What do water (H2O), table salt (NaCl), and gold (Ag) all have in common? You may be thinking absolutely nothing, and if we were looking at them from a surface level, that might be true! But, when we dive into their chemical make-up, we find that all the atoms in these substances are held together by the same thing: intramolecular forces.
Intramolecular forces are the forces that hold atoms together within a molecule. The three types of intramolecular
forces are covalent bonds, ionic bonds, and metallic bonds.
This will be an overview of intramolecular forces and will specifically introduce the relationship between intramolecular forces and potential energy.
Potential Energy, in this context, refers to the energy stored in bonds.
- We will first explain polar and nonpolar covalent bonds
- Then, we will look at how the bond length of covalent bonds is related to potential energy
- Finally, we will briefly look at some examples of how potential energy and covalent bonds interact
You may have noticed that the three types of intramolecular forces are exactly the same thing as the three types of chemical bonds. The types of chemical bonds you are familiar with are indeed the forces that hold together the atoms in a molecule. Chemical bonds are referred to as intramolecular forces to differentiate them from intermolecular forces that exist between molecules. Check out Types of Chemical Bonds for more detail on the formation of covalent, ionic, and metallic bonds and Intermolecular Forces to learn more about the different types of intermolecular forces!
Intramolecular Forces: Nonpolar and Polar Covalent Bonds
Before we jump into the different types of covalent bonds, let's make sure we are on the same page on what a covalent bond is.
Covalent bonds occur between two atoms that have similar electronegativities and are formed by sharing electrons between atoms.
Not only is it essential to understand what a covalent bond is, but it is necessary to grasp how this sharing occurs. In a molecule, the negative charge of the shared electron pair is attracted to the positive charge of both atoms' nuclei.
In addition to the attraction between the electrons and nuclei, there are two other forces at play within a molecule:
- repulsion of the negatively charged electrons in each atom
- repulsion of the positively charged nuclei in each atom
A covalent bond between H2. StudySmarter Original.
Imagine you and a friend are sharing some ice cream one afternoon. If both of you really wanted some, you would likely each take an equal portion. If you were absolutely craving ice cream, you may go ahead and take more than your equal share (it’s okay we would all do it!). This equal or unequal sharing is the difference between nonpolar and polar covalent bonds. Even though covalent bonds consist of shared electrons, the atoms do not always share the electrons equally.
Nonpolar Covalent Bonds
Nonpolar covalent bonds are a type of chemical bond that forms when electrons are shared equally between atoms
In the ice cream scenario above, a nonpolar bond is represented if you and your friend were equally sharing the ice cream, except you and your friend would be two atoms, and the ice cream portions would be your electrons. This is usually the case in bonds between two of the same atom, such as H2 or N2. But why are the electrons shared equally?
This is because of a concept called electronegativity.
Electronegativity refers to the ability of an atom to attract electrons to itself. It exists on a scale and is a periodic trend.
When two atoms have similar electronegativities, this means that the attractive force of both atoms towards the shared electrons is equal, and so the electrons essentially stay in the middle and are equally shared, as seen above in H2.
Polar Covalent Bonds
Polar Covalent Bonds are a type of chemical bond that forms when electrons are shared unequally between atoms.
A polar covalent bond in a hydrogen chloride molecule. StudySmarter Original.
To summarize important differences, we will make a table to easily lay out the significant aspects of nonpolar and polar covalent bonds.
| Nonpolar Bonds | Polar Bonds |
| Sharing of Electrons | equal | unequal |
| Strength | weaker | stronger |
| Electronegativity difference | very small (>0.4) | larger (<0.4) |
| Common Examples | H2, N2, I2 | H2O |
To get a more detailed explanation of nonpolar and polar covalent bonds check out Polar and Non-Polar Covalent Bonds!
What is the Definition of Potential Energy?
Now that we've gotten a chance to learn about polar and nonpolar covalent bonds, we will define how potential energy plays a role in intramolecular forces (chemical bonds). Remember, potential energy is the energy stored in bonds. The amount of potential energy in a specific covalent bond changes depending on the bond length and the balance of attractive and repulsive forces between the atoms.
Atoms without full valence electron orbitals are in a high energy state and are trying to achieve lower energy through bonding to become more stable. The more unbalanced the repulsive and attractive forces are, and as we will see the further you are from the middle of a potential energy curve, the higher the potential energy will be in the atoms.
We will spend some time on this topic and detail what bond length is and how to think of it in terms of potential energy.
How Is Potential Energy In Intramolecular Forces Related To Bond Length?
When we think of potential energy in intramolecular forces, we can use an energy curve to picture potential energy as a function of bond length.
Bond length is the average distance between the two nuclei of atoms in a covalent bond.
Two main factors affect the bond length of a covalent bond:
- the type of covalent bond: single, double, or triple (this is referred to as bond order)
- if the bond order is the same, the size of the atom determines the length
Generally, the rule of thumb is that as bond order increases, the bond length decreases, and bond strength increases.
We will quickly explore the why behind the rule of thumb mentioned above.
Single, double, and triple covalent bonds refer to the number of shared electron pairs in each bond.
Single bonds = 1 shared pair
Double bonds = 2 shared pairs
Triple bonds = 3 shared pairs
So, with this information in mind, why does bond length decrease as bond order increases?
As the number of shared electrons increases in the bond orders, the attraction between the two atoms grows stronger, and they pull closer together which shortens the distance between them. (bond length)!
Following a similar line of thinking, as the bond order increases and the attraction between the atoms increases, this creates a stronger bond as it requires more energy to pull the closely drawn atoms apart.
You are probably thinking, okay, but HOW does all of this relate to potential energy? Well, that's where the energy diagram mentioned earlier comes in.
Potential Energy Diagram. StudySmarter Original.
Let’s break down what this potential energy curve is showing us. First, it displays the relationship between internuclear distance (bond length) and potential energy. As we can see, there are three different phases the two atoms can be in:
Suppose the atoms are very close, like the two leftmost atoms. In that case, they will have a very small internuclear distance and experience very strong repulsive forces between the two nuclei which is why the potential energy is so high.
As the internuclear distance increases, as seen in the middle of the diagram, the potential energy decreases. As the internuclear distance increases, the attractive force between the electron of each atom and the nuclei of the other atoms begins to balance out with the repulsive forces. Remember, two atoms are seeking lower potential energy through bonding. The internuclear distance at which the potential energy level is at its minimum is the bond length equilibrium and is where a bond will form.
The bond length equilibrium also correlates to the bond energy
If the atoms move further away than this ideal distance, then the internuclear distance is too far for any bonding. as attractive and repulsive forces are too far to interact. As you can see, the potential energy approaches zero.
Bond length equilibrium is the separation between atoms at which the potential energy is lowest
Bond energy is the amount of potential energy required to break the bond
A very broad overview to keep in mind when looking at energy diagrams is:
If the atoms are too close: repulsion occurs and high potential energy
If the atoms are at the equilibrium bond length: bonding occurs and low potential energy
If the atoms are too far: no interaction occurs
Potential Energy Examples
Finally, now that you should understand how potential energy, bond length, and intramolecular forces all relate, we are going to look at an example and turn the knowledge into work!
We will first make sure you understand how to analyze a potential energy diagram and the energy curve it shows you. Looking at the below diagram, what are the equilibrium bond length and potential energy needed to break the bond?
Potential energy curve. StudySmarter Original.
The first step to answering this question is making sure you know where the bond length and bond energy are located on an energy diagram.
Remember, the bond length is the intranuclear distance when the potential energy is the lowest, so let's look at the potential energy axis and see the intranuclear distance when the potential energy reaches its minimum.
We can make a safe estimate that the bond length is around 150 pm and the potential energy at that length is -300 kJ/mol which is equal to the energy needed to break the bond.
Using this diagram still, let's see how much you remember about the relationship between bond length and bond order, and strength! If we compare these to potential energy curves, which one correlates to a stronger bond?
Comparing two potential energy curves. StudySmarter Original. Again, the first thing to do is figure out the bond length of the two unknown molecules.
The red line one remains the same at about 150 pm.
Can you figure out the blue line? You are correct if you said about 200pm (or between 200pm and 250pm)!
Now, remember that bond length is inversely related to strength. As bond length increases, bond strength decreases. So which one would have a stronger bond? The molecule represented by the red line would!
The shorter bond length means it's harder to pull the atoms apart and thus it's a stronger bond.
Well, that brings us to an end and hopefully, by now you should be familiar with covalent bonds as a type of an intramolecular force and be able to identify and explain the bond length and potential energy of a molecule on an energy diagram. To get more detailed explanations and examples, check out Chemical Potential Energy Diagrams and Bond Length.
We've spent most of this introduction covering the interactions between covalent bonds. Intramolecular forces also include ionic bonds, so be sure to check out Coulomb's Law and Interaction Strength to learn about ionic bond strength and interactions!
Intramolecular Force and Potential Energy - Key takeaways
- Intramolecular forces occur within a molecular and the three types are covalent, ionic, and metallic bonds.
- The two types of covalent bonds are nonpolar and polar bonds. Nonpolar covalent bonds happen between atoms that equally share electrons, and polar covalent bonds occur when two atoms unequally share electrons between them.
- Potential energy is the energy stored in bonds, and it fluctuates depending on the intranuclear distance between atoms.
- Equilibrium bond length is the distance two atoms are at when potential energy is the lowest which means the bond is the most stable.
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