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.
3.5 Types of Bonding in Matter
With this knowledge of atoms and ions, we now consider how they bond together to construct molecules and ultimately large, complex molecules or macromolecules, such as the genetic material DNA, that make up life.
There are five basic types of bonding that hold atoms and molecules together to give rise to the structure of ordinary matter, including life. We consider each of these in this chapter:
ionic bonding,
covalent bonding,
metallic bonding,
van der Waals interactions, and
hydrogen bonding.
The first three of these types of bonding are primarily involved in holding atoms together to make molecules, although ionic and covalent bonds also play prominent roles in holding parts of whole molecules together to generate three-dimensional structures. The last two types of bonding, van der Waals interactions and hydrogen bonding, are primarily involved in mediating interactions between individual molecules. Let's have a look at some of the features of these bonds and, in particular, how they are used in life.
3.6 Ionic Bonding
Ionic bonding is the electrostatic force of attraction between positively (+ve) and negatively (−ve) charged ions (primarily between non-metals such as chloride or fluoride ions and metals such as sodium or potassium ions). Most ionic compounds are crystalline solids at room temperature.
The crucial feature of an ionic bond is that each atom either gains or loses an electron so that the resulting ion has its lowest energy (noble gas-like) configuration. Table salt, NaCl, is a typical example of ionic bonding, and you can see its structure in Figure 3.4. In this salt, sodium gives up an electron, and chlorine gains this electron so that both ions gain a noble gas configuration, as we saw in Section 3.4. In other words, the Na atom has transferred its electron to the Cl atom, and the result is two ions, Na+ and Cl−.
Figure 3.4 The structure of NaCl showing the alternating sodium and chloride ions.
Source: Reproduced with permission of B. Blaus, https://commons.wikimedia.org/wiki/Category:Crystal_structure_of_sodium_chloride#/media/File:Blausen_0660_NaCl.png.
After transferring an electron, we now have two ions, Na+ and Cl−, with opposite charges. They are attracted to one another. However, if we have many of these ions, then things get more complicated. Clearly, negatively charged chloride ions will be repelled from each other, and positively charged sodium ions will be repelled from each other. If we place many Na+ and Cl− ions together, the natural configuration they take up to maximize attraction and minimize repulsion is an alternating packed cubic structure (Figure 3.4). Other similar examples are cesium chloride (CsCl) and sodium fluoride (NaF).
Ionic bonds are typically very strong. We can consider the stability of these bonds from an energetic (thermodynamic) point of view. Let's consider sodium fluoride, NaF. The energy required to break an ionic chemical bond in this structure is about 3 × 10−19 J. We can also calculate the typical thermal energy in a bond at a specified temperature. This is the thermal energy that would be in the bond when it is in equilibrium with a given environment at a particular temperature. This can be approximated by ∼kBT. kB is the Boltzmann constant, which has a value of 1.381 × 10−23 JK−1, and T is the temperature, measured in Kelvin. At room temperature (about 300 K), the value of kBT is 4.1 × 10−21 J. The ionic bond energy is therefore about 75 times the thermal energy in the bond. In other words, the thermal energy in NaF at room temperature will not cause the ions to break apart. You need to put in a lot more energy to break apart the ionic bonds, explaining why this compound is stable at room temperature. This is the case for many ionic bonds.
3.6.1 Ionic Bonds and Life
Ionic bonds are important for life because they can also play a role in holding different parts of molecules together.
In Figure 3.5, you can see a typical protein chain. A protein is made up of individual amino acids strung together in a long chain. Some of these amino acids are charged (we will come back to amino acids and discuss their features in more detail in Chapter 4).
Figure 3.5 A typical protein. The colored ribbons and lines depict the chains made of amino acids strung together, a chain that is itself folded together into a complex three-dimensional molecule. This one (called NOD2) is involved in the human immune system.
Source: Reproduced with permission of LPKozlowski, https://commons.wikimedia.org/wiki/File:Nod2_protein.png.
Some amino acids have a negatively charged side group, and some a positively charged side group. These two different types of amino acids can therefore come together to form an ionic bond. For example, the negative charge in an aspartic acid, which is one type of amino acid, is attracted to the positive charge in a lysine, another type of amino acid, to form an ionic bond, which keeps the chains together. Figure 3.6 illustrates this bonding.
Figure 3.6 An ionic bond in a protein formed between two amino acids, the positively charged amino acid lysine, and the negatively charged amino acid aspartic acid. Hydrogen bonding is also involved giving rise to both ionic and hydrogen bonding across the link – usually called a “salt bridge.”
This bonding is important in proteins that have functions to perform in cells, for example proteins that are enzymes or biological catalysts, that play a role in accelerating chemical reactions in life. These ionic bonds help the proteins keep their three-dimensional shape, which is necessary if they are to perform their functions correctly. From a very colloquial perspective, you can think of ionic bonds as “bolts” that can help hold together the three-dimensional structure of biological molecules, such as proteins.