The Periodic Table — How to Read It, Its Trends, and the Principles Behind It

When you first look at the periodic table, it’s hard to get a feel for why a chart packed with 118 elements has that particular shape. But the key idea is just one thing: elements in the same vertical column (group) behave similarly. Once you know this, you can read about 80% of the periodic table.

Try pointing to elements on the periodic table below. Same color = same category.

Why does the periodic table have that shape?

If you arrange elements in order of atomic number (number of protons), their chemical properties repeat at regular intervals. Lithium (Li) is a highly reactive metal, and sodium (Na), 8 spaces later, is also a highly reactive metal. Potassium (K), another 8 spaces later, is the same.

The periodic table is made by lining up these "elements with repeating similar properties" vertically. So the shape of the table wasn’t chosen arbitrarily — it follows a pattern found in nature.

What groups (vertical columns) tell you

A group is a vertical column. Elements in the same group behave similarly because they have the same number of outermost electrons (valence electrons).

Let’s look at what difference that makes:

• Group 1 (alkali metals): 1 valence electron → easily loses an electron → forms a $+1$ ion → reacts violently with water
• Group 17 (halogens): 7 valence electrons → tends to gain 1 electron → forms a $-1$ ion → highly reactive nonmetals
• Group 18 (noble gases): 8 valence electrons (octet complete) → no reason to give or take electrons → hardly react at all

If a test asks, "Why does sodium react so easily?", the answer is: "Because it is in Group 1, so it has 1 valence electron, and losing that electron gives it a stable electron configuration."

What periods (horizontal rows) tell you

A period is a horizontal row. Elements in the same period are filling the same electron shell (energy level).

Period 2 elements (from Li to Ne) all fill the second shell. Period 3 elements (from Na to Ar) fill the third shell.

Why does this matter? As you move to the right across the same period:

• One proton is added each step → the nucleus becomes more positively charged
• Electrons are added to the same shell → shielding barely increases
• Result: the nucleus pulls electrons in more strongly

This one principle explains all four major periodic trends.

Four periodic trends — all from the same principle

In the graph below, compare atomic radius and ionization energy across Period 3 (Na → Ar). You’ll see that they move in opposite directions.

Atomic radius: gets smaller to the right

When the nucleus pulls electrons more strongly, the electron cloud shrinks.

$\text{Li (152 pm)} \rightarrow \text{Be (112 pm)} \rightarrow \text{B (87 pm)} \rightarrow \cdots \rightarrow \text{F (64 pm)}$

On the other hand, as you move down a group, a new electron shell is added, so the atom gets larger.

Ionization energy: gets higher to the right

Ionization energy is "the energy needed to remove one electron." If the nucleus is holding electrons tightly, it’s harder to remove one.

So elements farther to the right have higher ionization energy, while elements lower down have lower ionization energy.

Test tip: It helps to remember this chain: "high ionization energy = hard to lose electrons = stronger nonmetallic character"

Electronegativity: gets higher to the right

Electronegativity is "how strongly an atom pulls shared electrons in a bond." The greater the nuclear charge and the smaller the atom, the more strongly it attracts electrons.

That’s why fluorine (F) has the highest electronegativity — a small atom with a strong nuclear charge.

Metallic character: gets weaker to the right

Metals are elements that lose electrons. Elements on the left lose electrons easily, so they have strong metallic character. Elements on the right tend to gain electrons, so they have stronger nonmetallic character.

To summarize:

$\text{같은 주기 오른쪽으로} \rightarrow \text{핵 전하 ↑} \rightarrow \begin{cases} \text{원자 크기 ↓} \\ \text{이온화 에너지 ↑} \\ \text{전기음성도 ↑} \\ \text{금속성 ↓} \end{cases}$

Don’t memorize all four separately. Just remember this one idea: "to the right = the nucleus holds electrons more strongly" and you can work out all the rest.

Reading it with an example: sodium vs. chlorine

If you compare sodium and chlorine side by side, you can really see the power of the periodic table.

Just from their positions on the periodic table, you can predict: "These two will react and form an ionic compound."

Common mistakes

"If the atomic number is bigger, the atom must also be bigger" — Not true. Within the same period, as atomic number increases, the atom actually gets smaller. Atomic size increases significantly when you move to a new period (a new shell).

"Noble gases don’t react because they have no electrons" — Noble gases do have electrons. They are unreactive because their outermost electron shell is full (an octet), which makes them stable.

"For transition metals too, group number = ion charge" — This works for main-group elements like Group 1 and Group 2, but transition metals (Groups 3–12) can have several possible ionic charges. For example, iron (Fe) forms both $\text{Fe}^{2+}$ and $\text{Fe}^{3+}$.

Check it yourself

1. Find lithium (Li), sodium (Na), and potassium (K) on the periodic table above. Same color, right? They behave similarly because they are in the same group (Group 1).
2. From sodium (Na) to argon (Ar) — as you move across the same Period 3, can you see the shift from metal → metalloid → nonmetal → noble gas?
3. Practice question: "Explain, using periodic table position, why magnesium (Mg) has a higher ionization energy than sodium (Na)."