Recovering Magnesium Salts from Seawater — The Lightweight Metal That Burns White

Magnesium (Mg, element 12) sits in Group 2 (alkaline earth metals), Period 3. It has an atomic weight of 24.305 and electron configuration [Ne] 3s². Like calcium (its heavier cousin in Group 2), magnesium readily gives up its two valence electrons to form the stable Mg²⁺ ion. The pure metal is silvery-white, lightweight (density 1.74 g/cm³), and moderately hard. It melts at 650 °C and boils at 1091 °C.

Magnesium's most dramatic property is its combustion: the metal burns in air with an intensely bright white flame (temperature ~2200 °C), producing magnesium oxide (MgO) and magnesium nitride (Mg₃N₂). Once ignited, burning magnesium cannot be extinguished with water — water reacts with hot magnesium to produce hydrogen gas, intensifying the fire: Mg + 2H₂O → Mg(OH)₂ + H₂↑. This is why magnesium fires must be smothered with dry sand or special Class D extinguishers, never water.

Collect 10–20 liters of clean seawater from a rocky shoreline or open-water location, away from river mouths (which dilute the salt content), harbors (pollution), and beach zones (where sand and debris contaminate samples). Average ocean water contains approximately 3.5% dissolved salts by weight, with the following approximate concentrations: sodium chloride (NaCl) 2.7%, magnesium chloride (MgCl₂) 0.33%, magnesium sulfate (MgSO₄) 0.23%, calcium sulfate (CaSO₄) 0.13%, and potassium chloride (KCl) 0.07%.

Filter the collected seawater through a cloth to remove visible debris and organic matter. Store in clean plastic or glass containers — not metal, as the salt water will corrode most metals.

Pour seawater into wide, shallow containers — traditional salt pans, baking trays, or any broad, flat vessels. Place them in direct sunlight in a location protected from rain. Solar evaporation concentrates the dissolved salts as water evaporates. In warm, dry climates, this takes 1–2 weeks; in cooler climates, it may take longer. You can also gently heat the pans over low fire to speed evaporation, but do not boil vigorously — rapid boiling creates fine salt spray that is lost.

As the water evaporates and salt concentration increases, salts begin to crystallize in a predictable order based on their solubility. The least soluble salt — calcium sulfate (gypsum, CaSO₄·2H₂O) — precipitates first as the solution reaches about one-third its original volume.

As evaporation continues past the gypsum stage, sodium chloride (NaCl) begins crystallizing as cubic crystals — common table salt. This is the main product of all traditional salt-making operations. When the liquid has reduced to about one-tenth its original volume, the pan bottom is covered with NaCl crystals. Carefully pour off the remaining liquid — this dark, bitter liquid is called bittern (or nigari in Japanese).

The bittern is the magnesium-rich fraction we want. It tastes intensely bitter because magnesium chloride (MgCl₂) triggers bitter taste receptors. Save all of the bittern in a separate clean container. This liquid contains most of the original seawater's magnesium, along with potassium, bromine, and remaining sodium salts.

Add slaked lime (calcium hydroxide, Ca(OH)₂) to the bittern while stirring. The hydroxide ions react with dissolved magnesium ions: MgCl₂ + Ca(OH)₂ → Mg(OH)₂↓ + CaCl₂. Magnesium hydroxide (Mg(OH)₂) is nearly insoluble in water and precipitates as a white, gelatinous solid that settles to the bottom. This is the same compound sold commercially as 'milk of magnesia' — an antacid and laxative.

Add the lime slowly and stir — too much lime wastes calcium hydroxide and introduces excess calcium into the solution. Stop adding lime when no more white precipitate forms (the liquid above the settled precipitate should be nearly clear). This precipitation reaction is the key step used in modern industrial magnesium production from seawater.

Let the white precipitate settle for several hours — overnight is ideal. Carefully pour off the clear liquid above (which contains calcium chloride and can be discarded). Collect the white Mg(OH)₂ paste by filtering through a fine cloth. Wash the precipitate once with clean water to remove residual salts, then allow it to dry.

The dried white powder is magnesium hydroxide — a mildly alkaline compound (pH 10.5 in suspension) used in medicine as an antacid, in water treatment to neutralize acidic waste, and as a fire retardant in plastics and building materials. It is also the starting material for producing magnesium oxide and metallic magnesium.

Instead of precipitation (step 5–6), the bittern can be slowly evaporated further to crystallize magnesium sulfate heptahydrate (MgSO₄·7H₂O) — Epsom salt. Named after Epsom in Surrey, England, where it was first distilled from mineral springs in 1618, Epsom salt forms prismatic, needle-like crystals with a bitter, salty taste. Slow cooling of concentrated bittern after NaCl removal preferentially crystallizes MgSO₄·7H₂O because its solubility decreases more steeply with temperature than MgCl₂.

Historically, Epsom salt was one of the first recognized chemical compounds — valued as a purgative (laxative), a bath salt for sore muscles, and a plant fertilizer (magnesium is essential for chlorophyll). It dissolves readily in water, producing a cool sensation on the skin.

Heat the dried magnesium hydroxide strongly in a crucible over a forge fire or charcoal furnace. At approximately 350 °C, Mg(OH)₂ decomposes: Mg(OH)₂ → MgO + H₂O. The resulting white powder is magnesia (magnesium oxide, MgO), one of the most refractory oxides known — it melts at 2852 °C, far higher than iron (1538 °C) or even alumina (2072 °C). This extraordinary melting point makes MgO essential as a furnace lining material in steelmaking.

Magnesia is also called 'calcined magnesia' or 'dead-burned magnesia' when heated above 1500 °C, producing a dense, chemically inert form used in refractory bricks. The lighter-burned form retains some reactivity and is used in agriculture (soil pH amendment), water treatment, and as a component in some cement formulations.

The same bittern you collected in step 4 is called nigari in Japanese (にがり, from 'nigai' — bitter). For millennia, Japanese and Chinese tofu makers have used nigari as the coagulant that turns soy milk into tofu. When a few drops of nigari (mainly MgCl₂) are added to hot soy milk, the magnesium ions cross-link soy proteins, causing them to aggregate into a soft curd. The amount of nigari and the temperature determine the tofu's firmness — silken tofu uses less, firm tofu uses more.

This is a beautiful example of how a single chemical — magnesium chloride — connects ocean chemistry to food culture. The ancient salt makers who discarded bitter bittern and the ancient tofu makers who prized it were handling the same magnesium-rich solution.

Magnesium compounds and metal connect to technology through several critical paths. Biology: magnesium is the central atom in chlorophyll — every plant on Earth depends on Mg²⁺ for photosynthesis. It is also essential in over 300 human enzymes, including those that synthesize DNA and produce ATP (cellular energy). Construction: MgO-based cements (Sorel cement, invented 1867) set harder than Portland cement and resist fire; ancient Chinese builders mixed MgO with organic compounds to create adhesives for the Great Wall. Metallurgy: magnesium alloys (with aluminum, zinc, and manganese) are used in aerospace, automotive, and electronics applications where light weight is critical — a magnesium aircraft wheel weighs 40% less than its aluminum equivalent. Medicine: Epsom salt baths, milk of magnesia (antacid), and magnesium supplements for muscle cramps and heart health. Light: magnesium's brilliant white flame made it the first practical photographic flash (1859) and remains the basis of military illumination flares and firework sparklers.

From the chlorophyll in every leaf to the alloys in spacecraft, magnesium is the lightweight element that punches far above its weight in human civilization.