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2.3: Atoms and Subatomic Particles

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  • Learning Objectives
    • Define atom and subatomic particle.
    • Describe the locations, charges, and relative masses of protons and electrons.
    • Determine the number of protons and electrons in an element.

    An atom is defined as the smallest unit of an element that still has the properties of that element.  For example, a piece of aluminum foil is shiny, silver in color, lightweight, and conductors heat and electricity.  If that piece of foil is torn in half, each of those halves would exhibit the same characteristics of the initial piece.  Each half could be divided over and over and over again, and each new fragment would still retain the same properties, until the atomic level is reached.

    The idea of an atom was first introduced by the Greek philosopher Democritus in 450 B.C.  However, his theories were largely-forgotten until the early 1800s, when John Dalton used the concept of an atom to explain why elements seemed to combine in whole-number ratios.  As mentioned in the first section of this chapter, these ratios are indicated by the subscripts of chemical formulas, and the derivation of these ratios will be discussed in greater detail in the next chapter.  As Dalton's theories became increasingly popular, additional scientists attempted to prove the existence of these small particles and were ultimately successful.  In the course of these studies, atoms were determined to be electrically-neutral, which means that they carry no overall charge, and were thought to be indivisible.

    Subatomic Particles

    While the former discovery was later proven to be correct, the latter was untrue.  Contrary to what Dalton and his contemporaries believed, atoms can, in fact, be broken apart into smaller units called subatomic particles.  Ultimately, three main types of subatomic particles have been discovered.


    Electrons, which were first discovered in 1897, are negatively-charged subatomic particles and are, therefore, symbolized using the notation "e."  In particular, every electron carries a −1 charge.  Electrons have incredibly small masses, but occupy the majority of an atom's volume.  Initially, scientists believed that electrons were tiny particles that were randomly-dispersed across a considerable volume, just as raindrops are little bits of water that are scattered throughout rain clouds.  However, this concept of an "electron cloud" was later proven to be inaccurate.  This theory will be revisited and corrected in a later section of this chapter. 

    Electrons are highly important, because a specific subset of electrons, called valence electrons, are solely-responsible for determining how elements interact, or bond, with one another.  The concept of bonding is the focus of Chapter 3 in this text.


    Protons, which were discovered in 1919, are subatomic particles that each bear a +1 charge and are, correspondingly, symbolized using the notation "p+."  Protons are 2,000 times more massive than electrons.  This ratio can be approximated by comparing the mass of a bowling ball to the mass of a penny.  However, despite their relatively large mass, protons occupy a very small percentage of an atom's volume.  The densely-packed space at the center of an atom in which protons are found is called the nucleus. The adage "opposites attract" can be used to explain why electrons remain anchored within atoms.  Since the central space within an atom is positively-charged, all negatively-charged electrons within that atom will be attracted to, and, therefore, bound within an atom by, the nucleus.

    Recall that the atomic number of an element is defined as the number of protons contained within an atom of that element.  Therefore, since atomic numbers are unique values, the identity of an element is solely-dependent on the number of protons present in an atom of that element.  For example, every atom of carbon, C, that exists in the known universe is defined to contain 6 protons, because its atomic number is 6, and no other element can contain exactly 6 protons.  Furthermore, it was stated above that atoms as a whole are electrically-neutral, but contain both electrons and protons, which are charged particles.  In order for all of this information to remain valid, the number of positively-charged protons and negatively-charged electrons in an atom must be equal, so that their combined charges effectively "cancel out" to a net zero, or neutral, charge.  Therefore, the atomic number of an element must indicate not only the number of protons found in an atom of that element, but also the number of electrons that are contained in an atom of that element.

    Example \(\PageIndex{1}\)

    Use a periodic table to determine the number of protons and the number of electrons contained in an atom of each of the following elements.

    1. Silicon
    2. Cd
    3. Bromine


    The number of protons in an atom is defined by the element's atomic number, which is found above the elemental symbol within a box on the periodic table. Furthermore, since an atom must have an overall neutral charge, the number of protons and electrons found within an atom of an element must be equal.

    1. Since silicon, Si, has an atomic number of 14, every silicon atom contains 14 protons and 14 electrons.
    2. Since Cd, cadmium, has an atomic number of 48, every cadmium atom contains 48 protons and 48 electrons.
    3. Since bromine, Br, has an atomic number of 35, every bromine atom contains 35 protons and 35 electrons.


    The final type of subatomic particle, the neutron, will be discussed in the next section.

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