Core Concepts
In this article, you will learn about atomic number, its definition, its usefulness in categorizing elements, and its history as a theory in chemistry.
Topics Covered in Other Articles
- How to Read the Periodic Table
- What is an Isotope?
- Nuclear Reactions
- Balancing Chemical Equations
- JJ Thompson and the Cathode Ray Tube
- James Chadwick and the Neutron
The Elements
What exactly makes one element different from another? Why are carbon and hydrogen and oxygen considered to be different substances? What can we specifically point to that explains the difference between these elements on its most basic level?
The answer is the atomic number. If you look at a periodic table, you will notice that each element has a unique value between 1 and 118 which chemists call “atomic number.” Hydrogen has an atomic number of 1. Carbon has an atomic number of 6. Oxygen has an atomic number of 8.
On first blush, you might assume that an element’s atomic number is arbitrary. Since there exist no gaps in atomic numbers from 1 to 118, it’s easy to presume atomic number only serves as some convenient numerical identification. You might even think of some data organization advantages that such numerical identifications would offer. However, atomic number isn’t arbitrary; it says something fundamental about the subatomic structure of each element.
What is the Atomic Number?
In essence, an element is a type of atom. Atoms, themselves, are small round structures composed of what chemists call subatomic particles, namely protons, electrons, and neutrons. Different elements involve atoms with different numbers of these subatomic particles.
With this in mind, an element’s atomic number represents the number of protons found in one atom of the element. Thus, hydrogen atoms have 1 proton, carbon atoms have 6 protons, and so on.
Importantly, chemists use atomic number as the defining characteristic of an element. An atom can have any number of neutrons and electrons, but as long as it has 6 protons, chemists will always consider it a carbon atom.
With carbon specifically in mind, its atomic structure most often has 6 neutrons, though chemists know about other forms of carbon with 7, 8, or more neutrons. Variants of an element with different neutron numbers, and thus different atomic weights, are called “isotopes” of the element. Also, oxygen has 8 electrons in its elemental form, but can also have 10 electrons given certain conditions. Variants of an element with different electron numbers, and thus different electric charges, are called “ions” of an element.
Atomic Number and the Periodic Table
Since we know atoms can vary not just in proton number, but also in neutron and electron number, why do we care so much about protons? After all, chemists organize elements by atomic number in the periodic table, which suggests some inherent importance tied to an atom’s proton number. The answer lies in the chemistry of different elements.
In truth, chemists didn’t always use atomic numbers to categorize elements. Dmitri Mendeleev, the architect of the modern periodic table, arranged his first table in 1869 according to atomic mass instead. Because atomic mass essentially equals the sum of protons and neutrons, it correlates strongly with atomic number. Indeed, Mendeleev’s first periodic table arranges elements in a similar order to the modern table.
Why not use Atomic Weight?
However, some quick observations of the table revealed that ordering the elements by atomic mass proved unhelpful and misleading. First, some elements don’t have unique atomic masses. At the time of the table’s formulation, chemists estimated the atomic masses of nickel and cobalt to roughly equal each other. Non-unique atomic masses suggested that it was impossible to meaningfully order elements in this way.
Second, and more troubling, the chemical behavior of the elements undermined mass-based ordering. Chemists understood at the time that certain elements with very distant atomic masses had similar chemical characteristics. Fluorine, chlorine, bromine, and iodine each had diatomic elemental forms, as well as a strong propensity to exclusively ionize to a -1 charge. Sodium, potassium, rubidium, and cesium had incredibly unstable neutral states and only seemed to form +1 charged ions. Chemists also grouped carbon, silicon, and selenium, as well as nitrogen, phosphorus, and arsenic for their similar behavior.
It was the arrangement of this first group, termed “halogens,” that raised the eyebrows of chemists. Mendeleev had arranged the table so that these similar chemical groups shared the same row, including a row for these halogens. However, he knew that tellurium had similar chemical behavior to the oxygen group of elements. Tellurium has a heavier atomic weight than iodine, which forced Mendeleev to confusingly place it one space before iodine to maintain the chemical groupings.
Shortly after the publication of Mendeleev’s first table, it became clear that the table needed rearrangement.
The Power of Atomic Number
For more than half a century, chemists lived in an awkward space concerning the periodic table. On the one hand, they understood that Mendeleev’s 1869 table had flaws that necessitated a new model. On the other, no better model existed, and the atomic mass table still maintained most chemical groupings.
This changed in 1911 when Ernest Rutherford published the data from his famous gold foil experiment. Rutherford theorized that each atom had a nucleus of charged particles within a cloud of oppositely charged particles. Importantly, this meant that scientists could theoretically measure this nuclear charge. Chemists further theorized that the charge value of a given element corresponded to the number of particles called protons in the nucleus. In the succeeding decades, each nuclear charge, termed “atomic number”, was measured, providing an alternate way of ordering elements.
Thus, the modern form of the periodic table was formulated. Unlike the previous table, ordering by atomic number better maintains the arrangement of chemical groups.
Further, this modern iteration allows for the emergence of trends across the entire table. These trends are electronegativity, electron affinity, atomic radius, and ionization energy. Each trend has a direct relationship with the number of protons in each element. This results in each trend increasing or decreasing in intensity when nearing either the upper-right or lower-left corner of the table. To see these trends, check out our interactive periodic table.