Atomic Structure on the MCAT

It’s amazing to think that all the things in the world — from animals, phones, and ourselves — are all built off of these small atoms that we can’t even see! To quantify the number of atoms that make up our whole planet, we’d probably need to be at least a type III civilization (mind blown…)!

The atomic structure of atoms not only tells us what atoms are composed of and their organization but also lays the foundation for how atoms interact with each other to form more molecules which in turn react with one another to form this expansive universe!

After going over this chapter overview, you should have the basics about atomic structure in relation to how it’s covered in general chemistry. Let’s get started!

Atomic Structure on the MCAT: What You Need to Know

Topics on general chemistry will be tested on the Chem/Phys and the Bio/Biochem section of the MCAT and can appear both as passage based and fundamental discrete questions. 

Try to expect around 2-4 questions that can appear on the MCAT covering stoichiometric topics, probably with a higher emphasis on calculation based questions!

Introductory general chemistry accounts for 30% of the content covered in the Chemical and Physical Foundations of Biological Systems and 5% of the content covered in the Biological and Biochemical Foundations of Living Systems.

Important Sub-Topics: Atomic Structure

A great way to approach the study of atomic structure is “build it brick by brick”. We say this because topics can get slightly more confusing and complicated when you get into electron configuration and different electron subshells. 

We’ll definitely employ this motif in the way we’ll organize and teach these topics! We’ll start off with the basic and easier to understand fundamentals and, slowly, but surely, introduce the more complex ideas and content.

1. The Subatomic Particles of the Atom

It’s amazing to believe that as small as atoms are, there are even smaller particles that constitute the main components of atoms, aptly called subatomic particles! There are 3 main subatomic particles: protons, neutrons, and electrons.

They all primarily differ in size, charge, and location within the atom: protons and neutrons are located in the atom’s nucleus, the atom’s dense center. Furthermore, protons are positively charged while neutrons retain a neutral charge.

It’s important to note that changing the number of protons changes the elemental atom! For example, increasing the number of protons from 4 to 5 changes the elemental atom from beryllium to boron!

Additionally, note that a neutral atom is one where there’s an equal number of protons and electrons — it doesn’t necessarily need to have the same number of neutrons!

2. Atomic Number v.s Mass Number v.s. Atomic Weight

It for sure can get a little confusing to differentiate between all these similar terms, but don’t worry! Similar to the advice we gave above, we’ll start from the most simply defined term and work our way up!

The atomic number simply refers to the number of protons within an elemental atom. The mass number is just a step up from atomic number, where we have to add both the amount of protons and neutrons.

Finally, the atomic weight refers to the average mass of all the naturally occurring isotopes of an element. To help you remember this, always think of the term “weighted average” to aid in your review!

It should as be noted, however, that you must also take into account the percentages of the naturally occurring isotopes, as they’re not all equal! Sometimes, certain isotopes come up in nature more often than others and this uneven distribution must be taken into consideration!

Click on this article for a more detailed coverage of the differences between all the terminology!

3. The Bohr Model and Atomic Spectra

Let’s now move our attention away from the nucleus and focus more on the electrons as the electrons essentially guide the reactivity of atoms and how molecules react with one another during chemical reactions!

In order to first understand how electrons are organized in the electron cloud, it may be helpful to first use a simplified model called the Bohr Model. In this model,  the electron cloud is represented by different principal shells, which are represented as circles — shells are where the electrons reside and travel.

As mentioned above, the lowest energy shell (n =1 ) is referred to as the ground state because it’s where an electron is lowest in energy. This will come back when we discuss atomic emission and absorption!

Finally, another important aspect to talk about which is easily visualized by the Bohr Model, is the atomic spectra. While electrons can only be located in one electron shell, they can also move between the shells!

In order to change between the electron shells, electrons must either gain or lose energy in the form of photons. Put in the simplest sense, photons are the basic units that makeup light, having a dual wave and particle nature. 

Though it has its flaws, the Bohr model is a great way to get introduced to the concepts of energy levels and atomic spectra. 

4. Understanding Quantum Numbers

While this is for sure one of the more tricky and complicated topics covering atomic structure, we’ll break it down bit by bit! The reason we use quantum numbers is to describe electrons in terms of their most probable location in an atom’s electron cloud — think of them as kinda like an address!

As mentioned in the last article, electrons are located in one of the principal electron shells, also called the principal quantum number (n). Furthermore, principal shells can be broken down into subshells, called the azimuthal quantum number (l).

When you increase the principal quantum number, the number of subshells increases as well! We can use the equation below to see which subshells are present within a principal electron shell:

Let’s give an example utilizing the n = 2 principal quantum number: given the relationship above, the accepted values for the azimuthal quantum number (l) are 0 and 1 (because 2 - 1 = 1). Each (l) value corresponds to a specific subshell as shown below: 

Going back to our example above, the n = 2 principal quantum number has 2 types of subshells: the s and p subshells. 

Additionally, this value also can tell you how many orbitals a specific subshell has. Take the p subshell. For example,  the p subshell (l = 1) has 3 possible magnetic quantum numbers (ml): -1, 0, and +1. Each of these 3 magnetic quantum numbers represents 3 different orbitals, each differing with orientations!

Because the spin of an electron can only be clockwise or counterclockwise, there are only 2 values that can be given for the electron spin, +½ and -½! This rule is highly important because orbitals can only hold up to 2 electrons — this explains why there are only 2 values for the electron spin quantum number!

Though it can be tricky, understanding quantum numbers takes a little time to grasp; however, we also encourage you not to get so focused and engulfed with understanding all the specifics!

5. Rules of Electron Configuration

We wanted to cover the topic of quantum numbers first because they lay the foundation for the electron configuration of elemental atoms! Simply put, electron configuration describes how electrons are distributed within the orbitals of the electron cloud.

A. Aufbau Principle

To understand this principle, it’s important to grasp that the different principal electron shells, as well as their subshells have different energy levels. 

As shown above, electrons must occupy orbitals from a lower energy level before occupying those at higher energy levels! You can use the arrow trick shown in the diagrams so that you don’t have to memorize the exact order!

B. Pauli Exclusion Principle

The best way to think about this principle is similar to a cell phone number or social security number: no 2 electrons can share the same set of quantum numbers.

Recall that the quantum numbers are used to describe the most likely location of an electron. This rule is a great segway into how we organize electrons when displaying the electron configuration of elemental atoms.

As shown above, we organize the orbitals to ascending energy levels in order to follow Aufbau’s principle. Likewise, notice also how each orbital only contains 2 electrons represented by 2 arrows: one pointing up (clockwise spin) and one pointing down (counterclockwise spin).

The electron configuration of neon (10 electrons) on the right side breaks the Pauli exclusion principle because the 2 box electrons have 2 identical sets of quantum numbers — this is shown by the 2 electrons having the same electron spin!

C. Hund’s Rule

Finally, Hund’s rule states that when assigning electrons to a subshell, they must occupy an empty orbital before pairing up. This is probably better understood when given a visual, as shown below!

These are some of the basic rules which serve as the foundations for how to construct the electron configurations of elemental atoms!

Important Definitions and Key Terms

Below are some high yield definitions and key terms to refer to when reviewing concepts and ideas about atomic structure in general chemistry!

Additional FAQs - Atomic Structure On the MCAT

Additional Reading Links (Coming Soon!) – Study Notes for Atomic Structure on the MCAT

Additional Reading: General Chemistry MCAT Topics:

• Acids and Bases on the MCAT
• Periodic Table on the MCAT
• Bonding and Chemical Reactions on the MCAT
• Chemical Kinetics on the MCAT
• Electrochemistry on the MCAT
• Equilibrium on the MCAT
• Solutions on the MCAT
• Stoichiometry on the MCAT
• The Gas Phase on the MCAT
• Thermochemistry on the MCAT
• Redox Reactions General Chemistry MCAT