All matter, life included, is made up of atoms. Here we discuss the inner workings of atoms.  This will be relevant as we work to understand the chemistry and biology of life.

Atomic Theory

Atomic theory provides a microscopic explanation of the many macroscopic properties of matter that you can directly test yourself in a chemistry laboratory. Every piece of matter we interact with, whether living or non-living, is made out of atoms.

The word atom comes from the Greek word for “indivisible.”  This naming is based on an idea held by some ancient philosophers—the Atomists, as discussed in the History of Astrobiology chapter—that all matter is composed of small, finite particles that differ in shape and size and that join together in different ways.

A modern understanding of atomic theory began in the early 1800s with John Dalton, an English schoolteacher. Through observations and experiments, Dalton demonstrated that atoms exist and correspond to different elements. Pure elements consist of only one type of atom, which has certain properties characteristic of that element.  As an example, Figure 1 shows a pure copper penny, which means the penny is entirely made up of copper atoms that are linked or bonded together.

Because an element consists of only one kind of atom, an element cannot be broken down into a different substance.  All atoms of a given element have identical chemical properties.  Inversely, atoms of a given element differ in properties from atoms of all other elements.

An atom is the smallest unit of an element that can participate in a chemical reaction. Atoms are neither created nor destroyed during a chemical reaction; they are instead rearranged to yield substances that are different from the substances present before the reaction (Figure 2). This property is captured in the law of conversation of matter.  Because the number and nature of atoms remains constant before and after a chemical reaction (it is only the configuration of atoms that changes), the total mass of matter before and after a chemical reaction also remains constant.

Going Subatomic

After establishing the existence of atoms, the inquiring mind is tempted to explore deeper: what are atoms themselves made of?

Discovery of Subatomic Particles

The first evidence of subatomic particles came in the late 1800s when different rays, meaning beams of light or radiation, were used to probe the structure of atoms.  Rays differ in energy level and source.  Some rays were found to emanate from certain elements, which we now call radioactive.

In 1897, the English physicist J. J. Thomson was the first to discover a subatomic particle from investigations of a type of ray called cathode rays.  Thomson found that cathode rays are made up of actual particles that are much smaller than atoms.  Experiments showed that these particles could be extracted from any atom, and that they carried a discrete amount of negative electric charge.  He named these particles electrons.

Soon thereafter, Marie Curie discovered radium, a radioactive element, in 1898.  Curie’s pioneering work in developing the theory of radiation was recognized by a Nobel Prize in Physics, and her work discovering the radioactive elements polonium and radium was awarded a Nobel Prize in Chemistry.  Marie Curie was the first woman to be awarded a Nobel Prize, and she is one of only four people (as well as the only woman) to win the Nobel Prize twice.

Ernest Rutherford, a physicist from New Zealand who largely spent his scientific career in Canada and England, showed that radium can be used to produce alpha (α) rays, beams of tiny, high-speed, positively charged particles called alpha particles.  Between 1908 and 1913, Rutherford directed a series of experiments using alpha rays in collaboration with, Hans Geiger, later famous for the Geiger counter, and Ernest Marsden, an undergraduate student at the time.

Geiger and Marsden aimed alpha rays at a very thin foil of gold and other various metals.  They found that alpha particles did one of three things: (1) most of the alpha particles passed right through the foil, (2) some particles were diverted slightly, and (3) a very small number of particles were deflected almost straight back towards the source of the alpha rays.  Rutherford described the behavior of the small number of particles being deflected back as follows:

It was quite the most incredible event that has ever happened to me in my life.  It was almost as incredible as if you fired a 15-inch shell at a piece of tissue paper and it came back and hit you¹

Let us examine what can be learned by the three different behaviors exhibited by alpha particles in this experiment.  Undeflected alpha particles were not affected by the foil at all, so they must have travelled through essentially empty space inside the atom.  Because most particles passed right through, this suggests that an atom largely consists of empty space.

We know that alpha particles are positively charged, and that like charges repel each other (like how the ends of magnets that have the same charge will repel each other rather than attract).  The deflected alpha particles therefore must have encountered another positive charge within an atom that repelled the alpha particles, causing their paths to be diverted or deflecting the particles straight back towards the source.  The smaller number of diverted and deflected particles suggests the existence of a nucleus, a small, positively charged body at the center of each atom.

These experiments led to an understanding of the atom as a small, positively charged nucleus, in which most of the mass of the atom is concentrated, surrounded by negatively charged electrons.  Because atoms are electrically neutral, the positive charge of the nucleus must be balanced by the negative charge of electrons.

After many more experiments, Rutherford established that the nuclei (the plural of nucleus) of all other elements contain the nucleus of a hydrogen atom.  He named this more fundamental “building block” particle a proton to refer to the positively charged, subatomic particle found in the nucleus of atoms.  We now know that the nucleus of hydrogen atoms consist of just a single proton.

The nucleus was known to contain almost all the mass of an atom, but it was found that the number of protons accounted for only half, or less, of the mass of the nucleus.  It was proposed that there existed neutral particles in the nucleus as well as the positively charged protons.  However, detecting uncharged particles is very challenging as they do not interact with charged particles.  It was not until 1932 that James Chadwick, a British physicist, found evidence of of these neutral particles now called neutrons.  Neutrons are uncharged, subatomic particles found in the nuclei of atoms with approximately the same mass as a proton.

Atomic Structure

We now understand that each atom is composed of a very small nucleus, made up of positively charged protons and uncharged neutrons, and that this nucleus is surrounded by a much larger volume of space in which negatively charged electrons exist.

The nucleus contains the majority of an atom’s mass because protons and neutrons are much heavier than electrons, whereas electrons occupy almost all of an atom’s volume. The diameter of an atom is on the order of 10⁻¹⁰ m, whereas the diameter of the nucleus is roughly 10⁻¹⁵ m—about 100,000 times smaller. To contextualize this, if an atom’s nucleus was the size of a blueberry, the diameter of the atom would be roughly the size of a football stadium (Figure 3).

Properties of Electrons, Protons, and Neutrons

Atoms—and the protons, neutrons, and electrons that compose them—are extremely small. For example, mass may typically be measured in terms of grams (g); a classic M&M is about 1 g.  The mass of a single carbon atom is less than 2 × 10⁻²³ g.  An electron is over a thousand times less massive, with a mass on the order of 1 × 10⁻²⁶ g.  Protons and electrons also carry electric charge in very small amounts that are unwieldy to express using more typical units.

When describing the properties of atoms and their subatomic particles, we therefore make use of appropriately small units of measure.  Rather than grams, we use the atomic mass unit (amu), which is approximately the mass of one proton.  For electric charge, we use the fundamental unit of charge (e), which is equivalent to the electric charge on a single proton.  In other words, a proton has a mass of ~1 amu and a charge of +1 e. Table 1 summarizes the properties of these subatomic particles.

For neutral atoms, atoms that carry no electric charge, the positive charge from protons must be balanced out by the negative charge of electrons.  For example, we know that a helium atom contains two protons.  A neutral helium atom must therefore also contain two electrons.  For neutral atoms, this means that the atomic number also corresponds to the number of electrons in the atom.

An atom’s mass number (A) gives the total number of protons and neutrons in the atom.  Recall that an element is defined by its atomic number (Z), the number of protons in its atoms neuclei, regardless of the number of neutrons.  We can use the mass number to calculate the number of neutrons in an atom as the difference between the mass number and the atomic number, i.e. A – Z = the number of neutrons.  For example, if a helium atom, which has two protons, has a mass number of 4, we can calculate that this helium atom has 2 neutrons (4-2).

Ions

Recall that neutral atoms contain the same number of positively charged protons and negatively charged electrons; the charges from the protons and electrons cancel each other out so that the atom has a net zero electric charge.  An ion is an atom that does carry an electric charge because there are either more electrons or less electrons than there are protons.  The electric charge of an atom is define as follows:

Atomic charge = number of protons − number of electrons

Atoms typically acquire charge by gaining or losing electrons. An atom that gains one or more electrons electrons will carry a negative charge.  These negatively charged ions are called anions.  Positively charged atoms, called cations, have lost one or more electrons.

As an example, consider the helium (He) atom characterized in Figure 4.  Recall that helium has an atomic number of 2, which indicates that all helium atoms have two protons.  A neutral helium atom therefore has two electrons.  If a helium atom gains one electron, it will have three electrons total (2+1) and become an anion with a charge of -1 e.  If a helium atom loses one electron, it will have one electron left (2-1) and become a cation with a charge of +1 e.  The helium atom in Figure 4 has a charge of +2 e, meaning it has lost both of its electrons.

Isotopes

The number of protons in an atom defines the type of element it is; if the number of protons in an atom changes, the type of atom it is also changes.  The number of electrons can change and defines the charge of the atom.  What about the number of neutrons?

During the early 1900s, it was discovered that an element could have atoms with different masses.  Despite the different masses of these atoms, they were otherwise indistinguishable from one another.  The difference in masses corresponded more closely to the mass of neutrons and protons, which have similar masses, than the mass of electrons, which are less massive by about a factor of 2,000 (see Table 1).  It was therefore deduced that these atoms have different masses because they possess different numbers of neutrons.

Such atoms are called isotopes—atoms of the same element that have differing amounts of neutrons.

An isotope of an element is specified by the mass number (A) written as a superscript to the left of the element symbol (see Figure 4).  For example, magnesium naturally occurs in the form of three different isotopes that have mass numbers of 24, 25, and 26.  These isotopes would be written as ²⁴Mg, ²⁵Mg, and ²⁶Mg respectively.  These isotope symbols are read “element, mass number.”  For instance, ²⁴Mg is read as “magnesium 24,” and can be written as “magnesium-24” or “Mg-24.”

Unstable Isotopes

Not all isotopes of an element are stable.  Unstable elements undergo a process known as radioactive decay, in which the subatomic particles in a an atom changes.  These unstable elements are described as radioactive.  As unstable nuclei decay, they often change from one isotope into another and sometimes into different elements depending on whether the number of neutrons or protons changes in the nucleus.  The rate of radioactive decay is commonly characterized by half-life, the amount of time in which approximately half the number of atoms will have decayed.

Let’s consider as an example uranium-238, which can also be written , meaning it has 92 protons and and 238-92=146 neutrons.  Uranium-238 has a half-life of 4.5 billion years.  If we had a sample of 12 uranium-238 atoms, then after 4.5 billion years we expect that 6 of those atoms will have undergone radioactive decay.   In other words, over 4.5 billion years, each uranium-238 atom has a half/half or 50% chance of decaying.

When uranium-238 decays, it emits an alpha particle; an alpha particle consists of two protons and two neutrons.  Because it gives rise to an alpha particle, this process is known as alpha decay and can produce alpha rays.

Recall that uranium-238 has 92 protons and and 238-92=146 neutrons.  With the emission of an alpha particle, a uranium-238 atom loses two protons and two neutrons.  This decay results in an atom with 90 protons and 144 neutrons.  Because the number of protons has changed, the nature of the atom has also changed.  An atom with 90 protons, and therefore an atomic number of 90, is a thorium (Th) atom.  The mass number of this atom is the combined number of protons and neutrons it has, 90+144=234.  Therefore, we see that when undergoing radioactive decay, a uranium-238 atom has decayed into a thorium-234 atom.

The process of radioactive decay is happening around us all the time.  If a radioactive nucleus decays into another radioactive nucleus, then another decay can happen.  This chain of reactions is called a “decay chain”. Eventually, the resulting nucleus may be stable and no longer decay. Most of the nuclei in the natural world appear to be stable, but there is a possibility that all nuclei will eventually decay if given enough time. We do not yet know if  “stable” nuclei are just extremely long lived.

Cosmic Abundance of Elements

After the Big Bang, there were only a few elements in existence: hydrogen, helium, and a sprinkling of lithium. As described in the section “Assembling the Periodic Table,” we know that the majority of the remaining elements, including most of those which make up the Earth, were formed at the cores of stars.  These elements were then distributed throughout the cosmos through their violent and explosive deaths.

We have used spectroscopy to identify the chemical composition of other stars, giant molecular clouds, and even the gas around and between galaxies.  These measurements allow us to calculate the abundance of various elements and isotopes.  We see that the universal abundance of elements match what is expected from the processes that create them.

These observations of abundances throughout the universe suggest that the laws governing chemistry and physics that we observe on Earth also operate the same way on any other planet at any other location in our galaxy and beyond. Therefore, if we can understand the chemical origins of life on Earth and in our solar system, we would gain insights into when, where, and how life might arise on other worlds.

Figure 5 shows the relative abundances of different elements in the Solar System. What are the most common elements?

Concept Check

All life on Earth is based on four elements: hydrogen (H), carbon (C), nitrogen (N) and oxygen (O). How common are each of these elements in the solar system, according to Figure 5?

Why might life not use the element helium (He) for life? Explain your answer!

Hydrogen, carbon, nitrogen, and oxygen form the basis of life, but how do we move from single atoms to complicated living things? To take the first step, we consider the chemistry of how atoms bond together to form molecules.

Review Questions

Summary Questions

1. What is an atom?
2. How are atoms relate to chemical elements?
3. What is the difference between an atom and a molecule?
4. What three “subatomic” particles are found inside an atom? Describe the similarities and differences between them.
5. How are different isotopes of a particular element distinguished? What are some reasons that certain isotopes are more common than others?
6. Which elements are the most common and which elements are the least common in our universe? How does astrophysics help explain this?

Activities

1. Building atoms. Open the “Build an atom” simulator (https://phet.colorado.edu/en/simulations/build-an-atom) and select the Atom option. Build the following:
   1. hydrogen, deuterium, tritium
   2. lithium (is this atom stable?)
   3. carbon, carbon-12, carbon-13