Have you ever looked at a Periodic Table and thought, “How am I supposed to understand all these boxes and symbols?” Don’t worry, by the end of this, you’ll see just how cool and simple it actually is!
Table of Contents
What is the Periodic Table?
The periodic table is a table that shows all known chemical elements, organised in a smart way.
It was first arranged by a scientist named Dmitri Mendeleev in 1869. He noticed that when you line up the elements by their atomic number (that’s the number of protons), a pattern starts to show!
It’s like arranging students by age and noticing their behaviour changes with each year, that’s how elements behave in the table.
How is the Periodic Table Arranged?
- Horizontal rows are called Periods (there are 7).
- Vertical columns are called Groups or Families (there are 18).
Elements in the same group usually act alike. For example, Group 1 elements (like sodium and potassium) are all very reactive metals!
Let’s Explore the Periodic Trends (The Patterns You Should Know)
1. Atomic Radius – How Big is an Atom?
This refers to the size of an atom from its centre to the outermost electron.
- As you move across a period (left to right), atomic radius decreases.
Why? The number of protons increases, pulling the electrons closer! - As you move down a group, atomic radius increases.
Why? More shells (energy levels) are added.
Example:
- Lithium (Li) is larger than Fluorine (F)
- Cesium (Cs) is much larger than Sodium (Na)
2. Ionization Energy – How Hard is it to Remove an Electron?
This is the energy needed to remove an outer electron from an atom.
- Across a Period → Ionisation energy increases
(More protons = stronger hold on electrons) - Down a Group → Ionisation energy decreases
(Electrons are farther from the nucleus, easier to remove)
Example:
- It’s harder to remove an electron from Fluorine than from Sodium
3. Electronegativity – How Greedy is an Atom for Electrons?
This is how strongly an atom wants to attract electrons in a bond.
- ⬆️ It increases across a period (atoms want full outer shells!)
- ⬇️ It decreases down a group (big atoms = weaker pull)
Highest Electronegativity?
- Fluorine – the most “electron-hungry” element!
4. Metallic Character – How ‘Metal-Like’ is an Element?
This tells us how much an element behaves like a typical metal i.e. shiny, good conductor, malleable, etc.
- ➡️ Across a period: Decreases
- ⬇️ Down a group: Increases
Example:
- Francium (bottom left) is super metallic
- Fluorine (top right) is not metallic at all
5. Electron Affinity – How Happy is an Atom to Gain an Electron?
This is the amount of energy released when an atom gains an electron.
- It increases across a period
- It decreases down a group
Why? Atoms like Chlorine are close to a full outer shell, so they release energy happily when they gain an electron.
6. Valence Electrons – The Real MVPs of Bonding
Valence electrons are the electrons in the outermost shell of an atom. They are responsible for bonding and chemical reactions.
- All elements in a group have the same number of valence electrons.
Examples:
- Group 1 = 1 valence electron (e.g., Na)
- Group 17 = 7 valence electrons (e.g., Cl)
- Group 18 = 8 valence electrons (e.g., Ne) – That’s why noble gases don’t react!
7. Reactivity Trends – Which Elements React More Easily?
- Metals (Groups 1 & 2): Reactivity increases down the group
(Electrons are easier to lose!) - Nonmetals (Group 17): Reactivity decreases down the group
(Harder to gain electrons as atoms get bigger)
Example:
- Francium is more reactive than Lithium
- Fluorine is more reactive than Iodine
Sections of the Periodic Table You Should Know
Type of Element | Where to Find Them | Characteristics |
| Metals | Left and center | Shiny, good conductors, malleable |
| Nonmetals | Right side (except Hydrogen) | Dull, brittle, poor conductors |
| Metalloids | Stair-step line (e.g., Si, B) | Properties of both metals & nonmetals |
Groups You Should Remember
- Group 1 – Alkali Metals: Extremely reactive, soft metals
- Group 2 – Alkaline Earth Metals: Reactive, used in fireworks
- Group 17 – Halogens: Very reactive nonmetals (form salts!)
- Group 18 – Noble Gases: Non-reactive gases with full outer shells
Quick Recap
- Elements are arranged by atomic number
- Same group = same valence electrons
- Left to right = atoms get smaller and stronger
- Top to bottom = atoms get bigger and more metallic
- Trends help predict reactivity and bonding
Questions
- Which group contains non-reactive elements?
- What happens to atomic radius across a period?
- Which is more reactive: Lithium or Francium?
- What makes Fluorine so electronegative?
Understanding the Periods on the Periodic Table
Remember:
- Periods are the horizontal rows on the Periodic Table.
- There are 7 periods in total.
- The period number = number of electron shells (energy levels) in that element’s atoms.
Let’s go through them one by one:
Period 1 – The Simplest Row
- Elements: Hydrogen (H) and Helium (He)
- Shells: 1 energy level
- Special Notes:
- Only two elements here!
- Hydrogen is unique: It’s a nonmetal but placed with metals in Group 1 sometimes.
- Helium is a noble gas with a full shell (2 electrons), which makes it super stable!
Period 2 – The Start of Real Chemistry Fun
- Elements: Lithium (Li) to Neon (Ne)
- Shells: 2 energy levels
- Special Notes:
- Includes important elements like Carbon (C), Oxygen (O), and Nitrogen (N) — all essential for life
- Starts with metals on the left, ends with nonmetals and a noble gas
- First full period to show the metal → nonmetal transition
Period 3 – Similar to Period 2, Just Bigger
- Elements: Sodium (Na) to Argon (Ar)
- Shells: 3 energy levels
- Special Notes:
- Includes everyday elements like Sodium (in salt) and Silicon (in electronics)
- Again, moves from reactive metals to nonmetals to noble gas (Argon)
Period 4 – The Start of Transition Metals
- Elements: Potassium (K) to Krypton (Kr)
- Shells: 4 energy levels
- Special Notes:
- Begins to include transition metals like Iron (Fe), Copper (Cu), and Zinc (Zn)
- These metals are used in building, wiring, and making coins
- More elements = more complexity
Period 5 – More Transition Metals
- Elements: Rubidium (Rb) to Xenon (Xe)
- Shells: 5 energy levels
- Special Notes:
- Similar pattern to Period 4, but atoms are bigger
- Includes Silver (Ag) and Tin (Sn)
- Ends with Xenon, a noble gas used in lights and lasers
Period 6 – The Longest Period
- Elements: Cesium (Cs) to Radon (Rn)
- Shells: 6 energy levels
- Special Notes:
- Includes the Lanthanides (also called rare earth elements) — they’re usually shown below the table
- Contains Gold (Au), Mercury (Hg), Lead (Pb)
- Elements here are heavy and often used in electronics, medicine, and industry
Period 7 – The Radioactive Zone
- Elements: Francium (Fr) to Oganesson (Og)
- Shells: 7 energy levels
- Special Notes:
- Includes Actinides (also shown below the table)
- Many elements here are radioactive and man-made (synthetic)
- Used in nuclear energy, research, and sometimes cancer treatment
Summary Chart:
| Period | Number of Elements | Number of Shells | Key Features |
| 1 | 2 | 1 | Hydrogen and Helium only |
| 2 | 8 | 2 | Life-related nonmetals (C, O, N) |
| 3 | 8 | 3 | Sodium, Silicon, Chlorine |
| 4 | 18 | 4 | Starts transition metals (Fe, Cu) |
| 5 | 18 | 5 | Silver, Tin, more complex atoms |
| 6 | 32 | 6 | Includes Lanthanides, Gold, Lead |
Final Thoughts: Why Periods Matter
- The period number tells you how many electron shells an atom has
- As you go down the periods, atoms get bigger and sometimes more reactive (for metals)
- Each period helps predict how atoms bond, react, and behave
Read Also: The History Of Atomic Theory (A Fundamental Approach)
