How to Understand Periodic Trends
- by Fehbe Meza
- Oct 22, 2019
- MCAT Blog, MCAT Chemistry
I can recall how dreadful it was to enter the MCAT testing room with simply an identification card and noise-canceling headphones, questioning whether you have the right tools to tackle the exam or not. Luckily, there are various resources provided during the exam that can guide you towards the right answers. An immensely useful tool is the periodic table for the MCAT. Even though the MCAT rarely includes questions asking directly about periodic properties, knowing the periodic trends is fundamental to answering many questions in the science sections. So we’re focusing on explaining periodic trends and discussing variations in the atomic radius, ionic radius, ionization energy, electron affinity, and electronegativity.
If you need a more in-depth look at the periodic trends for the MCAT, Next Step Online MCAT Course students can attend Office Hour sessions specifically about this topic.
Atomic and Ionic Radii
Even though it is not possible to specify a definite value for the radius of an isolated atom, the atomic radius can be thought of as half the single-bond length of a homonuclear bond and atomic radii variations are closely related to the concept of effective nuclear charge. Recall that the effective nuclear charge is the net positive charge experienced by valence electrons or the difference between the number of protons and the shielding electrons located in the energy levels between the nucleus and the valence electrons. Visualize these shielding electrons as the ones that guard and “shield” the valence electrons from the attractive force of the positive nucleus. Strong “shielding” will not allow the valence electrons to be strongly pulled by the nucleus and they will be free to spread out, thus making the atomic radius bigger. Conversely, the greater the effective nuclear charge, the smaller the atomic radius.
When studying the periodic table for the MCAT, remember that shifting from left to right across the same period will increase the number of protons in the nucleus while keeping the number of shielding electrons constant since only the number of valence electrons increases across the same period. The effective nuclear charge is greater towards the right side and that is the reason why the atomic radius decreases as we shift from left to right.
The variation in atomic size of the elements in the same group can also be explained by the concept of electron shielding since going down a column increases the principal quantum number n and the number of orbitals between the nucleus and the valence electrons. The added inner shells are effective at shielding valence electrons which allows for distinct increases in size. Therefore, the atomic radius increases from top to bottom down a group.
The concept of effective nuclear charge and comparisons of ionic radii with atomic radii of the same element show that cations are always smaller than the parent atom since they are formed by losing electrons and anions are always larger than the parent atom since they are formed by gaining electrons.
The periodic table also helps us predict the ionization energy which is the energy required to remove electrons from the highest-energy orbital of neutral atoms. In order to remove electrons, work must be done to overcome the electrostatic attraction between the nucleus and the electrons; thus, the ionization energy increases as the nuclear charge increases from left to right across the periodic table. On the other hand, the ionization energy decreases while going down a column because electrons are removed from shells that are further from the nucleus and its attractive force.
This trend can also be explained in terms of electrostatic forces and Coulomb`s Law: . As represented by this formula, the electrostatic force of the nucleus is inversely proportional to the radius squared. As radius increases, the electrostatic force decreases; thus, the ionization energy decreases as well. It is important to know that an atom has as many ionization energies as it has electrons and the second or third ionization energy is always bigger than the second or first, respectively. In addition, noble gases have the highest ionization energies due to their completed outer shells.
While ionization energy refers to the formation of positive ions, electron affinity refers to the formation of negative ions by adding electrons to the lowest vacant energy orbitals. Since the added electron will be attracted by the nucleus, this is generally an exothermic process that releases energy. Even though the electron affinity trends are not as clear as the ones for ionization energy or atomic radius, electron affinities become more exothermic as we shift from left to right across a period and less exothermic as we go down a group from top to bottom.
An electronegative atom is analogous to a child who strongly pulls on their toys and is not willing to share them. Electronegativity measures the tendency of atoms to pull and form bonds with electrons; thus, it is helpful to think of electronegativity trends relative to the atomic radius and the attractive force of the nucleus. An atom with a smaller atomic radius exerts a great electrostatic force on its valence electrons and is not willing to share them with other atoms in a covalent bond; therefore, it has a higher electronegativity than atoms with larger radiuses who cannot exert the same attractive force on their outer electrons. As we shift from left to right across the periodic table, the radius decreases and the atoms have a higher tendency to pull electrons into their valence shell. Conversely, as the radius increases from top to bottom down a column, it becomes harder for the nucleus to keep hold of electrons. Thus, electronegativity increases from left to right and decreases from top to bottom, with the exception of noble gases, lanthanides, and actinides which do not possess electronegativity values. It is important to note that Fluorine, which sits on the top right corner of the periodic table, is the most electronegative element.
The periodic table and periodic trends might seem overwhelming, but with a little review, you’ll be able to understand their complexities. If you need additional help, our online MCAT PREP covers periodic elements in greater detail. You might also consider MCAT Tutoring for one-on-one learning. Unable to decide between the two? Schedule a free consultation with an Academic Manager to sort your through options and find the MCAT prep style that’s right for you.
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