How Does the Periodic Table Make Sense?

Brief History

First, let’s take a look at how the table developed. I doubt you’re here to get a history lesson, so I made this short and put in the more fun details. For all details, click here and here.

  • 1789, Lavoisier: Grouped elements based on properties and categorized gases, non-metals, metals, and earths.
  • 1829, Döbereiner: Grouped elements in triads with similar properties.
  • 1862, de Chancourtois: Made a 3D model – the telluric screw. As you turn the screw, atomics weights increased by certain proportions.
  • 1863, Newlands: Made a Law of Octaves. Saw similarities in elements differing in atomic weight by 7. This was before the discovery of the noble gases.
  • 1864-1870, Meyer: Produced multiple models. Organized elements by how many other atoms they could combine with. Made another model organizing elements by atomic weight.
  • 1869, Mendeleev: Discovered the pattern we use today to organize the elements. The most profound aspect of his table was that he left open spaces for elements which weren’t discovered yet.


  • Columns are called Groups (or Families)
  • Rows are called Periods
  • Groups have similar properties
  • For the A groups, the number next to it is the same number of valence electrons the element has (number of electrons on its outer, or valence, shell)
  • Main groups:
    • IA- Alkali metals
    • IIA- Alkaline earth metals
    • VIIA- Halogens
    • VIIIA- Noble gases (or Group 0)

Reading the Elements

  • Atomic number: This is the number of electrons and the number of protons (e.g., Carbon’s number is 6, so it has 6 electrons and 6 protons).
  • Symbol: This symbol comes from the Latin name of the element.
  • Atomic Mass (or Weight): Take the mass of each isotope of the element and multiply it by the percentage of its occurrence on Earth. Take the average of the products and this is the atomic mass. More info here.
  • Isotope: Same atom, different amount of neutrons. What defines an atom is its number of electrons and protons. There is no such thing, for example, as a Carbon atom with 5 electrons. Isotopes have similar chemical properties, but different physical properties.
  • Atomic Mass Units (AMU): The number of grams per mole of that element.
  • Mole: A specific number of atoms (same for every element- Avogadro’s Number).
  • Avogadro’s Number: Approximately 6.022×1023 atoms.
  • Neutron number? Atomic mass – Atomic number = Neutrons. The difference is rounded to the nearest whole number. This number is how many neutrons are in the pure element.
  • Properties
    • Physical– dependent on mass number.
    • Chemical– dependent on number of electrons.

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