Periodic Table Organization & Trends

How the Periodic Table Is Organized

The periodic table lists all known elements in order of increasing atomic number (the number of protons in the nucleus), not by mass. This single rule fixes every element's position. On the NY Regents Chemistry exam you receive a full Periodic Table of the Elements plus the Properties of Selected Elements table (the section the legacy edition labeled Table S), so most trend questions are lookups paired with a why.

Two directions matter:

• Groups are the vertical columns (numbered 1-18). Elements in a group have the same number of valence electrons (outermost-shell electrons), which gives them similar chemical properties.
• Periods are the horizontal rows (1-7). Across a period, electrons fill the same principal energy level (the same shell, labeled by quantum number n).

Key Families to Know

• Group 1 - Alkali metals (Li, Na, K...): 1 valence electron, very reactive, form +1 ions.
• Group 2 - Alkaline earth metals (Be, Mg, Ca...): 2 valence electrons, form +2 ions.
• Group 17 - Halogens (F, Cl, Br, I): 7 valence electrons, very reactive nonmetals, form -1 ions.
• Group 18 - Noble gases (He, Ne, Ar...): full valence shell (2 for He, 8 for the rest), nearly unreactive (inert).

Metals sit on the left and center, nonmetals on the upper right, and metalloids (B, Si, Ge, As, Sb, Te) along the staircase that separates them. The center block (Groups 3-12) holds the transition metals, which often form ions of more than one charge (for example iron as Fe2+ or Fe3+).

Why Group Position Predicts Behavior

The number of valence electrons sets an element's chemistry. Group 1 has 1, Group 2 has 2, and Groups 13-18 have 3 through 8 (count the second digit). Atoms react to reach a full octet, so Group 1 metals readily lose 1 electron and Group 17 nonmetals readily gain 1. This is why elements in the same group form the same kinds of ions and compounds. On the Regents exam, naming an element's group really tells you its valence-electron count and likely ion charge.

The Four Trends the Exam Tests

Four periodic trends appear again and again. Each follows from two competing factors: increasing nuclear charge (pulls electrons in) and increasing shielding from inner shells (pushes the trend the other way).

Atomic Radius

Atomic radius is the size of an atom. Across a period the nucleus gains protons while electrons stay in the same shell, so the stronger pull shrinks the atom. Down a group, each row adds a new energy level, so atoms get bigger. Example: Na is larger than Cl (same period), and K is larger than Na (same group, one shell lower).

Ionization Energy

First ionization energy is the energy needed to remove the most loosely held electron from a gaseous atom. Smaller, tightly held atoms (upper right) have high ionization energy; large atoms (lower left) lose electrons easily and have low ionization energy. The Properties of Selected Elements table (legacy Table S) lists these in kilojoules per mole (kJ/mol).

Electronegativity

Electronegativity measures how strongly an atom attracts the shared electrons in a bond. On the NYS Pauling scale in the Properties of Selected Elements table (legacy Table S), fluorine is the most electronegative element at 4.0; values fall as you move left and down. Electronegativity drives bond polarity, covered in a later section. As a rule of thumb, the top-right reactive nonmetals (F, O, N, Cl) have the highest values, and the bottom-left metals (Cs, Fr, K) the lowest.

Ionization energy and electronegativity move in the same direction across the table because both describe how tightly an atom holds electrons. Atomic radius moves the opposite way: small atoms hold electrons tightly (high ionization energy and electronegativity), while large atoms hold them loosely.

Metallic Character

Metals lose electrons readily, conduct electricity, and are malleable. Metallic character increases down and to the left - cesium and francium are the most metallic. Nonmetallic character increases up and to the right.

Worked Example: Comparing Two Atoms

Question: Which has the larger atomic radius, sulfur (S) or chlorine (Cl)?

1. Locate both: S is atomic number 16, Cl is 17 - both in Period 3.
2. Cl has the higher atomic number, so a stronger nuclear pull.
3. Same period, so no new shell is added.
4. Conclusion: S is larger because it has a weaker nuclear pull on the same shell. The Properties of Selected Elements table (legacy Table S) confirms it (S radius > Cl radius).

This read-the-table-then-explain pattern earns full credit on constructed-response items.

Ion Formation and Size

When an atom forms an ion, its size changes predictably:

• A metal forming a positive ion (cation) loses electrons and gets smaller than the neutral atom.
• A nonmetal forming a negative ion (anion) gains electrons and gets larger than the neutral atom.

For example, Na (about 190 pm on the NYS Properties of Selected Elements table) shrinks to Na+ (about 95 pm), while Cl (about 97 pm) grows to Cl- (about 181 pm). Regents items often ask you to compare an atom with its own ion, and this rule answers them directly.

Common Exam Traps

• Confusing groups and periods. Same group = similar reactivity; same period = same shell filling. Mixing them up reverses your trend.
• Assuming bigger atomic number means bigger atom. Across a period the opposite is true.
• Forgetting noble gases. Group 18 elements are listed without electronegativity on older tables because they rarely bond.
• Ignoring units. Ionization energy is kJ/mol and radius is in picometers; constructed-response questions may dock credit for missing units.
• Saying ionization energy increases down a group. It decreases, because outer electrons are farther from the nucleus and easier to remove.

Master these four trends and the two driving factors, and any comparison on the Properties of Selected Elements table (legacy Table S) becomes a quick, defensible answer.