per trillion!).
Figure 3.2 Schematic diagram of the three isotopes of carbon. The number of neutrons varies.
Carbon 12 (12C) is the most common form of carbon and constitutes more than 98% of carbon in living matter. 13C is rarer but, like 12C, it is a long-lived stable isotope. By contrast, 14C, with two additional neutrons, is not stable. It is called a radioactive isotope or radioisotope. One of the neutrons decays into a proton with the release of an electron and the atom becomes nitrogen 14 (14N). It transforms into a different element because it has now gained a proton which, you'll recall, defines the type of element. This decay has a half-life of 5730 years. In other words, after 5730 years, half of a sample of 14C will have decayed.
The unstable nature of 14C means that it is a small proportion of carbon isotopes but, despite this, it turns out to be enormously useful. As it decays with a known rate, it can be used to determine the age of ancient material from living things that contain carbon (“carbon dating”) such as bones. Living things constantly take up new 14C when they are alive, but once they are dead, they no longer actively take up carbon through metabolism. The 14C in the now dead organic matter begins to radioactively decay, allowing us to back-track and work out when that decay must have started and thus how old the object is. Later in the book, we see in more detail how radioisotopes can be used to put absolute dates on the fossil and geological record of Earth and other planetary bodies.
3.4 Electrons, Atoms, and Ions
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