Organization

Walk into any chemistry classroom in the world and you will see the same chart on the wall. That universality is not an accident. The periodic table is the single most powerful organizational tool in all of science, and its layout is not arbitrary - every row, every column, and every block reflects something real about how electrons fill orbitals. Once you understand the architecture, you can extract an element’s electron configuration, predict its bonding behavior, and estimate its physical properties just from its position on the grid.

On MCAT test day, the periodic table is the only reference you get. No formulas, no constants, no conversion factors. Just the table. That makes this section essential: the better you read the table, the more “free answers” you pick up throughout the exam.

A Brief History: Mendeleev to Moseley

In 1869, Dmitri Mendeleev arranged the known elements in order of increasing atomic weight and noticed that chemical properties repeated at regular intervals. He called this the periodic law. Mendeleev was so confident in the pattern that he left gaps in his table and predicted the properties of elements that had not yet been discovered. When gallium, scandium, and germanium were later found with almost exactly the properties he predicted, the scientific community was convinced.

There was one problem. Arranging by atomic weight occasionally placed elements in groups where they did not belong. In 1913, Henry Moseley used X-ray spectroscopy to measure the nuclear charge (atomic number) of each element. His revision was simple but crucial: arrange the table by atomic number, not atomic weight. The modern periodic law states that the physical and chemical properties of the elements are periodic functions of their atomic numbers.

Periods (Rows)

The rows of the periodic table are called periods. There are seven periods, numbered 1 through 7 from top to bottom.

The period number corresponds to the principal quantum number (n) of the outermost electrons. Elements in Period 2 have their valence electrons in the n = 2 shell. Elements in Period 4 have their valence electrons in the n = 4 shell (though some inner d-orbitals fill during this period as well).

Each period gets longer as you move down because higher energy levels have more subshells available. Period 1 has only 2 elements (1s fills). Period 2 has 8 elements (2s and 2p fill). Period 3 has 8 elements (3s and 3p fill). Periods 4 and 5 each have 18 elements (s, d, and p fill). Periods 6 and 7 each have 32 elements (s, f, d, and p fill), though the f-block elements are traditionally pulled out and shown below the main table to keep it from being absurdly wide.

Groups (Columns)

The columns of the periodic table are called groups (or families). Elements in the same group have the same valence electron configuration, which is why they share similar chemical properties.

There are two numbering systems, and the MCAT may use either:

| System | Description | Example for Carbon’s Group |
|--------|-------------|---------------------------|
| IUPAC (1-18) | Numbers columns 1 through 18 from left to right | Group 14 |
| American (IA-VIIIA, IB-VIIIB) | A groups = representative elements, B groups = transition metals | Group IVA |

For representative elements (the A groups), the group number tells you exactly how many valence electrons the element has. Group 1A (or Group 1) elements have 1 valence electron. Group 6A (or Group 16) elements have 6 valence electrons. This is one of the most useful shortcuts the periodic table gives you.

The s, p, d, and f Blocks

The periodic table can be divided into four blocks based on which subshell is being filled by the element’s highest-energy electron.

| Block | Location | Subshell Filling | Groups (IUPAC) |
|-------|----------|-------------------|-----------------|
| s-block | Far left (2 columns) | s-orbital | 1-2 (plus He) |
| p-block | Far right (6 columns) | p-orbitals | 13-18 |
| d-block | Middle (10 columns) | d-orbitals | 3-12 |
| f-block | Bottom rows (14 columns) | f-orbitals | Lanthanides and actinides |

The widths of the blocks match the maximum number of electrons each subshell can hold: s holds 2 (2 columns), p holds 6 (6 columns), d holds 10 (10 columns), and f holds 14 (14 columns).

Representative vs. Nonrepresentative Elements

Representative elements (also called main-group elements) are those in the s-block and p-block - the A groups in the American system, or Groups 1-2 and 13-18 in the IUPAC system. Their chemistry is governed primarily by their s and p valence electrons, making their behavior highly predictable from group number alone.

Nonrepresentative elements include the transition metals (d-block) and inner transition metals (f-block). These elements can use d or f electrons in bonding, giving them variable oxidation states and more complex chemistry. You cannot simply read off the number of “active” electrons from the group number the way you can with representative elements.

Valence Electrons and Group Number

For representative elements, the connection between group number and valence electrons is your most valuable shortcut:

| Group (American) | Group (IUPAC) | Valence Electrons | Example |
|------------------|---------------|-------------------|---------|
| IA | 1 | 1 | Na: [Ne] 3s1 |
| IIA | 2 | 2 | Mg: [Ne] 3s2 |
| IIIA | 13 | 3 | Al: [Ne] 3s2 3p1 |
| IVA | 14 | 4 | C: [He] 2s2 2p2 |
| VA | 15 | 5 | N: [He] 2s2 2p3 |
| VIA | 16 | 6 | O: [He] 2s2 2p4 |
| VIIA | 17 | 7 | F: [He] 2s2 2p5 |
| VIIIA | 18 | 8 (2 for He) | Ne: [He] 2s2 2p6 |

This is why elements in the same group behave similarly. Sodium (Group 1, Period 3) and potassium (Group 1, Period 4) both have one valence electron. They both form +1 cations, react with water to produce hydrogen gas and a hydroxide, and form white ionic solids with halogens. The underlying electron arrangement drives the chemistry.

Why Same-Group Elements Behave Alike

Chemical reactivity depends overwhelmingly on valence electrons - the electrons in the outermost shell. Core electrons are buried deep, tightly bound, and shielded from the outside world. When two atoms interact, it is their valence electrons that overlap, transfer, or rearrange.

Since all elements in a group share the same valence configuration (just in progressively higher principal energy levels), they undergo the same types of reactions, form compounds with the same formulas, and display similar physical properties. The differences within a group - and there are important ones - come from the fact that atoms get larger as you move down, which changes how tightly the nucleus controls those valence electrons. Those differences are the periodic trends we will build in the rest of this chapter.