🔹 1. What is the Periodic Table?
The Periodic Table is a systematic arrangement of all known elements based on their atomic number (number of protons).
It shows patterns in chemical properties and electron arrangements and is a useful tool for predicting element behavior.
🔹 2. History of the Periodic Table
- Dmitri Mendeleev (1869): First organized the table based on atomic mass and predicted missing elements.
- Modern Periodic Table: Organized by atomic number, not atomic mass, which corrected the order and explained the patterns better.
🔹 3. Structure of the Periodic Table
| Term | Explanation |
|---|---|
| Groups | Vertical columns (numbered 1 to 18); elements in the same group have same number of outer electrons and similar chemical properties. |
| Periods | Horizontal rows; elements in the same period have the same number of electron shells. |
| Metals | Found on the left and center of the table (Groups 1–12 mostly). |
| Non-metals | Found on the right side (Groups 14–18 mostly). |
| Metalloids | Elements with both metallic and non-metallic properties (e.g., Si, B, As). |
🔹 4. Group Trends and Periodic Trends
A. Group Trends
- Group I (Alkali Metals):
- 1 outer electron. Very reactive (especially with water). Reactivity increases down the group. Soft metals, low melting points, form +1 ions (e.g., Na⁺, K⁺)
- 2Na + 2H₂O → 2NaOH + H₂ (g)
- Group VII (Halogens):
- 7 outer electrons. Non-metals form diatomic molecules (e.g., Cl₂, Br₂). Reactivity decreases down the group Form -1 ions (e.g., Cl⁻, Br⁻)
- Cl₂ + 2KBr → 2KCl + Br₂
(Chlorine displaces bromine from salt)
- Group 0 or 18 (Noble Gases):
- Full outer shell
- Very unreactive
- Used in lighting (Ne), balloons (He), etc.
B. Period Trends
As you go left to right across a period:
- Atomic number increases
- Metallic → Non-metallic character
- Electronegativity increases
- Atomic size decreases
🔹 5. Transition Metals (Group 3 to 12)
- Hard, strong metals with high melting points
- Form coloured compounds
- Often used as catalysts
- Show variable oxidation states. Examples:
- Fe²⁺ (green), Fe³⁺ (yellow/brown)
- Cu²⁺ (blue)
🔹 6. Predicting Properties Using the Table
You can predict:
- Chemical reactivity: by group number
- Number of electrons: from atomic number
- Ion formation: Group 1 → +1, Group 17 → -1
- Physical properties: metals vs. non-metals
🔹 7. Metallic vs. Non-Metallic Elements
| Property | Metals | Non-Metals |
|---|---|---|
| Conductivity | Good | Poor (except graphite) |
| Appearance | Shiny | Dull |
| Malleability | Malleable, ductile | Brittle |
| State | Solid (except Hg) | Solid/liquid/gas |
| Type of Oxide | Basic oxides | Acidic oxides |
🔹 8. Examples for Practice
- Which group does magnesium belong to?
- Group 2 → 2 outer electrons
- Predict the formula of the compound between calcium (Group 2) and chlorine (Group 17):
- Ca²⁺ + 2Cl⁻ → CaCl₂
- Why is argon unreactive?
- It has a full outer shell of electrons.
- Order these elements by increasing reactivity: Li, Na, K
- Li < Na < K (down Group I, reactivity increases)
🔹 9. Summary Table of Key Groups
| Group | Name | Outer Electrons | Reactivity Trend | Ion Formed |
|---|---|---|---|---|
| 1 | Alkali Metals | 1 | Increases ↓ | +1 |
| 2 | Alkaline Earth | 2 | Increases ↓ | +2 |
| 17 | Halogens | 7 | Decreases ↓ | -1 |
| 18 | Noble Gases | 8 (except He=2) | Very low | None |
🧠 Tips for Exams
- Always link the electronic structure to the group and period.
- Know trends across a period and within a group.
- Be able to compare two elements based on their position.
- Practice displacement reactions and ionic compound formulas.
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