Everything around you is made of chemicals. And that’s only possible because those chemicals interact and bind together. Exactly how and why they do this depends on their nature but, in general, there are two kinds of interactions that keep them close: “primary” (or ‘strong’) and “secondary” (or weak) interactions.
These further break down into more subcategories, meaning there’s quite a lot of ground to cover. Today, we’ll be looking at the strong ones, which are formed through the transfer of electrons or electrostatic attraction between atoms.
As we go forward, keep in mind that atoms interact in order to reduce their energy levels. That’s what they get out of bonding to other chemicals, and they will do so until they find a bond-mate which will bring perfect balance to their lives; kinda like people do.
An atom’s stable configuration, the state all atoms tend towards, is known as its noble gas configuration. Noble gases make up the last column on the periodic table’s rightmost side, and they’re extremely or completely non-reactive chemically (they don’t need to interact because they have internal equilibrium).
Strong bonds are the most resilient ties atoms or molecules can forge with their peers. The secret to their strength comes from the fact that primary interactions are based on an atom’s valence. The valence number signifies how many electrons zipping around an atom’s core can be ‘shared’ with others. The overwhelming majority of a substance’s chemical behavior is a direct product of these electrons.
The first type of strong interactions we’ll look at, and the most common one, is the covalent bond. The name, “co-valence” sums up the process pretty well: two atoms share some or all of their valence electrons, which helps both get closer to equilibrium. This type of bond is represented with a line between two atoms. They can be single (one line), double (two lines), or triple (three lines).
In essence, what happens inside a covalent bond is that you have an atom starved of electrons (positively charged) and one who has too many electrons (negatively charged). Neither of them wants to keep going on like that because their internal imbalance of electrical charges makes them unstable. When put close to each other, they will start behaving like a single ‘macroatom’ — their electrons will start orbiting around both.
These shared orbits are what physically keeps the atoms together. The atom with too many electrons only ‘has’ them for half the time, and the one with too few gets to have enough half the time. It’s not ideal, but it’s good enough and it requires no changes to the structure of the atom (which is just dandy if you ask nature).
Things get a bit more complicated in reality. Electrons don’t zip around willy-nilly, but need to follow certain laws. These laws dictate what shape their orbits will take (forming ‘orbitals’), how many layers of orbitals there will be and how many electrons each can carry, what distance these orbitals will be from the nucleus, and so on. In general, because of their layered structure, only the top-most orbitals are involved in bonding (and as such, they’re the only ones giving elements their chemical properties). Keep in mind that orbitals can and do overlap, so exactly what ‘top-most’ means here is relative to the atom we’re discussing.
But to keep it short, covalent bonding involves atoms pooling together their free electrons and having them orbit around both, using each other’s weakness to make the pair stronger.
Covalent bonds are especially prevalent in organic chemistry, as it is the preferred way carbon bonds to other elements. The products they form can exist in a gas, liquid, or solid state, whereas the following two types can only produce solid substances.
Next are ionic bonds. Where covalent bonds involve two or more atoms sharing electrons, ionic bonds are more similar to donations. This type of chemical link is mediated by an electrostatic charge between atoms (negatively charged particles attract positively-charged ones). The link is formed by one or more electrons going from the donor to the receiver in a redox (oxidation-reduction) reaction; during this type of reaction, the atoms’ properties are changed, unlike in covalent bonds. Ionic bonds generally involve a metal and a nonmetal atom.
Table salt is a great example of a compound formed with ionic bonds. Salt is a combination of sodium and chlorine. The sodium atom will cede one of its electrons to the chlorine, which will make them hold different electrical charges; due to this charge, the atoms are then strongly drawn together.
It again ties into equilibrium. Due to the laws governing electron orbitals, there are certain configurations that are stable, and many others that are not. At the same time, atoms want to achieve electrostatic neutrality, as well. In an ionic bond, an atom will take an increase in its electrostatic energy (it will give or take negative charge) to lower its overall internal imbalance (by reaching a stable electron configuration) because that’s what lowers its energy the most.
Covalent bonds for the most part take place between atoms with the same electrostatic properties, and there’s no direct transfer of electrons because that would increase the overall energy levels of the system.
Ionic bonds are most common in inorganic chemistry, as they tend to form between atoms with very different electrostatic properties and (perhaps most importantly) ionic compounds are always soluble in water. However, ionic compounds such as salts do have a very important part to play in biology.
The main difference between ionic and covalent bonds is how the atoms involved act after they link up. In a covalent bond, they are specifically tied to their reaction mates. In an ionic bond, each atom is surrounded by swarms of atoms of opposite charge, but not linked to one of them in particular. Atoms with a positive charge are known as cations, while those with a negative charge are anions.
Another thing to note about ionic bonds is that they break if enough heat is applied — in molten salts, the ions are free to move away from each other. They also quickly break down in water, as the ions are more strongly attracted to these molecules than each other (this is why salt dissolves in water).
If the name didn’t give it away, this type of chemical bond is the hallmark of metal and metallic alloys. It’s not the only type of bond that they can form, even between pure metals, but it’s almost always seen in metals.
Chemically speaking, metals are electron donors — they need to shed electrons to reach equilibrium. Because of the nature of these atoms, their electrons can move around between atoms, forming ‘clouds’ of electrons. These detached electrons are referred to as being ‘delocalized’.
This type of bond shares properties of both ionic and covalent bonds. In essence, every metal atom needs to give away electrons to be stable (thus behaving like a cation). But because it’s surrounded by other metal atoms (meaning other cations), there’s nobody who wants to accept that electrical charge. So the electrons get pooled together and everyone gets to have them some of the time (thus forming a covalent bond). You can think of it as an ionic bond where the atomic nuclei form the cations and the electrons themselves the anions. Another way to look at it, although this is more of an abstraction used to illustrate a point, is that all the atoms involved in a metallic bond share an orbital.
Keep in mind that this ‘sea of electrons’ theory is a model of the process — it’s oversimplified and not a perfect representation of what’s actually going on, but it’s good enough to give you a general idea of how metallic bonds work.
Because metallic bonds share properties of both ionic and covalent bonds they create crystalline bonds (like salts) while still remaining malleable and ductile (unlike most other crystals). Most of the physical properties we seek in metals are a direct product of this structure. The cloud of delocalized electrons acts as a binder, holding the atoms together. It also acts as a cushion, preventing mechanical shock from fracturing the structure. When blacksmiths hammer iron or steel, they rearrange the atomic cores. Electrons can still move around them, like water around the rocks in a stream, and help hold everything together during the process.
Metallic bonds have the lowest bond energy of the types we’ve seen today — in other words, they’re the most stable.
Chemistry often gets a bad rep for being that boring subject with math and mixing of liquids. So it’s easy to forget that it literally holds the world together. The objects around us are a product of the way their atoms and molecules interact. Our knives can cut the food on our plates because billions of atoms inside that knife hold onto each other for dear life, and those in food don’t. Diamonds cut through solid stone because carbon atoms can bind to other carbon atoms in structures that are stronger than almost anything else we’ve ever seen. Our cells and tissues are held together by the same interactions. We’re alive because water molecules are shaped in such a way as to make them universal solvents.
We’re still very much working with models here — our understanding of the ties that bind is still imperfect. But even these models can help us appreciate the immense complexity hidden in the most mundane objects around us.