In Sorting the Elements you were handed property data for a set of elements with the names stripped off, and asked to arrange them so that the arrangement said something. Groups sorted by melting point, by density, by how the numbers rose and fell. Every one of those sorts is real. The one chemistry settled on is not just a tidier filing system — it is a claim about matter, and it earned its place by making predictions that could have been wrong.
That is the difference worth holding on to. A classification that only organises is a cupboard. A classification that predicts is a theory.
An argument, not an inventory
Mendeleev published his table in 1869 with gaps in it. He had arranged the known elements in order of increasing atomic mass, noticed that similar properties came round again at regular intervals, and found that the pattern only held if he left empty spaces for elements nobody had found. Leaving a gap is a risky thing to do: it says the scheme is right and reality is incomplete.
He went further and described what would fill one of the gaps below silicon, an element he called eka-silicon. Fifteen years later Clemens Winkler isolated germanium.
| Property | Mendeleev predicted, 1871 | Measured for germanium |
|---|---|---|
| Atomic mass | about 72 | 72.63 |
| Density of the element | about 5.5 g/cmÂł | 5.32 g/cmÂł |
| Appearance | dark grey | greyish white, lustrous |
| Formula of the oxide | ||
| Formula of the chloride | ||
| Density of the chloride | about 1.9 g/cmÂł | 1.88 g/cmÂł |
| Boiling point of the chloride | below 100 °C | about 86 °C |
Nobody predicts a boiling point to that accuracy by luck. The table had told the truth about an element that had never been seen.
It also had a problem, and how the problem was fixed is the more useful half of the story. Ordered strictly by atomic mass, tellurium (127.6) comes after iodine (126.9) — but tellurium behaves like the other elements in iodine’s column, and iodine like the ones in tellurium’s. Mendeleev swapped them, on the grounds that the properties mattered more than the mass. He was overruling his own ordering rule with no justification beyond “otherwise the pattern breaks”.
He turned out to be right for a reason he could not have known. In 1913 Henry Moseley measured the X-rays emitted by elements and found a quantity that increased by exactly one from element to element: the number of protons in the nucleus, the atomic number. Reorder by atomic number and the anomalies disappear — tellurium is 52 and iodine is 53, in that order, no swap required. The rule was not bent; it was the wrong rule, and the properties had been signalling that for forty years.
So the periodic law is stated this way: when the elements are arranged in order of increasing atomic number, their physical and chemical properties recur at regular intervals. Every trend on Periodic Trends is a consequence of that sentence.
Atomic number, mass number, and isotopes
Two numbers describe a nucleus and they are constantly confused.
- The atomic number, , is the number of protons. It is what makes an atom that element. Change it and you have a different element.
- The mass number, , is the number of protons plus neutrons in one particular atom. It is always a whole number, because you cannot have part of a neutron.
Atoms of the same element with different numbers of neutrons are isotopes. Carbon-12 and carbon-13 both have six protons; one has six neutrons and the other seven. They sit in the same square of the table, have the same electron arrangement, and therefore behave the same way chemically — a carbon-13 atom forms exactly the bonds a carbon-12 atom does. What differs is mass, and anything that depends on mass.
Some isotopes have nuclei that are not stable. These are radioisotopes, and they decay over time, emitting radiation and becoming a different nuclide. Carbon-14 is the well-known one: it forms continuously in the upper atmosphere, is taken up by living things along with ordinary carbon, and begins decaying once an organism dies, which is what makes radiocarbon dating possible. Cobalt-60 and iodine-131 are used in medicine for the same property that makes them hazardous — the radiation they emit does something to the tissue it passes through.
Radioactivity is a nuclear property, not a chemical one. No reaction you run in this course changes a nucleus. Everything else in this course is electrons.
Why chlorine weighs 35.45
Look under the symbol for chlorine and you find 35.45. No chlorine atom has that mass. Chlorine-35 has a mass of 34.969 u and chlorine-37 has 36.966 u, and there is nothing in between.
The number printed on the table is the relative atomic mass: the average mass of the element’s atoms as they actually occur, weighted by how common each isotope is. Chlorine on Earth is about 75.76% chlorine-35 and 24.24% chlorine-37, so
The average is closer to 35 than to 37 because there is three times as much of the lighter isotope. This is a weighted average, not a midpoint — averaging the two masses without the abundances gives 35.97, which is wrong by far more than any measurement error.
Two consequences you will use constantly:
- Molar masses are decimals for a reason. When you add up a molar mass in Molar Mass and Composition you are adding averages, so the answer is an average too, and it applies to a sample rather than to one molecule.
- The number is about this planet. Isotopic abundances are measured from terrestrial samples. A rock from elsewhere in the solar system can differ slightly, which is exactly how geochemists identify meteorites.
What the table lets you do
You now have three things from one chart: the identity of an element (atomic number), the average mass of its atoms (relative atomic mass), and its position in a pattern that repeats. The third is the one with predictive power, and Periodic Trends takes it apart — why the pattern repeats, what changes along a row, and where the neat story is honestly wrong.
After that, Ionic and Covalent Bonding uses position to work out what two elements will do when you put them together.
Curriculum connection
B3.1
explain the relationship between the atomic number and the mass number of an element, and the difference between isotopes and radioisotopes of an element
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B3.2
explain the relationship between isotopic abundance of an element’s isotopes and the relative atomic mass of the element
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B3.3
state the periodic law, and explain how patterns in the electron arrangement and forces in atoms result in periodic trends (e.g., in atomic radius, ionization energy, electron affinity, electronegativity) in the periodic table
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