The Atom

Subatomic Particles

It was originally thought that atoms were the smallest unit of matter, the indivisible unit of our world. However, this was disproven upon the discovery of the proton, neutron, and electron. These three subatomic particles are the building blocks of atoms and contribute to different features of an atom.

Let’s compare the three particles now to get a sense of their properties.

Proton +1 e1.7 x 10-27kg
Neutron01.7 x 10-27kg
Electron-1 e9.1 x 10-31kg

Both the proton and neutron are much heavier than the electron and are therefore the main determiners of an atom’s weight. While only the proton and electron contribute to an atom’s charge. So the number of protons and neutrons will determine an atom’s weight while the balance of protons and electrons will determine its charge.

Plum Pudding Model

It was originally thought that all of these particles were floating in a gelatinous mass suspended in a sort of atomic pudding. The discovery of a small dense mass at the center of atoms quickly dispelled this belief.

The Nucleus

This central positively charged mass termed the nucleus contains all of the protons and neutrons in an atom. There protons and neutrons alike are held together by the strong nuclear force. At small distances, this force is able to overcome the electrostatic repulsion caused by the neighboring protons and keeps them from pushing one another away.

Like charges repel the only reason the nucleus can stay together is the strong force which at really close distances overcomes the electrostatic repulsion of protons.

Consequently, it takes a large sum of energy to bind these particles together. This binding energy accounts for the fact that atoms weigh less than we expect. An adequate explanation for this perplexing idea was one of Einstein’s profound contributions to the field of quantum mechanics and is summed up in his famous equation:


In simple terms, this equation means that mass can be converted to energy and vice versa. In the case of our nucleus, we end up losing a small fraction of the nucleus’ mass to the binding energy. We call this phenomena mass defect and it explains why atoms weigh less than their constituent particles.

The Electron Cloud

Surrounding this dense positively charged core are the electrons. Here they are held in a sort of orbit by the electrostatic attraction of the protons in the nucleus. They don’t fall into the nucleus because electrons don’t really exist as little floating balls, but as waves that spread out in space. Due to this, electrons take up a lot of space and contribute to the overall size of an atom while protons and neutrons contribute to its mass.

Putting them All Together

We can represent atoms as a whole by describing the number of protons, neutrons, and electrons each has by defining an atom’s mass number, atomic number, and atomic charge.

Mass Number

Mass number (A) is the sum of all of the protons and neutrons in the nucleus. It is an estimate of atomic mass, the actual weight of an individual atom, in atomic mass units. For example, the element carbon has a mass number of twelve. This indicates that the total number of both protons and neutrons in its nucleus is equal to twelve.

The mass number is displayed in two notations: the standard nuclear notation and isotope notation.

A = mass number, Z = # of protons, and C = atomic charge

The mass number and atomic mass don’t align perfectly because an atomic mass unit (amu or Da) is defined as 1/12 the weight of a neutral carbon-12 atom. This discrepancy occurs for two major reasons:

  1. Mass defects differ between different nuclei
  2. Neutrons actually weigh slightly more than protons

We have already discussed the effect of mass defect and we will look into number two a bit later when we discuss isotopes and their masses.

Atomic Number

The atomic number (Z) on the other hand describes only the number of protons in the nucleus. Carbon has an atomic number of six and therefore it has six protons in its nucleus.

Protons are the defining property of elements. So if we change the number of protons we end up with a different element. For example, if we were to give carbon an additional proton it would become nitrogen.

Atomic number is also displayed in standard nuclear notation and on the periodic table.

A = mass number, Z = # of protons, and C = atomic charge

By using both the atomic number and atomic weight we can determine the overall number of neutrons in an element by subtracting A from Z.

\[A=protons + neutrons\]
\[A-Z=(protons +neutrons)-protons = neutrons\]


While changing the number of protons fundamentally changes the identity of an atom, changing the number of neutrons doesn’t. We can add or subtract neutrons and still retain the original element. For instance, if we gave carbon an additional neutron it would still be carbon albeit a heavier version of the element. All versions of the same elements including the original are called isotopes and differ in weight and number of neutrons.

Due to this isotopes will have different mass numbers but the same atomic number. Remember that neutrons end up weighing a little bit more than protons so isotopic masses end up being larger than their mass number.

Atomic Weight

To make matters more confusing we can also describe an atom’s atomic weight. Since elements can exist in a wide array of different isotopes it doesn’t make sense to describe the mass of a sample of carbon by the atomic mass of carbon-12 alone. What if this sample has a mix of carbon-11, carbon-12, carbon-13, and carbon-14? Then using carbon-12 would misrepresent the weight of the sample.

Instead, we calculate atomic weight, a weighted average of all of these isotopes, in order to generate a more accurate representation of carbon as it exists in the “real world”. We use a weighted average because certain isotopes are more abundant than others. In the case of carbon, the only isotopes that exist in large enough quantities to be counted are carbon-12 (98.93%) and carbon-13 (1.07%).

Weighted Average: \[Avg_{weighted}\;=\;atomic \;mass_1(abundance_1) \;+ \;atomic \; mass_2(abundance_2)… \]
\[Atomic\; Weight_{Carbon}=12(0.9893)\;+\;13(0.0107)\;=\;12.01\]

This is the value that will be displayed in the periodic table and since we can’t place a single atom on a scale it will be measured in g/mol.

Carbon contains 12.01 g/mol

The Mole

A mole is a somewhat arbitrary unit that describes 6 x 1023 things, be it molecules, atoms, bats, or cats. In the case of g/mol for carbon, it tells us the weight in grams for 6 x 1023 carbons atoms. In most cases, it will be used as a conversion factor anytime questions ask for the quantities of atoms, molecules, photons, or other subatomic particles.

Mass number, Atomic Mass, and Atomic Weight???

It can be confusing to keep all of the different ways of measuring an atom’s mass straight. Let’s quickly overview them here to better understand how they are different.

Mass NumberTotal number of neutrons and protons an atom containsStandard Nuclear, Isotope
Atomic MassWeight of an individual atom in terms of amu or DaN/a
Atomic WeightWeighted average of the isotopic masses found in a representative sample of an elementPeriodic table

Atomic Charge

Unlike the various ways of representing the mass of an atom, there is only one way of measuring the charge, atomic charge (C). Atomic charge represents the balance between the number of protons and electrons in an atom. In a neutral atom, the number of protons and electrons are the same while charged atoms will have different numbers of these particles.

In positively charged atoms there are more protons than electrons and vice versa in negatively charged atoms. Since we can’t change the number of protons in an atom without changing its identity, the atomic charge of an element is controlled by the number of electrons it possesses.

These charge variants are called ions and are analogous to isotopes, but for charges instead of masses. We also have special names for positively and negatively charged atoms termed cations and anions respectively. So Na+ is the cation of sodium (Na), while Cl is the anion of chloride (Cl).

On standard atomic notation, the atomic charge will be placed to the right of the X and using the combination of the atomic number and atomic charge we can determine the number of electrons in an atom.

We will spend more time looking at electrons since they are pivotal in determining what characteristics different elements have, but for now, let’s look at how nuclei fall apart.

Term Summary

  • Proton: Positively charged particle found in the nucleus that determines the atomic number, weight, and identity of an element
  • Neutron: Neutrally charged particle found in the nucleus that determines the weight of an element
  • Electron: Negatively charged particle orbiting contributes to the overall charge of an atom
  • Nucleus: Small dense center of an atom containing protons and neutrons
  • Strong Force: Force that holds the nucleus together works over extremely small distances
  • Binding Energy: Amount of energy needed to bind the nucleus together to overcome the electrostatic repulsion between protons
  • Mass Defect: The amount of mass lost due to binding energy
  • Mass number (A): Number representing the sum of the protons and neutrons in the nucleus
  • Atomic Mass: Actual weight of an individual atom
  • Atomic Number (Z): Number of protons in the nucleus this is the number above each element in the periodic table
  • Isotope: An element with the same atomic number, but a different number of neutrons
  • Atomic Weight: Weighted average of an element’s isotopes represents the weight of a sample of an element as found naturally occurring
  • Mole: 6×1023 of something typically molecules or atoms though
  • Atomic Charge: The overall charge of an atom as determined by the balance between protons and electrons
  • Ion: A charged version of an element
  • Cation: A positively charged version of an element
  • Anion: A negatively charged version of an element