Revision Notes — Periodic Classification of Elements

1. Why do we need a periodic classification?

The number of known elements grew rapidly in 19th century and scientists needed an organised system to understand patterns in properties. A classification arranges elements so similar properties appear at regular intervals (periodically). This helps predict properties of unknown elements and simplifies study.

2. Early attempts and Mendeleev’s contribution

Early attempts grouped elements by similar properties (Dobereiner's triads, Newlands’ law of octaves). Mendeleev (1869) arranged elements primarily by increasing atomic mass and left gaps for undiscovered elements — he predicted properties of gallium and germanium successfully. However, ordering only by mass caused exceptions (e.g., position of iodine and tellurium).

3. Modern periodic law

The modern periodic table is based on atomic number (Z), not atomic mass: properties of elements are periodic functions of their atomic numbers. This resolves anomalies in Mendeleev's table and aligns with understanding of electronic structure — atomic number equals number of protons and, for neutral atoms, electrons. The modern periodic table is arranged in periods (rows) and groups (columns).

4. Structure of the periodic table — periods and groups

Periods: Horizontal rows (period 1 to period 7). Elements in the same period show trends because they fill the same principal energy level (shell). Moving left → right across a period, nuclear charge increases by one for each successive element.

Groups (families): Vertical columns — elements in the same group have the same number of valence electrons which largely determines their chemical behaviour. For example, alkali metals (Group 1) have 1 valence electron and form +1 ions; halogens (Group 17) have 7 valence electrons and typically form −1 ions.

5. Electronic configuration and periodicity (simple explanation)

Periodic properties arise from electron configuration — particularly the outermost (valence) electrons. Elements in the same group have similar valence electron configurations, hence similar chemistry. E.g.:

• Group 1: ns¹
• Group 2: ns²
• Group 17: ns² np⁵
• Group 18: ns² np⁶ (inert gases; He is 1s²)

6. Periodic trends — what's changing and why?

Understanding trends uses two main ideas: effective nuclear charge (the net positive pull felt by valence electrons) and electron shielding (inner electrons partially cancel nuclear pull).

Atomic radius

Definition: Half the distance between nuclei of two bonded atoms of the same element.

• Across a period: Atomic radius decreases — increasing nuclear charge pulls electrons closer (shielding remains roughly same), reducing size.
• Down a group: Atomic radius increases — additional shells are added, outweighing the increased nuclear charge.

Ionization energy

Definition: Energy required to remove an electron from an isolated gaseous atom.

• Across a period: Ionization energy generally increases (atoms hold onto electrons more strongly due to higher effective nuclear charge).
• Down a group: Ionization energy decreases (outer electrons further from nucleus and better shielded — easier to remove).

Electronegativity / Electron affinity (qualitative)

Across a period electronegativity tends to increase (atoms more keen to attract electrons to complete valence shell); down a group it decreases.

Metallic and non-metallic character

Metals are generally found on the left and centre of the table and tend to lose electrons to form cations. Non-metals (right-hand side) tend to gain electrons to form anions. Across a period metallic character decreases; down a group metallic character increases (for many groups).

7. Important groups — short notes and important points

Group 1: Alkali metals (e.g., Li, Na, K)

• Have 1 valence electron (ns¹), show +1 oxidation state.
• Soft metals, low melting points (compared to transition metals), highly reactive — reactivity increases down the group.
• React with water to form hydroxides and hydrogen gas (vigorous for heavier alkalis).
• Common uses: Na in street-salt (NaCl), Na compounds in industry; Li in batteries.

Group 2: Alkaline earth metals (e.g., Be, Mg, Ca)

• Have 2 valence electrons (ns²), show +2 oxidation state.
• Less reactive than alkali metals; reactivity increases down the group.
• Important compounds: CaCO₃ (limestone), Mg for alloys, Ca in cement.

Group 17: Halogens (e.g., F, Cl, Br, I)

• Have 7 valence electrons (ns² np⁵) and typically form −1 ions (gain 1 electron).
• Highly reactive non-metals; reactivity decreases down the group (fluorine is most reactive).
• Exist in different physical states at room temperature: F and Cl gases, Br liquid, I solid.
• Applications: Chlorine in water purification, fluorine in toothpaste compounds, iodine as antiseptic.

Group 18: Noble gases (e.g., He, Ne, Ar)

• Have full valence shells (stable electron configurations) — largely inert (chemically unreactive).
• Used in lighting (Ne signs), inert gas shielding, helium in balloons and cryogenics.

8. Position of hydrogen

Hydrogen is unique — it can behave like an alkali metal (H⁺) or like a halogen (H⁻ in hydrides). In the periodic table it is usually placed separately at the top of Group 1 but should be treated as a special case in exams.

9. Trends illustrated with examples (quick calculations and reasoning)

Example 1 — Why does sodium have larger atomic radius than magnesium (same period)?

Answer: Sodium (Z = 11) has one electron in 3s and magnesium (Z = 12) has two electrons but higher nuclear charge pulls electrons more strongly — magnesium's radius is smaller.

Example 2 — Explain why chlorine is more reactive than bromine (both halogens).

Answer: Chlorine's valence shell is closer to the nucleus and less shielded; it more readily accepts an electron than bromine, making it more reactive.

10. Simple rules for predicting valency and formulae

• Valency for metals (left side) often equals number of electrons lost (Group 1 → +1, Group 2 → +2).
• Valency for non-metals may be 8 − (number of valence electrons) (e.g., O has 6 valence electrons → valency 2 → forms O²⁻, water H₂O).
• Combine oxidation states to write simple formulas: e.g., Na (Group 1) + Cl (Group 17) → NaCl; Mg (Group 2) + Cl (Group 17) → MgCl₂.

11. Common exam-style questions and how to answer

Typical questions ask to:

• Name periodic trends and give reasons (mention increasing nuclear charge, shielding effect).
• Predict valency and write chemical formulas — show valence electrons and their transfer/sharing.
• Explain positions of elements and compare properties within a group or period.
• Write short notes on important groups — be concise, include valency, reactivity trend and uses.

12. Worked examples (short)

Worked example — Predict formula: Aluminium (Group 13, valency 3) and oxygen (valency 2) combine → formula Al₂O₃. Reason: LCM of valencies 2 and 3 gives subscripts 3 and 2 respectively.

Worked example — Valency from electronic configuration: Element with configuration 2,8,1 (i.e., Na) has 1 valence electron → valency 1 → forms Na⁺.

13. Real-world relevance

Periodic trends explain why certain elements are used where they are — e.g., alkali metals for reducing agents, halogens for disinfection, noble gases for inert environments. Understanding trends is key to materials selection and chemical reactivity prediction in industry and everyday applications.

14. Revision checklist (quick)

• Know definition of periodic law (modern form) and why atomic number is central.
• Be able to explain and give one example for each trend: atomic radius, ionization energy, electronegativity, metallic character.
• Memorise key properties of Groups 1,2,17,18 and the special case of hydrogen.
• Practice writing formulas and predicting valency using valence electrons.
• Attempt NCERT exemplar and text questions related to this chapter for higher confidence.

15. Common mistakes to avoid

• Confusing atomic mass with atomic number — modern periodic table is based on atomic number (Z).
• Forgetting that trends are general — exceptions may exist; exam questions for Class 10 typically expect general trend reasoning.
• Ignoring units and significant figures in numerical questions; include units like pm or Å for atomic radius if asked.

16. Practice questions (brief)

1. Why does ionization energy generally increase across a period?
2. Predict formula for magnesium and oxygen compound and show valence electron transfer.
3. Compare reactivity of sodium and potassium and explain trend.
4. State two uses of noble gases.