Isotope

The atoms that make up every mineral consist of three fundamental particles: protons and neutrons bound together in a dense central nucleus, and electrons orbiting around it. ^[1] The three particles differ in mass and electric charge.

One atomic mass unit (AMU) is defined as one-twelfth the mass of a ¹²C atom. ^[1] Electrons are nearly two thousand times lighter than protons or neutrons, so for most purposes the mass of an atom is determined entirely by its nucleus.

What makes one element chemically distinct from another is its proton count, not its mass. Every atom of a given element has a fixed, characteristic number of protons in its nucleus - the atomic number Z. ^[1] Hydrogen has Z = 1, oxygen Z = 8, and aluminium Z = 13. Change the proton count and the element changes entirely - what used to be oxygen becomes nitrogen or fluorine. Neutrons, however, are chemically neutral. They contribute to nuclear mass and stability but do not affect how the atom bonds, reacts, or enters a mineral structure. The number of neutrons in a nucleus can vary from atom to atom within the same element. Those variants are isotopes: atoms with the same Z but different neutron counts. ^[1]

Mass Number and Notation

The combined count of protons and neutrons in a nucleus is the mass number of that atom. ^[1] Because electrons are so light, the mass number is also a close approximation to the atom’s total mass in atomic mass units. Different isotopes of the same element are distinguished by writing the mass number as a superscript to the left of the chemical symbol - so ³⁹K and ⁴⁰K are two isotopes of potassium. The element letter tells you the element (Z = 19 for all potassium), and the superscript tells you the total nucleon count.

Potassium illustrates the concept clearly. Every potassium nucleus contains 19 protons (Z = 19 by definition), but potassium occurs in nature as three isotopes, each with a different number of neutrons. ^{[1] 39}K carries 20 neutrons, ⁴⁰K carries 21, and ⁴¹K carries 22. All three behave identically in chemical bonding and will substitute freely for one another in a growing K-feldspar or mica crystal - the mineral cannot distinguish between them. Their difference is purely one of nuclear mass. The neutron surplus grows with atomic number: for light elements, neutron and proton counts are roughly equal; for heavier elements, the nucleus needs progressively more neutrons to hold the increasingly crowded protons together. ^[1]

Atomic Mass and Atomic Weight

Atomic mass has a precise definition: it is the mass of a particular atom divided by one-twelfth the mass of a ¹²C atom. ^[1] For ¹²C itself the atomic mass is exactly 12 by definition. For every other isotope, the atomic mass is slightly different from its mass number - the difference arises because nuclear binding energy slightly changes the actual mass of the assembled nucleus relative to the sum of its parts. Each isotope of an element has its own atomic mass, because each has its own distinct nuclear composition.

The atomic weight of an element is not the mass of any single isotope but the weighted average of the atomic masses of all naturally occurring isotopes, each weighted by its natural abundance in the Earth’s crust. ^[1] For potassium: ³⁹K has atomic mass 38.964 and natural abundance 93.26%; ⁴⁰K has atomic mass 39.964 and abundance 0.01%; ⁴¹K has atomic mass 40.963 and abundance 6.73%. Taking those abundances as weights and averaging the three atomic masses yields 39.098 - the atomic weight reported on every periodic table. Because ³⁹K dominates at over 93%, the atomic weight sits just above 39. The rare ⁴¹K pulls it fractionally higher while the vanishingly rare ⁴⁰K contributes almost nothing. This is precisely the information the periodic table encodes when it gives a single decimal number for potassium: it is telling you the mixture of isotopes typical of Earth surface materials, not the mass of any one atom.

Atomic weight is therefore not a fixed physical constant of an element in the way Z is - it depends on the isotopic proportions of the sample in question. ^[1] For most geological work the variations between samples are negligible, but not always.

Geological Significance

The isotopic compositions of elements in minerals are not frozen at the moment of crystallisation. Radioactive isotopes decay over time, transforming into different elements entirely - and as they do, the isotopic proportions of both the decaying element and the element it produces change continuously and measurably. ^[1] A uranium-bearing mineral that contained no initial lead when it first crystallised will have accumulated lead ever since, in proportion to how much uranium has decayed. By measuring the uranium remaining and the lead that has accumulated, the time since crystallisation can be calculated. This is radiometric dating - one of the most powerful geochronological tools available.

More broadly, the actual isotopic composition of any element in a geological sample can differ from the internationally agreed standard values, and those deviations carry information even when they are too small to affect everyday calculations. ^[1] Stable isotopes of oxygen, carbon, and sulphur are fractionated by temperature, biological activity, and fluid-rock interaction in predictable ways, so their measured ratios in a mineral record something about the conditions of formation. The same minerals that look identical in hand specimen and under the microscope can carry quite different isotopic signatures - and it is those signatures that unlock their thermal history, source environment, or age.

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