vessels. Never iron, zinc, or aluminum — these are attacked by HCl. Lead is acceptable for short-term storage. Seal containers tightly as HCl vapors escape from solution.

Dilution: Always add acid to water, never water to acid. The dissolution of HCl in water releases heat.

Aqua Regia: The Gold-Dissolving Mixture

One part concentrated nitric acid mixed with three parts concentrated hydrochloric acid produces aqua regia — the only common acid mixture capable of dissolving noble metals:

Au + 3 HNO₃ + 4 HCl → HAuCl₄ + 3 HNO₂ + H₂O (simplified)

This enables:

• Assaying and refining gold and platinum
• Recovering precious metals from ore or electronic salvage
• Producing soluble gold compounds for electroplating

Aqua regia must be freshly mixed — it decomposes over hours. Work only in very well-ventilated conditions as nitrosyl chloride and chlorine gas are released.

Practical Applications

Application	HCl Concentration	Procedure
Metal cleaning/pickling	5–15%	Immerse metal briefly, rinse with water
Removing scale from iron	10–20%	Soak until scale bubbles off; neutralize with lime water
Tanning leather (bating)	2–5%	Soak hides to break down fiber structure
Producing ferric chloride (etching agent)	10–20%	Dissolve iron in acid; concentrate
Food-grade gelatin production	1–3%	Acid hydrolysis of collagen-rich bones
Ore dissolution (acid leaching)	10–20%	Dissolve carbonate and oxide ores
Testing limestone/carbonate	Any dilute	Fizzing confirms calcium carbonate present

Safety Summary

• Always work outdoors or in strong crosswind
• Keep sodium bicarbonate (baking soda) or lime solution ready for spills and burns
• Neutralize waste acid before disposal — pour onto limestone chips or into lime water
• Never mix with bleach (sodium hypochlorite) — produces chlorine gas
• Fumes are detectable by smell well below dangerous concentrations; if you smell it, improve ventilation immediately
• Eye protection is essential — even dilute HCl causes painful eye burns

Hydrochloric acid is one of the workhorses of early industrial chemistry. Its accessibility from salt — a universally available commodity — means that any civilization with reasonable metallurgy and heat sources can establish HCl production relatively early in its industrial development.