Geology Reference
In-Depth Information
lion years ago. On the one hand, many rocks older than 2.5 billion years contain minerals
that are easily destroyed by the corrosive effects of oxygen, suggesting an oxygen-free en-
vironment prior to that time. Geologists find unweathered and rounded pebbles of pyrite
(the iron sulfide also known as fool's gold) and uraninite (the commonest uranium miner-
al) in ancient streambeds—places where such minerals would quickly corrode and break
down in today's oxygen-rich surface environment. Such ancient sandy layers also have a
telltalechemistry—unusuallyconcentratedinoxygen-avoidingelementslikecerium,while
strikingly deficient in others like iron compared with today's soils. These chemical quirks
further evince an atmosphere devoid of oxygen.
By contrast, rocks younger than 2.5 billion years contain many unambiguous signs of
oxygen. Between 2.5 and 1.8 billion years ago, massive deposits of iron oxides called ban-
ded iron formations were deposited in staggering abundance. These distinctive, dense ac-
cumulations of alternating black and rusty red layers hold 90 percent of the world's known
ironorereserves.Manganeseoxidesalsosuddenlyappear,asthicklystratifieddepositsthat
nowprovidetheworld'schiefrepositoriesofmanganeseores.Hundredsofothernewmin-
erals—oxidized ores of copper, nickel, uranium, and more—also appear in the rock record
for the first time after the Great Oxidation Event. Yet in spite of this expanded mineralo-
gical repertoire, some scientists remained unconvinced that the Great Oxidation Event was
reallyaneventatall.Perhapstherewasjustaslowandsteadyincreaseinatmosphericoxy-
gen. Perhaps the spotty, eroded rock record is incomplete and misleading.
The smoking gun of the Great Oxidation Event came from an unexpected
source—remarkable recent data on the isotopes of the common element sulfur. The 1990s
sawdramaticincreasesintheresolvingpowerandsensitivityofmassspectrometers,which
are the workhorse analytical instruments of the isotope world. The new generation of mass
spectrometers allowed scientists to analyze smaller and smaller samples, even microscopic
mineral grains or individual living cells, with higher and higher precision. Sulfur, one of
life's essential elements, proved a particularly tempting target for study, as there are four
stable sulfur isotopes in nature: sulfur-32, -33, -34, and -36. All of these isotopes have sul-
fur's requisite sixteen protons in the nucleus, but the number of neutrons varies from six-
teen to twenty.
The distribution of sulfur isotopes is usually predictable on the simple basis of mass.
All atoms wiggle, but less massive isotopes wiggle more. Consequently, in any chemical
reaction light isotopes are more likely to jump about than heavy ones. This selective pro-
cess, called isotope fractionation, occurs anytime a collection of sulfur atoms experiences
achemical reaction, whetherinasolidrockoralivingcell. Inthecaseofsulfur,anisotope
of mass 32 will typically fractionate more than an isotope of mass 34 or 36. What's more,
the fractionation ratio is usually directly related to the mass ratio of the isotopes: the frac-
tionation of sulfur-36 to sulfur-32 is almost always twice the fractionation of sulfur-34 to
