1. A quick history
Dmitri Mendeleev (1869) built the first widely accepted Periodic Table. He arranged elements by increasing atomic mass and grouped elements with similar properties into columns. He famously left gaps for undiscovered elements (gallium, germanium, scandium) and predicted their properties with remarkable accuracy.
Henry Moseley (1913) revised the table by reordering it by increasing atomic number (protons), which fixed a few inconsistencies in Mendeleev's mass-based ordering. This is the version we use today.
Why atomic number, not atomic mass?
A few elements break the mass order. The classic example: argon (atomic mass ~40, atomic number 18) actually comes before potassium (atomic mass ~39, atomic number 19) on the modern table. Ordering by atomic number (which is the count of protons that defines the element) keeps elements with similar properties aligned.
2. How the Periodic Table is organized
| Groups (or families) | The vertical columns (1–18). Elements in the same group have the same number of valence electrons and therefore similar chemical behavior. |
| Periods | The horizontal rows (1–7). The period number tells you how many energy levels (shells) the atom has occupied. Period 3 → 3 shells. |
Two predictive rules from position alone
- Same group → similar chemistry (because same number of valence electrons).
- Same period → same number of occupied energy levels.
3. The named families
| Group | Family name | Valence electrons | Key behavior |
|---|---|---|---|
| Group 1 (except H) |
Alkali metals | 1 | Most reactive metals. Soft, silvery, react violently with water. Lose 1 electron to form +1 ions. Li, Na, K, Rb, Cs, Fr. |
| Group 2 | Alkaline earth metals | 2 | Reactive metals (less than Group 1). Lose 2 electrons to form +2 ions. Be, Mg, Ca, Sr, Ba. |
| Groups 3–12 | Transition metals | varies | Hard, dense, often colored compounds. Multiple oxidation states (e.g. iron is Fe²⁺ or Fe³⁺). Fe, Cu, Zn, Ag, Au, etc. |
| Group 17 | Halogens | 7 | Most reactive nonmetals. Gain 1 electron to form −1 ions. F, Cl, Br, I, At. Exist as diatomic molecules (F₂, Cl₂…). |
| Group 18 | Noble gases | 8 (He has 2) | Full octet → essentially unreactive. He, Ne, Ar, Kr, Xe, Rn. Used as the reference for "stable" electron configurations. |
4. Metals, nonmetals, metalloids
On the Periodic Table, the "staircase" running roughly from boron (B) down to polonium (Po) divides metals from nonmetals.
| Metals (left of staircase) | Nonmetals (right of staircase) | |
|---|---|---|
| Appearance | Shiny, lustrous | Dull |
| Conductivity | Good conductors of heat and electricity | Poor conductors (insulators) |
| Malleability | Malleable (hammered into sheets) and ductile (drawn into wires) | Brittle (if solid) |
| State at STP | Solid (except Hg, which is liquid) | Mostly gas; some solid (C, S, P); Br is liquid |
| Ion behavior | Lose electrons → form cations (+) | Gain electrons → form anions (−) |
Metalloids
The elements touching the staircase — B, Si, Ge, As, Sb, Te — are metalloids (sometimes called semi-metals). They have properties intermediate between metals and nonmetals. Silicon and germanium are critical to semiconductor technology for exactly this reason: they conduct electricity, but not as well as a true metal.
5. The four periodic trends
| Property | Symbol / units | Across a period → | Down a group ↓ | Why |
|---|---|---|---|---|
| Atomic radius | r (pm) | decreases | increases | More protons pull harder (across); new shell added (down) |
| First ionization energy | IE₁ (kJ/mol) | increases | decreases | Tighter hold across; outer e⁻ farther & shielded down |
| Electronegativity | EN (Pauling, unitless) | increases | decreases | Same pull-on-shared-electrons logic as IE |
| Metallic character | — | decreases | increases | Opposite of IE/EN — metals want to lose electrons |
These four trends drive a huge fraction of Regents questions. Memorize the direction of each trend and you can answer questions about any pair of elements.
| Trend | Definition | Across a period (→) | Down a group (↓) |
|---|---|---|---|
| Atomic radius | Size of an atom (roughly, nucleus to outer electron) | Decreases | Increases |
| Ionization energy (IE) | Energy required to remove the most loosely held electron | Increases | Decreases |
| Electronegativity (EN) | How strongly an atom attracts electrons in a bond | Increases | Decreases |
| Metallic character | How "metal-like" an element behaves (easier to lose electrons) | Decreases | Increases |
The shortcut
Three of the four (radius, IE, EN) move in the same way: they reach maximum or minimum at the upper right (excluding noble gases) — except for atomic radius which is opposite.
- Smallest atom / highest IE / highest EN: upper right corner → fluorine.
- Largest atom / lowest IE / most metallic: lower left corner → francium (or cesium for practical purposes).
6. Why the trends exist (one explanation each)
Atomic radius decreases across a period
Across a period, electrons are being added to the same energy level while protons are being added to the nucleus. The nucleus pulls the electron cloud in more tightly. Result: the atom shrinks.
Atomic radius increases down a group
Down a group, each element has an additional occupied energy level. Even though the nucleus has more protons, those extra shells make the atom physically larger.
Ionization energy increases across, decreases down
Smaller atoms (across a period) hold their electrons more tightly, so it takes more energy to remove one. Larger atoms (down a group) have their valence electron farther from the nucleus and shielded by inner electrons, so it's easier to pull off.
Electronegativity follows ionization energy
Atoms that hold their own electrons tightly (high IE) are also good at attracting electrons in a bond (high EN). Fluorine has the highest electronegativity of any element (4.0 on the Pauling scale).
7. Using Reference Table S
Reference Table S gives you the first ionization energy, electronegativity, atomic radius, and melting/boiling points of every element. On exam day, you do not need to memorize numbers — you just need to know which trend to expect and look up the actual values.
How the Regents tests Table S
- "Which element has the highest IE: Li, Na, K, or Rb?" → Same group (1), top of the column has highest IE → Li.
- "Which element has the largest atomic radius: F, Cl, Br, or I?" → Same group (17), bottom has largest radius → I.
- "As you move from Na to Cl in period 3, atomic radius…" → decreases.
Key terms to know cold
| Group / family | Vertical column; same number of valence electrons |
| Period | Horizontal row; same number of occupied energy levels |
| Alkali metals | Group 1; very reactive; 1 valence electron; form +1 ions |
| Alkaline earth metals | Group 2; reactive; 2 valence electrons; form +2 ions |
| Transition metals | Groups 3–12; multiple oxidation states; often colored compounds |
| Halogens | Group 17; reactive nonmetals; gain 1 electron → −1 ions; diatomic |
| Noble gases | Group 18; full valence shell; unreactive |
| Metalloids | B, Si, Ge, As, Sb, Te — properties between metals and nonmetals |
| Ionization energy | Energy to remove the most loosely held electron |
| Electronegativity | An atom's pull on shared electrons in a bond |
| Atomic radius | Size of an atom from nucleus to outer electron |
| Metallic character | How easily an element loses electrons; high on the lower left |
Practice questions
Q1 · Multiple Choice
Which list of elements is arranged in order of increasing atomic radius?
- Li, Be, B, C
- F, Cl, Br, I
- Sr, Ca, Mg, Be
- N, O, F, Ne
Show answer
(2) F, Cl, Br, I. Going down Group 17 (halogens), each element has an additional occupied energy level, so atomic radius increases. Option (1) goes across period 2 (radius decreases). Option (3) goes up Group 2 (radius decreases). Option (4) goes across period 2 (radius decreases).
Q2 · Multiple Choice
As the elements of Group 1 are considered from top to bottom, the first ionization energy of each successive element
- decreases
- increases
- remains the same
- fluctuates randomly
Show answer
(1) decreases. Larger atoms hold their outer electron more loosely, so less energy is needed to remove it.
Q3 · Multiple Choice
An element that is malleable, ductile, and a good conductor of electricity is most likely classified as a
- metal
- nonmetal
- noble gas
- metalloid
Show answer
(1) metal. Malleability, ductility, and electrical conductivity are the three classic Regents-tested metallic properties.
Q4 · Part B-2 / Short response
Using Reference Table S, identify which element has a greater electronegativity: oxygen (O) or sulfur (S). Then explain in terms of atomic structure why that element has the greater electronegativity.
Show answer
Oxygen has the greater electronegativity (3.4 vs. 2.6 for sulfur).
Explanation: Both elements are in Group 16, but oxygen is in period 2 (2 occupied
energy levels) while sulfur is in period 3 (3 levels). Oxygen is the smaller atom, so its
nucleus is closer to the bonding electrons and attracts them more strongly. Sulfur's outer
electrons are farther from the nucleus and partially shielded by inner electrons, weakening
the pull.
Q5 · Part C / Extended response
Mendeleev organized his Periodic Table by increasing atomic mass and left gaps for elements he predicted would later be discovered. Moseley reorganized the table decades later. (a) State the basis Moseley used for organizing the elements. (b) Give one reason Moseley's ordering is more scientifically sound than Mendeleev's.
Show answer
(a) Moseley organized the elements by increasing atomic number
(number of protons).
(b) Atomic number is the defining property of an element — every atom of a
given element has the same number of protons. Atomic mass varies between isotopes and can
create ordering anomalies (e.g., Ar at mass ~40 comes before K at mass ~39 even though K has
more protons). Ordering by atomic number eliminates these inconsistencies and keeps elements
with similar chemistry properly aligned in their groups.