Bonds Basics
Up to this point, we have focused on exploring atoms and elements as independent entities. Pure elements and single atoms don’t really exist alone though. Instead, they are found bound together with other elements and atoms. This allows for an astounding amount of compounds with unique chemical properties that depend not only on what elements they are made of but also what types of bonds hold them together. In this lesson, we will focus on both bonds themselves and the properties they produce.
What Is A Bond?
Bonds are lasting electrostatic interactions that hold together two or more atoms, ions, or molecules. Bonds can result from the careful balance between attractive and repulsive forces that results in the sharing of electrons, as seen in covalent bonds. As well, bonds can result from the attraction of ions wherein unbalanced electronegativity values between atoms cause an electron to be “stolen”, as seen in ionic bonds. Such as those found between the halogens and alkali and alkaline earth metals. Regardless of bond type, it all comes down to electrostatics.
Why Do Things Bond?
Great, we know that things bond due to their electrostatic attraction to one another, but why do they bond? In the world, things want to be stable and they do their very best to achieve this. This is why nuclei undergo radioactive decay and why high energy electrons fall down to the lowest energy level they can. However, there is only so much an individual atom can do to achieve electronic stability.
Instead, they require the help of other atoms to fill up their shells and achieve their “goal” of becoming a noble gas, the most stable of all the elements. Sodium freely gives it electron to chlorine converting both of their electron configurations to those of the noble gases. In turn, this makes sodium positively charged and chlorine negatively charged, and thus they attract to form a bond. Most atoms are trying to do the same and as a result, they form bonds to fill out their shells.
Octets and Duets
In general, elements will follow the octet rule when forming bonds which states that elements with fewer than eight valence electrons in its shell will react to form more stable compounds. Typically, by forming one bond for every electron that it is missing.
For example, Carbon has an electron configuration of [He]2s22p2. Therefore it has four electrons in its valence shell and is missing four electrons to complete its octet. On the basis of this, we would expect carbon to form four bonds and in most circumstances it does.
One notable exception to the octet rule is hydrogen. This isn’t because it is a “rule-breaker” as we will discuss later, but because its valence shell only has two spots. Since it already has one valence electron it needs one more to fully fill its valence shell. As a result, hydrogen only forms one bond with other atoms.
The d-block does the same thing except the transition metals need 18 electrons for a fully filled shell. D subshells are a bit funky though so they don’t always obey the 18-electron rule and stable atoms have anywhere from 12-18 electrons in their valence shells.
| First Shell | p-block | d-block |
| Duet Rule | Octet Rule | 18-Electron Rule |
| s2 | s2p6 | d10s2p6 |
Rule Breakers
Some elements are said to break the octet rule and expand their octet. In doing so they end up forming more bonds than expected for their valence configuration. Take SF6 for example. Sulfur has a valence electron configuration of s2p4 so we would predict that it should form two bonds just like oxygen. However, here it forms six.
The reason this occurs is a bit complicated and the chemistry is still being worked out. Spoiler they aren’t using their d-orbitals. Regardless of the reason this occurs, these atoms aren’t breaking their octets just doing some fancy electron gymnastics with 3 electron bonds, ionic character, and the like.
Since this is an area of ongoing research and debate we only need to remember that some elements are going to “expand” their octets. The most common examples seen on the MCAT are phosphorous (P), sulfur (S), silicon (Si), and chlorine (Cl). So if you see them bonding to a bajillion other molecules don’t be alarmed.
Intramolecular Verus Intermolecular
As we go through the next lessons we will explore a myriad of bonds. From covalent to ionic and H-bond to dipole-dipole, it is important to understand whether the bonds we discuss are holding a molecule or multiple molecules together. This distinction defines the difference between intramolecular bonds and intermolecular bonds. Each of which defines the different properties of compounds and molecules.
Intramolecular Bonds
Intramolecular bonds are the bonds found within a molecule and hold its constituent atoms together. For instance, the bonds within a water (H2O) molecule hold oxygen and hydrogens together. Since they are within a molecule they are classified as intramolecular. These bonds come in two major flavors as alluded to above: covalent and ionic.
Intramolecular bonds determine molecular rigidity, molecular geometry, molecular polarity, among other properties.
Intermolecular Bonds
Intermolecular bonds on the other hand are found between molecules and hold multiple molecules together. Like international travel where we fly between different countries. For example, the H-bonds that hold multiple water molecules together are classified as intermolecular. These bonds come in four different types: London-dispersion, π-stacking, dipole-dipole, and hydrogen bonding.
Intermolecular bonds determine the melting point, boiling point, ability to interact with other molecules, overall structure, among others of molecules.