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An atomic nucleus is surrounded by electrons, and they orbit at a great distance from the nucleus relative to the diameter of the nucleus itself. A typical atomic nucleus has a diameter of ∼10⁻¹⁴ m, whereas a whole atom has a radius of 2 × 10⁻¹⁰ to 5 × 10⁻¹⁰ m. Atoms have the same number of negatively charged electrons as the positively charged protons so that overall, an atom has no charge. It is neutral.

Electrons have something of a split personality. They exhibit particle-like properties and the behavior of waves, like light. Therefore, to consider electrons orbiting the nucleus in the same sense that a planet orbits a star is not quite technically correct. It is sufficient as a general description of an atom, and for simplicity they are often shown as tiny particles orbiting a nucleus, just as they are in the figures in this chapter. The dual wave and particle-like properties of electrons mean that they actually occupy fuzzy domains around the nucleus determined by probability distributions that, to add confusion, are called “orbitals.” The shapes of orbitals are determined by this wave – particle duality. The structure of electron orbitals will not be considered in detail in this book, but it is worth briefly spending some time pointing out their main features as they explain why atoms react at all and why they bond in particular ways. This is essential for grasping the fundamental atomic and molecular structure of life.

Each orbital can only take a maximum of two electrons. This is called the Pauli exclusion principle, which is rooted in quantum mechanics. Again, we do not need to explore the reasons for this in detail here, but we can take it as a fact to progress the discussion. As we move through higher atomic number elements in the Periodic Table, the electrons are stacked into additional orbitals, two by two.

The orbitals themselves are collected together into subshells that are given letter designations (s, p, d, f, g). Technically, these are electrons that share the same “angular momentum quantum number,” or the same orbital shape. These subshells themselves are put together to form shells with number designations: 1, 2, 3, etc. Technically, shells are electrons that share the same “principal quantum number.” In a crude way, you can think of shells as the “layers” of electrons as you move out from the inner layers of electrons (lower shell numbers) to the outer layers (higher shell numbers). Figure 3.3 shows this somewhat confusing nomenclature more clearly. If you want to understand these quantum constraints better, the further reading section provides suggestions. Here, the purpose is to draw out some key points.

Figure 3.3 A diagram showing the second shell in an atom and the nomenclature used to describe the various electron locations.

From the point of view of atomic structure, and the consequences for molecules involved in life, the crucial point to understand is that atoms have a tendency to react, losing or gaining electrons, until they have full electron shells. Once electron shells are full, then there are no “extra” electrons to get involved in chemical reactions or “missing” spaces for electrons that tend to be filled in chemical reactions. This explains the inert and stable characteristics of the noble gases, such as argon and neon, which have full electron shells. In a simplified sense, one can say that atoms tend to react until they achieve noble gas configurations of electrons. An understanding of this basic idea goes a long way to explaining how atoms behave and how molecules are formed.

Let's look at an example to illustrate this idea and how it applies to bonding. Consider the sodium atom (Na). It is in group 1 of the Periodic Table (see Appendix). It has 11 electrons (and therefore 11 protons – it has an atomic number of 11). Its electron structure is written as 1s² 2s² 2p⁶ 3s¹. The first number in this sequence is the shell number (in this atom there are three shells: 1, 2, and 3). Each letter (s and p) refers to a different subshell. The superscript shows the number of electrons in each subshell. Starting at the beginning, you can see that it has two electrons in shell 1, subshell s (1s²). This shell is full. Moving outwards in the layers of electrons, we then see that in the second shell, subshell s, it has two electrons (2s²). This is also full. In the second shell, we also have a p subshell. You will see that this has six electrons in it (2p⁶). The 2p subshell is made of three separate orbitals called x, y, and z of the same shape, each with a pair of electrons in them (they are full), giving the 2p subshell six electrons in total, hence 2p⁶. Finally, we come to the outermost shell, number 3, which has one lone electron in its s subshell (3s¹). This electron shell is not full. By losing this lone electron, the sodium atom becomes more stable because the next shell down is full. In other words, in an unscientific turn of phrase, sodium wants to lose this dangling spare electron to achieve a noble gas, stable configuration.

However, there are other consequences of losing this 3s¹ electron. In losing it, the sodium atom gains a net positive charge, as it now has 11 protons, but only 10 electrons, making a net positive charge of 1. The product of this electron loss, written as Na⁺, is called an ion. An ion is an atom that has gained or lost electrons.

To briefly illustrate this concept again on the other side of the Periodic Table, consider the element chlorine. Chlorine, in group 17 of the Periodic Table, has the electronic structure 1s² 2s² 2p⁶ 3s² 3p⁵. You can see that in its last electron subshell, 3p⁵, it has five electrons. It would like to gain one to arrive at six electrons in the 3p subshell and therefore fill the shell. By gaining an electron, it would attain a noble gas electron configuration, making it more stable. If it does this, however, it will now have 17 protons and 18 electrons, resulting in a net negative charge of 1. It will have become the ion Cl⁻.

The tendency of atoms to lose or gain electrons in this way to attain a noble gas electron configuration explains the key features of many bonds that we look at in the next section.

You will also notice that atoms have a tendency to lose or gain electrons in their outer electron shells since the ones below are full. The characteristics of the outer electron shells define the chemical behavior of different atoms. This explains why elements in the same group of the Periodic Table, for example carbon and silicon in group 14, which have the same electron configurations in their outer shells, tend to share similar chemical characteristics. This fact becomes important when we consider elements used in life in Chapter 4.

At this point, it is worth revisiting the question we brought up in Chapter 1 that puts an astrobiological perspective on this discussion. How much of what we have just discussed can be said to be a universal characteristic of life? I think you'd agree that we can confidently say that everything we have discussed is universal. If there is an alien intelligence somewhere else in the Universe, they would be drawing diagrams like Figure 3.3. We can say this because we know that the Periodic Table is universal since the Pauli exclusion principle that determines electron structure is universal. This is not conjecture. Using spectroscopy, which is discussed at the end of this chapter, we know that distant galaxies and stars, and therefore the planets they host, are made of the same elements as we have discovered in the Periodic Table. The fundamental chemical structure of atoms that make up all life in the Universe, if it exists elsewhere, is the same as on this planet.