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Jai Ganesh

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Registered: 2005-06-28

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Atomic Number

Atomic Number

Gist

Atomic number is the number of a chemical element in the periodic system and on the periodic table that equals the number of protons in the nucleus of the atom. The elements are arranged on the table in order of increasing number of protons in the nucleus. Accordingly, the number of protons, which is always equal to the number of electrons in the neutral atom, is also the atomic number. An atom of iron has 26 protons in its nucleus; therefore, the atomic number of iron is 26 and its ranking on the periodic table of chemical elements is 26.

Summary

An atomic number is the number of protons contained within the nucleus of an atom. Each element on the periodic table has a different atomic number. Currently, the atomic numbers of known elements range from 1 to 118.

The nucleus of an atom contains protons and neutrons. An atom's electrons are found in orbitals around the nucleus. The atomic number definition is based on the number of protons in an atom. Protons are positively charged particles in the nucleus of atoms. The atomic number is based solely on the number of protons in an atom. In contrast, atomic mass is based on the number of protons and neutrons in an atom's nucleus.

The atomic number of carbon is 6 - meaning all atoms of carbon have 6 protons.

Periodic table entry for carbon indicating atomic number
Elements on the modern periodic table are organized sequentially by atomic number. The organization of the first versions of the periodic table was based on atomic mass. In 1869, a Russian scientist named Dmitri Mendeleev developed an early version of the periodic table organized by atomic mass. Atomic mass (or weight) indicates the average mass of one atom of an element expressed in atomic mass units (amus).

In 1869, Mendeleev organized his original periodic table by atomic mass. Scientists at the time did not know about the atomic structure of protons.

Over forty years later, in the early 20th century, the English researcher Henry Moseley confirmed the existence of protons and suggested that the periodic table be organized by proton number rather than atomic mass. He termed the number of protons in an element as its atomic number.

Details

The atomic number or nuclear charge number (symbol Z) of a chemical element is the charge number of an atomic nucleus. For ordinary nuclei composed of protons and neutrons, this is equal to the proton number (np) or the number of protons found in the nucleus of every atom of that element. The atomic number can be used to uniquely identify ordinary chemical elements. In an ordinary uncharged atom, the atomic number is also equal to the number of electrons.

For an ordinary atom which contains protons, neutrons and electrons, the sum of the atomic number Z and the neutron number N gives the atom's atomic mass number A. Since protons and neutrons have approximately the same mass (and the mass of the electrons is negligible for many purposes) and the mass defect of the nucleon binding is always small compared to the nucleon mass, the atomic mass of any atom, when expressed in daltons (making a quantity called the "relative isotopic mass"), is within 1% of the whole number A.

Atoms with the same atomic number but different neutron numbers, and hence different mass numbers, are known as isotopes. A little more than three-quarters of naturally occurring elements exist as a mixture of isotopes (see monoisotopic elements), and the average isotopic mass of an isotopic mixture for an element (called the relative atomic mass) in a defined environment on Earth determines the element's standard atomic weight. Historically, it was these atomic weights of elements (in comparison to hydrogen) that were the quantities measurable by chemists in the 19th century.

The conventional symbol Z comes from the German word Zahl 'number', which, before the modern synthesis of ideas from chemistry and physics, merely denoted an element's numerical place in the periodic table, whose order was then approximately, but not completely, consistent with the order of the elements by atomic weights. Only after 1915, with the suggestion and evidence that this Z number was also the nuclear charge and a physical characteristic of atoms, did the word Atomzahl (and its English equivalent atomic number) come into common use in this context.

The rules above do not always apply to exotic atoms which contain short-lived elementary particles other than protons, neutrons and electrons.

Additional Information

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 math 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.

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