Chapter Overview
Classification of Elements and Periodicity in Properties is the third chapter of CBSE Class 11 Chemistry. It tells the story of how chemists organized the elements — from Dobereiner's triads and Newlands' octaves to Mendeleev's breakthrough and the modern periodic table. More importantly, it explains why elements behave the way they do through periodic trends in atomic radius, ionization enthalpy, electron gain enthalpy, electronegativity, and metallic character.
Understanding periodicity is essential because it allows you to predict properties of unfamiliar elements and explain chemical behaviour systematically. Every trend in this chapter connects directly to atomic structure (Chapter 2) and chemical bonding (Chapter 4).
1. Historical Development of the Periodic Table
1.1 Dobereiner's Triads (1829)
- Grouped elements in sets of three (triads) where the atomic mass of the middle element was approximately the arithmetic mean of the other two.
| Triad | Element | Atomic Mass |
|---|---|---|
| Li | 7 | |
| Na | 23 | |
| K | 39 | |
| Ca | 40 | |
| Sr | 88 | |
| Ba | 137 |
- Limitation: Only a few triads were identified; the model could not accommodate all known elements.
1.2 Newlands' Law of Octaves (1865)
- When elements are arranged in increasing order of atomic mass, every eighth element has properties similar to the first (like musical octaves).
| Element | Li | Be | B | C | N | O | F |
|---|---|---|---|---|---|---|---|
| Similar to | Na | Mg | Al | Si | P | S | Cl |
- Limitations: (1) Worked only up to calcium (20 elements). (2) Placed iron (a metal) with cobalt and nickel (different properties). (3) Only 56 elements were known; Newlands assumed no more would be discovered.
1.3 Mendeleev's Periodic Table (1869)
Mendeleev's Periodic Law: The physical and chemical properties of elements are a periodic function of their atomic masses.
Key features:
- Elements arranged in increasing order of atomic mass.
- Horizontal rows called periods; vertical columns called groups (I to VIII).
- Elements in the same group have similar properties.
Achievements: 1. Systematic classification of all known elements. 2. Left gaps for undiscovered elements (eka-boron → Sc, eka-aluminium → Ga, eka-silicon → Ge). 3. Predicted properties of undiscovered elements that later matched closely.
Anomalies: 1. Placed cobalt (58.9) before nickel (58.7), reversing atomic mass order to maintain similar properties. 2. Placed tellurium (127.6) before iodine (126.9) for the same reason.
1.4 Modern Periodic Law (1913, Moseley)
The physical and chemical properties of the elements are a periodic function of their atomic numbers (not atomic masses).
This resolved Mendeleev's anomalies — atomic number always increases correctly, and it is the number of protons (nuclear charge) that determines properties, not atomic mass.
2. The Modern Periodic Table
2.1 Structure
- 7 periods (horizontal rows): 1st period has 2 elements, 2nd and 3rd have 8 each, 4th and 5th have 18 each, 6th has 32, 7th is incomplete.
- 18 groups (vertical columns): Elements in the same group have similar outer electronic configurations.
- IUPAC groups 1-18 replace the old group numbering (IA, IIA, etc.).
2.2 Blocks of the Periodic Table
Elements are classified into four blocks based on the subshell being filled:
| Block | Subshell | Groups | Key Features |
|---|---|---|---|
| s-block | ns¹⁻² | 1, 2 | Highly reactive metals; low ionization enthalpy; form basic oxides |
| p-block | ns² np¹⁻⁶ | 13-18 | Metals, metalloids, and non-metals; includes noble gases |
| d-block | (n-1)d¹⁻¹⁰ ns⁰⁻² | 3-12 | Transition metals; variable oxidation states; coloured ions |
| f-block | (n-2)f¹⁻¹⁴ (n-1)d⁰⁻¹ ns² | Below table | Inner transition metals (lanthanoids and actinoids) |
2.3 Electronic Configuration Types
| Type | Pattern | Examples | Groups |
|---|---|---|---|
| Noble gas | ns² np⁶ | He, Ne, Ar, Kr, Xe, Rn | 18 |
| s-block | ns¹ or ns² | Na (3s¹), Mg (3s²) | 1, 2 |
| p-block | ns² np¹⁻⁶ | Al (3s² 3p¹), Cl (3s² 3p⁵) | 13-17 |
| d-block | (n-1)d¹⁻¹⁰ ns⁰⁻² | Fe (3d⁶ 4s²) | 3-12 |
| f-block | (n-2)f¹⁻¹⁴ | Ce (4f¹ 5d¹ 6s²) | Below table |
Group number and electronic configuration:
- For s-block: Group number = number of valence electrons
- For p-block: Group number = 10 + number of valence electrons
- For d-block: Group number = (n-1)d electrons + ns electrons
3. Periodic Trends in Properties
3.1 Atomic Radius
Atomic radius is the distance from the nucleus to the outermost shell of electrons.
Types:
- Covalent radius: Half the distance between the nuclei of two identical atoms bonded by a single covalent bond.
- Metallic radius: Half the distance between the nuclei of two adjacent metal atoms in a metallic crystal.
- Van der Waals radius: Half the distance between the nuclei of two non-bonded atoms of adjacent molecules (largest of the three).
Trend across a period (left to right): Decreases
- Nuclear charge increases; electrons are added to the same shell; effective nuclear charge increases, pulling electrons inward.
Trend down a group: Increases
- New shells are added; the distance between the nucleus and outermost electrons increases; shielding effect increases.
3.2 Ionic Radius
The radius of a charged ion.
- Cation (M⁺): Smaller than the parent atom (fewer electrons, same nuclear charge, increased effective nuclear charge per electron).
- Anion (X⁻): Larger than the parent atom (more electrons, same nuclear charge, decreased effective nuclear charge per electron).
Isoelectronic series: Species with the same number of electrons but different nuclear charges. The ion with the higher nuclear charge has the smaller radius.
Example: O²⁻ > F⁻ > Na⁺ > Mg²⁺ (all have 10 electrons, but nuclear charge increases from +8 to +12)
3.3 Ionization Enthalpy (Ionization Energy)
Ionization enthalpy is the minimum amount of energy required to remove the outermost electron from an isolated gaseous atom.
IE₂ > IE₁ because removing an electron from a positive ion requires more energy.
Trend across a period: Generally increases
- Nuclear charge increases; atomic radius decreases; electrons are held more tightly.
Trend down a group: Decreases
- Atomic radius increases; outermost electron is farther from the nucleus; shielding effect increases.
Anomalies in IE across a period:
| Anomaly | Explanation |
|---|---|
| Be > B | Be has 2s² (filled s-subshell, stable); B has 2s² 2p¹ (electron removed from less stable 2p) |
| N > O | N has 2s² 2p³ (half-filled p-subshell, extra stable); O has 2s² 2p⁴ (one paired p-electron is easier to remove) |
| Mg > Al | Mg has 3s² (filled s-subshell); Al has 3s² 3p¹ |
| P > S | P has 3s² 3p³ (half-filled); S has 3s² 3p⁴ |
3.4 Electron Gain Enthalpy (Electron Affinity)
Electron gain enthalpy is the enthalpy change when an electron is added to an isolated gaseous atom.
- For most elements, electron gain enthalpy is negative (energy is released).
- More negative value means greater tendency to gain electrons.
Trend across a period: Becomes more negative (greater affinity)
- Nuclear charge increases; atomic size decreases; added electron is more strongly attracted.
Trend down a group: Becomes less negative (lesser affinity)
- Atomic size increases; added electron is farther from the nucleus; shielding effect increases.
Important exceptions:
| Exception | Reason |
|---|---|
| F has less negative electron gain enthalpy than Cl | Small size of F causes strong inter-electronic repulsion for the incoming electron |
| O has less negative electron gain enthalpy than S | Same reason — small size of O causes repulsion |
| Noble gases have positive values | Electron would enter a new shell; stable octet already present |
| N has near-zero electron gain enthalpy | Half-filled 2p³ is stable; adding an electron disrupts it |
3.5 Electronegativity
Electronegativity is the tendency of an atom to attract the shared pair of electrons towards itself in a chemical bond.
- It is a relative property (no units). Measured on the Pauling scale.
- Fluorine has the highest electronegativity (4.0 on the Pauling scale).
- Cesium has the lowest (0.7).
Trend across a period: Increases
- Nuclear charge increases; atomic radius decreases.
Trend down a group: Decreases
- Atomic radius increases; shielding effect increases; the shared pair is farther from the nucleus.
3.6 Metallic Character
Metallic character is the tendency of an atom to lose electrons and form positive ions (cations).
Trend across a period: Decreases
- Ionization enthalpy increases; atoms hold electrons more tightly.
Trend down a group: Increases
- Ionization enthalpy decreases; outermost electrons are easier to remove.
Summary Table of Periodic Trends
| Property | Across a Period (→) | Down a Group (↓) |
|---|---|---|
| Atomic radius | Decreases | Increases |
| Ionic radius (cation) | Decreases | Increases |
| Ionization enthalpy | Increases (with anomalies) | Decreases |
| Electron gain enthalpy | More negative | Less negative |
| Electronegativity | Increases | Decreases |
| Metallic character | Decreases | Increases |
| Non-metallic character | Increases | Decreases |
| Reactivity of metals | Decreases | Increases |
| Reactivity of non-metals | Increases | Decreases |
4. Anomalous Properties of Second Period Elements
The first element of each group (Li, Be, B, C, N, O, F) shows behaviour different from the rest of the group. Reasons:
1. Small atomic size — leads to high charge density. 2. High electronegativity — greater than expected for the group. 3. High ionization enthalpy — difficult to remove electrons. 4. Absence of d-orbitals — maximum covalency is limited to 4 (cannot expand octet). 5. High polarizing power — due to small size and high charge.
Examples of anomalous behaviour:
- Lithium forms a nitride (Li₃N) directly with nitrogen, unlike other alkali metals.
- Beryllium forms covalent compounds (unlike other group 2 elements which form ionic compounds).
- Boron is a non-metal and semiconductor; other group 13 elements are metals.
- Carbon shows catenation to a remarkable degree (forms chains, rings, networks).
- Nitrogen is a diatomic gas; other group 15 elements are solids.
- Oxygen is a diatomic gas; other group 16 elements form metallic sulphides.
- Fluorine is the most electronegative element and shows only a -1 oxidation state.
5. Diagonal Relationships
Elements in the second period show similarity with elements diagonally placed in the third period:
| Period 2 | Period 3 | Similarity |
|---|---|---|
| Li | Mg | Both form nitrides, carbonates that decompose on heating, normal oxides |
| Be | Al | Both are amphoteric, have similar charge/size ratio, form covalent halides |
| B | Si | Both are metalloids, form acidic oxides, form covalent hydrides |
Cause: Diagonal neighbours have similar charge-to-radius ratios (ionic potential). Moving down a group increases size, moving right across a period increases charge — these effects cancel diagonally.
6. Important Definitions
| Term | Definition |
|---|---|
| Periodic law | Properties of elements are a periodic function of their atomic numbers |
| Group | Vertical column in the periodic table; elements have similar properties |
| Period | Horizontal row in the periodic table; properties change gradually |
| Atomic radius | Distance from the nucleus to the outermost shell of electrons |
| Covalent radius | Half the internuclear distance between two identical bonded atoms |
| Ionic radius | Radius of a charged ion |
| Isoelectronic species | Atoms/ions with the same number of electrons |
| Ionization enthalpy | Energy required to remove an electron from an isolated gaseous atom |
| Electron gain enthalpy | Enthalpy change when an electron is added to an isolated gaseous atom |
| Electronegativity | Tendency of an atom to attract shared electrons in a bond |
| Metallic character | Tendency of an atom to lose electrons and form cations |
| Diagonal relationship | Similarity between the first element of a group and the second element of the next group |
7. Common Mistakes to Avoid
1. Confusing atomic radius trends — Atomic radius decreases across a period (more protons pull electrons in) and increases down a group (new shells added). 2. Forgetting IE anomalies — Be > B and N > O are the most commonly tested exceptions. Remember: filled and half-filled subshells are extra stable. 3. Confusing electron gain enthalpy with electronegativity — Electron gain enthalpy is measured (quantitative, has units); electronegativity is relative (dimensionless, Pauling scale). 4. Thinking F has the most negative electron gain enthalpy — Cl has a more negative value than F because F's small size causes electron-electron repulsion. F has the highest electronegativity, not the most negative electron gain enthalpy. 5. Confusing iso-electronic species comparison — In an iso-electronic series, the species with the higher nuclear charge has the smaller radius: N³⁻ > O²⁻ > F⁻ > Ne > Na⁺ > Mg²⁺ > Al³⁺. 6. Writing group numbers incorrectly — Modern IUPAC uses groups 1-18, not IA, IIA, etc. Group 1 is alkali metals, Group 2 is alkaline earth metals, Group 17 is halogens, Group 18 is noble gases. 7. Ignoring d-orbital limitation in second period — Second period elements cannot expand their octet because they lack d-orbitals. This is why nitrogen forms NF₃ (not NF₅) while phosphorus forms PF₅.
8. Exam Tips
- For trend questions: Always explain trends in terms of nuclear charge and atomic size. Use phrases like "effective nuclear charge increases across a period because..." and "atomic radius increases down a group because additional shells are added."
- For anomaly questions: State the general trend, then state the anomaly, then explain the reason (filled/half-filled subshell stability, small size, electron-electron repulsion).
- For iso-electronic series: Arrange by nuclear charge — higher nuclear charge means smaller radius.
- Common 1-mark questions: Define ionization enthalpy, name the most electronegative element, arrange elements by atomic radius/IE.
- Common 3-mark questions: Explain why IE of N > O, compare F and Cl electron gain enthalpy, explain diagonal relationship Li-Mg.
- Common 5-mark questions: Explain all periodic trends with reasons, compare properties of Group 1 and Group 17 elements, discuss anomalous properties of second period elements.
9. Quick Revision Notes
- Modern periodic law: properties are a periodic function of atomic number (Z), not atomic mass
- 7 periods, 18 groups, 4 blocks (s, p, d, f)
- s-block: Groups 1-2 (alkali metals, alkaline earth metals)
- p-block: Groups 13-18 (includes halogens and noble gases)
- d-block: Groups 3-12 (transition metals)
- f-block: Lanthanoids (4f) and actinoids (5f)
- Atomic radius: decreases across a period, increases down a group
- IE: increases across (with Be>B, N>O, Mg>Al, P>S anomalies), decreases down
- Electron gain enthalpy: more negative across, less negative down; Cl > F; S > O
- Electronegativity: increases across, decreases down; F is highest (4.0)
- Metallic character: decreases across, increases down
- Iso-electronic: more nuclear charge = smaller size
- Anomalous 2nd period: small size, high EN, high IE, no d-orbitals, high polarizing power
- Diagonal: Li-Mg, Be-Al, B-Si — similar charge/size ratio
- Mendeleev: atomic mass; Modern: atomic number (Moseley)
10. Practice Questions
1. Arrange the following in order of increasing atomic radius: O, C, F, N, B. 2. Why does beryllium have a higher first ionization enthalpy than boron? 3. Which has the larger radius: (a) Na or Na⁺ (b) F or F⁻ (c) Mg²⁺ or Al³⁺? Explain. 4. Arrange the iso-electronic species in order of increasing size: N³⁻, Na⁺, F⁻, O²⁻, Mg²⁺. 5. Explain why electron gain enthalpy of fluorine is less negative than that of chlorine. 6. What is the diagonal relationship between lithium and magnesium? Give two examples of similar properties. 7. Predict the group and period of an element with electronic configuration [Ar] 3d¹⁰ 4s² 4p³. 8. Why is the first ionization enthalpy of nitrogen greater than that of oxygen? 9. Arrange the following in order of decreasing electronegativity: N, O, F, C, B. 10. State two properties in which lithium resembles magnesium. Why does this resemblance occur?
Quick Answers to Practice Questions
1. F < O < N < C < B (atomic radius increases as we move left in a period) 2. Be has 2s² (filled subshell, more stable); B has 2s² 2p¹ (electron removed from less stable 2p orbital) 3. (a) Na > Na⁺ (cation smaller than atom) (b) F⁻ > F (anion larger than atom) (c) Mg²⁺ > Al³⁺ (both have 10 electrons; Al has higher nuclear charge, so smaller) 4. Mg²⁺ < Al³⁺ < Na⁺ < F⁻ < O²⁻ < N³⁻ (higher nuclear charge = smaller radius in iso-electronic series; wait — let me correct: Mg²⁺ has 12 protons, 10 electrons; Al³⁺ has 13 protons, 10 electrons — Al³⁺ is smaller. So: Al³⁺ < Mg²⁺ < Na⁺ < F⁻ < O²⁻ < N³⁻) 5. F is small; the incoming electron experiences strong repulsion from the dense electron cloud. Cl is larger, with more space for the incoming electron, so electron gain enthalpy is more negative. 6. Li resembles Mg: both form normal oxides (Li₂O, MgO), both form nitrides (Li₃N, Mg₃N₂), both form carbonates that decompose on heating (Li₂CO₃, MgCO₃). Resemblance is due to similar charge-to-size ratio. 7. Group 15 (10 + 5 valence electrons = 15), Period 4 (4th shell is being filled). 8. N has 2s² 2p³ (half-filled p-subshell, extra stable). O has 2s² 2p⁴ (one electron is paired, easier to remove). Half-filled stability makes N harder to ionize. 9. F > O > N > C > B (electronegativity decreases going left across period 2) 10. Li and Mg: both form normal oxides, both form carbonates that decompose on heating, both form nitrides with nitrogen. Resemblance occurs because of similar ionic potential (charge-to-size ratio).
11. Final Revision Checklist
Before the exam, make sure you can:
- Trace the historical development from Dobereiner to Moseley
- Explain the modern periodic law and how it resolved Mendeleev's anomalies
- Classify elements into s, p, d, f blocks and explain their properties
- State and explain all periodic trends (atomic radius, IE, EGE, EN, metallic character) with reasons
- Explain exceptions: Be > B, N > O, F vs Cl electron gain enthalpy
- Define and explain anomalous properties of second period elements
- Explain diagonal relationships with examples (Li-Mg, Be-Al)
- Compare iso-electronic species and arrange them by size
- Predict group, period, and block from electronic configuration