Magnesium

• 1. The Brilliant White Light

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  Burning magnesium reaches a temperature of roughly 3,100 degrees Celsius, and it throws off a brilliant white light rich in ultraviolet wavelengths. That light once lit the earliest photographs, fired theatrical lightning, and made trick birthday candles relight themselves. It is also a fire that water cannot put out. Pour water on burning magnesium and the metal rips the water apart, releasing hydrogen gas that feeds the flame. Carbon dioxide does no better. The metal pulls the oxygen out of it, leaving behind carbon and burning hotter still. This is element 12, symbol Mg, a shiny gray metal with low density and high chemical reactivity. It is the eighth most abundant element in the Earth's crust and makes up 13% of the planet's mass. How did a metal this dangerous to ignite become the third most commonly used structural metal on Earth? Why did a German military aircraft program embrace it, and why did spectators at a French racetrack once get showered with its burning fragments? And how does the same element sit quietly at the center of every green leaf?

• 2. Forged In Dying Stars

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  In large, aging stars, magnesium is built by adding three helium nuclei in sequence to a carbon nucleus. When those stars explode as supernovas, the magnesium scatters into the interstellar medium, where it can recycle into new star systems. Some of the oldest objects in the Solar System carry a record of this process. The calcium-aluminium-rich inclusions inside certain carbonaceous chondrite meteorites hold an unusual surplus of a particular magnesium nuclide. That surplus is attributed to the decay of a parent isotope with a half-life of 717,000 years. Researchers conclude the meteorites formed in the solar nebula before that parent had decayed away. The signature lives on in how geologists read these rocks. They plot a magnesium ratio against an aluminium-to-magnesium ratio on an isochron, where the slope reveals not an age but the initial ratio when the systems first separated from a common reservoir. The same metal that lights a flare on Earth thus dates the birth of the planets.

• 3. The Skin That Protects

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  Magnesium reacts readily with air, yet it does not need an inert atmosphere for storage. The moment it meets air it grows a thin coating of magnesium oxide, and that passivation layer shields the metal beneath from further corrosion. This self-defense extends underwater. When magnesium sits in water, hydrogen bubbles form slowly on its surface as it produces magnesium hydroxide. The hydroxide builds up and blocks the reaction, so the process stays far less dramatic than the violent way alkali metals greet water. Calcium, a fellow group 2 metal, reacts faster. Powdered magnesium changes the picture entirely, reacting quickly enough that its reversible reaction with water can store energy and run a magnesium-based engine. The metal also reacts exothermically with acids like hydrochloric acid, giving magnesium chloride and hydrogen gas, much as aluminium and zinc do. As a reducing agent it strips oxygen from other compounds, which is why it can ignite thermite, a mixture of aluminium and iron oxide that lights only at very high temperature. That talent for handing over electrons reaches deep into the laboratory bench.

• 4. Grignard And Beyond

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  Organomagnesium compounds run through the heart of organic chemistry, most famously as Grignard reagents. They form when magnesium reacts with haloalkanes or aryl halides in diethyl ether, producing species like phenylmagnesium bromide and ethylmagnesium bromide. A Grignard reagent acts as a nucleophile, attacking an electrophilic carbon such as the one in the polar bond of a carbonyl group, and building new carbon-carbon bonds. Magnesocene, also called magnesium anthracene, offers a more aggressive option as a source of highly active magnesium. First prepared in 1954 by two independent groups, one led by Ernst Otto Fischer and the other by Albert Wilkinson, it is a white to off-yellow pyrophoric powder that hydrolyses violently in water. A related butadiene-magnesium adduct supplies the butadiene dianion, and chemists have even observed complexes of dimagnesium in the +1 state. Detecting the ordinary +2 ion is simpler. Adding ammonium chloride, ammonium hydroxide, and monosodium phosphate to a solution drops a white precipitate, while azo violet dye turns deep blue when an alkaline magnesium salt is present.

• 5. Light Metal, Hard Choices

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  Pure polycrystalline magnesium is brittle and fractures along shear bands, yet it is only two-thirds the density of aluminium. Add just 1% aluminium and it becomes far more malleable, and shrinking the grain size to about a micrometre improves it further. That trade between lightness and weakness has defined every alloy. As of 2013, the world used less than one million tonnes of magnesium alloys a year, against 50 million tonnes of aluminium alloys, held back by tendencies to corrode, creep, and combust. Corrosion is triggered by iron, nickel, copper, or cobalt, which form intermetallic compounds that act as cathodic sites and consume the metal. Enough manganese can overcome iron's effect, and adding about one part in three hundred of arsenic cuts the corrosion rate in salt solution by nearly tenfold. Creep at high temperature eases with zinc and rare-earth elements, while a small amount of calcium curbs flammability. The German answer to the lightness was a magnesium alloy they named Elektron, a term still used today. That name would later be written across the wreckage of a racetrack.

• 6. Wings, Engines And Disaster

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  Wright Aeronautical built a magnesium crankcase into the WWII-era Wright R-3350 Duplex Cyclone engine, and it nearly doomed the early Boeing B-29 Superfortress. An in-flight engine fire could ignite the crankcase, burning as hot as 5,600 degrees Fahrenheit, about 3,100 degrees Celsius, hot enough to sever the wing spar from the fuselage. On the road the same metal carried its own catastrophe. Mercedes-Benz used Elektron in the bodywork of an early Mercedes-Benz 300 SLR, which competed in the 1955 World Sportscar Championship, winning the Mille Miglia. At Le Mans that year one of these cars was caught in the 1955 Le Mans disaster, and spectators were showered with burning fragments of Elektron. Despite the danger, manufacturers kept coming back. Porsche ran magnesium alloy frames in the 917/053 that won Le Mans in 1971. BMW used a high-temperature alloy called AJ62A in its N52 engine, built from 2005 to 2011 across the 1, 3, 5, 6, and 7 series and the Z4, X1, X3, and X5. Chevrolet chose the alloy AE44 for the 2006 Corvette Z06, both alloys forming intermetallic precipitates at the grain boundaries to resist creep.

• 7. From Brine To Metal

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  World production of magnesium ran around 1,100 kilotonnes in 2017, with China making 930 kilotonnes and Russia 60. The United States once supplied 45% of world output as recently as 1995, but after China mastered the Pidgeon process the US share fell to 7%. By 2013 a single US producer remained, US Magnesium of the Renco Group, on the shores of the Great Salt Lake. The Pidgeon process now dominates the world. It calcines dolomite to magnesium oxide, then reduces it with silicon at high temperature, using a ferrosilicon alloy because it is cheaper than pure silicon. The calcium oxide grabs the silicon as an oxygen scavenger, leaving stable calcium silicate and gaseous magnesium that is condensed and collected. The rival route is electrolysis. The Dow process mixes seawater and dolomite, precipitates magnesium hydroxide with lime, converts it to magnesium chloride with hydrochloric acid, then splits the molten salt electrically at 680 to 750 degrees Celsius. That method releases harmful chlorine gas and demands enormous energy. Newer ideas chase a cheaper path. A solid oxide membrane process using yttria-stabilized zirconia was reported in 2011 to cut cost per pound by 40% over electrolytic reduction.

• 8. The Metal Inside Us

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  The human body holds about 22 grams of magnesium, with 60% in the skeleton, 39% inside cells, and 1% outside them. It is the eleventh most abundant element by mass in the body and essential to every cell. More than 300 enzymes require magnesium ions to work, including all those that use or make ATP and the nucleotides that build DNA and RNA. The ATP molecule normally sits in a chelate with a magnesium ion. Plants depend on it just as deeply. Magnesium occupies the center of the porphyrin ring in chlorophyll, the way iron sits at the center of the ring in heme, and without it photosynthesis cannot proceed. A shortage shows up as late-season yellowing between the veins of older leaves, correctable with epsom salts or crushed dolomitic limestone. In people, low plasma magnesium appears in 2.5 to 15% of the general population, and from 2005 to 2006 some 48% of Americans ate less than the recommended amount. The story of this element circles back to an English farm. In 1618 a farmer at Epsom found his cows refusing bitter well water that still healed their scratches, and the salts left behind on evaporation became epsom salts, later recognized as hydrated magnesium sulfate. The metal itself waited until 1808, when Sir Humphry Davy isolated it by electrolysis and first proposed calling it magnium.