Everything you see is made of tiny particles called atoms, but how are they bundled up together to make up individual objects? Well, we’ll be exploring the world of chemical bonds and intermolecular bounds. Let’s find out.
To explore the question above, we first need to note the hierarchy of how atoms become macroscopic objects. Atoms bond together into molecules, and molecules connect with each other to form different structures. In this section, we’ll be exploring chemical bonds, which are the processes that group atoms together into different types of molecules. For starters, are two common bonds from which molecules can form — the covalent bond and the ionic bond.
Both types of interaction deal with the number of valence electrons in the atom — that is, the number of electrons in the outermost layer. For an atom to be in a stable configuration, it needs to obey the octet rule — that is, every atom tends to have 8 valence electrons to fill up its outer shell. Chemical bonds, like covalent and ionic bonds, help these atoms maintain the stability of each other.
Covalent bonds occur when two atoms share their outer electrons. Simply put, some valence electrons from one atom are also outer electrons in the other, and vice versa. The diagram below might help:
In a covalent bond, two, four, or six electrons can be shared. This is known as a single, double, and triple bond, respectively. For example, the element of chlorine contains seven valence electrons, meaning it “wants” to get the last valence electron to become stable. It can do this with another chlorine atom, sharing two of its outer electrons (one from each) in a single bond. This is why you often see gases not as single atoms but as molecules of two atoms. For example, oxygen is often O2, nitrogen is often N2, and chlorine is usually Cl2. In particular, non-metallic elements usually form compounds using covalent bonds.
(put a diagram showing the covalent bonds)
Another way for atoms to stabilize themselves is through ionic bonds. Recall that an atom has no neutral charge — the negative charge of the electrons is equal to the positive charge of the protons. However, if an atom gains or loses an electron, that balance is broken — it becomes negatively and positively charged, respectively. This renders the atom an ion, and is key to the ionic bond.
In this type of chemical bond, one atom (A) gives off electron(s) to another atom (B). A becomes positively charged, and B becomes negatively charged. Electrostatic attraction forces bring A and B together, making the chemical AB.
These types of bonds occur pretty frequently, specifically in bonds between metals and non-metals. A good example is table salt (sodium chloride), consisting of one sodium atom and one chlorine atom. The sodium atom has one valence electron (which tends to be given away), and the chlorine atom has seven valence electrons (so it tends to gain one more electron). Thus, the extra electron in the sodium atom is transferred to the chlorine atom, and both ions attract together to form sodium chloride.
Apart from covalent and ionic bonds, there is also one special type of bond that involves metals — called the metallic bond. In this case, the outer electrons in the atoms escape from their nuclei. These delocalized electrons flow freely between the positively charged metal ions. Electrostatic attraction keeps the electrons and ions in place and helps form a strong metallic bond between the metal atoms. In fact, that’s why metals have such a high melting and boiling point compared to non-metals. This is also why conductivity works so well in metals, as the delocalized electrons can easily become current carriers.
After we explain the forces that hold atoms together into molecules, it’s time to discuss how molecules hold together into structures. One of the common ways for this to happen is through ion-dipole interactions. The charge distribution of an ionic compound is not even — one side is (relatively) positively charged while the other is negatively charged. This essentially makes it a dipole, in which a positive charge and a negative charge are separated by a distance. This allows dipole-dipole interaction to occur, as electrostatic forces mean opposite sides of the dipoles attract each other.
However, these dipole bonds are not limited to ionic compounds. Instead, all they need is that they are polar, where the charge is not uniform and different on each side. This could work for covalent molecules as well, such as water (H2O) and ammonia (NH3). Specifically, the bond holding these molecules together is the hydrogen bond, involving a positive hydrogen ion (a proton) and a negative ion (the atom that the hydrogen atoms are bonding to). For example, in water, the hydrogen ion (that loses one electron) attracts the surrounding oxygen ions (that gain two electrons). This creates a strong intermolecular bond known as a hydrogen bond.
In this article, we’ve discussed how atoms bond together into molecules, and how separate molecules can create structures. Even though most of the forces behind this might be easy to understand, it’s very hard to predict how a structure looks based on what it’s made of. Every atom in the structure affects the forces and positions of other atoms, so might be challenging to search for a stable state without extensive simulation or analysis. If you have any questions or suggestions, please put them down in the comments below.