ap chem study guides

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⚖️  Unit 7 - Equilibrium

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🤺  AP Chemistry Essentials

2.2 Intramolecular Force and Potential Energy



⏱️  4 min read

written by

anika p

Dalia Savy

dalia savy

August 23, 2020

You may be thinking, what are intramolecular forces?

Intramolecular forces are the ones between two atoms in a molecule! This is very different from intermolecular forces, which we learn in unit 3.

These two are confused often😕, but here are some tips💡:

  • Intermolecular are between molecules. Think inter = between two groups

  • Intramolecular are between two groups in a molecule, so think intra = within a group.

So far we learned about two intramolecular forces: covalent bonds and ionic bonds.

Types of Covalent Bonds

You must remember the following key information for the AP exam.

Physical or chemical processes can be described through energy diagrams. In molecules, bonds can either be classified as single, double, or triple bonds.

SINGLE = 2 electrons involved

DOUBLE = 4 electrons involved

TRIPLE = 6 electrons involved

An easy way to remember this is that each dash on a lewis dot diagram corresponds to two shared electrons.


Image Courtesy of Shodor

Also note that:

  • SINGLE bonds are the longest in length and have the smallest bond energy

  • TRIPLE bonds are the shortest in length and have the largest bond energy

Let's see this information in some graphics👀!

Bond Length and Bond Energy

For covalent bonds, bond length is influenced by the bond order (single, double, triple) and the balance between repulsive and attractive forces. Bond length is the physical distance between two atoms bonded to one another. Bond energy in the diagram shows how the greatest potential energy is the repulsion of two atoms. 


Image Courtesy of SplainScience

Let's take a look at each of these stages:

  • Repulsion - Since the atoms are very close together and the internuclear distance is very small, the atoms are experiencing lots of electron-electron repulsion. This is very unstable and leads to a potential energy of greater than 0.

  • Some overlap/attraction - This is the most stable state. There is a balance⚖️ between the repulsive and attractive forces and a stable bond is formed. This has been went over before but hopefully now you understand why potential energy is lowest when the bond is stable.

    • The potential energy at this stage is the amount required to break the bond.

  • No overlap/attraction - Since the internuclear distance is so large, there is no bond formed. This leads to a potential energy of almost 0.


It is good to understand these properties because you may be asked to guess where an element falls on this graph.

Say the following image is a diagram of chlorine atoms bonded together (Cl-Cl), where would Br-Br fall in comparison to chlorine's curve?


Image Courtesy of Chegg

To answer this question, we have to think about periodic trends:

  • Internuclear distance: is the Cl-Cl bond or Br-Br bond longer? Well, the one with the largest atomic radii would have to be the one with the longer bond. As you go down in a group, atomic radius increases. Therefore, Br-Br is longer than Cl-Cl.

    • This helps out with drawing the curve in relation to the x axis.

  • Potential energy: which bond would be easier to break? Cl-Cl or Br-Br? This should automatically make you think of ionization energy. The lower the ionization energy, the easier it would be to break the bond --> the less energy needed. As you go down a group, ionization energy decreases because there are more occupied electron shells. Therefore, Br-Br has a lower ionization energy and this bond is much easier to break.

    • This helps out with drawing the curve in relation to the y axis.

Knowing that Br-Br is longer and easier to break, you would have to graph its curve up (less energy) and to the right more (larger internuclear distance).

This question is a very good way to test your knowledge about this key topic and periodic trends. Here is what the graph should look like:


Forces Within Ionic Bonds

Understanding the strength of ionic interactions involves the use of Coulomb's law:

Energy of two interacting charged particles = (Q1*Q2)/d Where charges, Q1 and Q2, are separated by a distance, d

*Attraction occurs if the charges are opposite and repulsion occurs if the charges are the same*

  • Interaction strength INCREASES as distance DECREASES since those ions can be closer together, so smaller ions have stronger interactions

  • Larger charges on the interacting ions INCREASES interactions since interaction strength is proportional to charge

Coulomb's Law is everywhere!

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