NCERT Solutions for Class 11 Chemistry Chapter 3: Classification of Elements and Periodicity in Properties
If there is one chapter in Class 11 Chemistry that gives you a bird's-eye view of all elements in existence, it is Chapter 3 — Classification of Elements and Periodicity in Properties. The periodic table is arguably the most powerful tool in chemistry, and this chapter explains how it was developed, how it is organised, and what patterns it reveals. At Myclass24, we have developed NCERT solutions for this chapter that make periodic trends intuitive, not just something you memorise before the exam. Students often struggle with understanding why ionisation energy decreases down a group but increases across a period — or why atomic radius behaves the way it does. Our solutions walk through each concept with a clear explanation backed by the reasoning behind it. One can check out all chapters of NCERT Solutions for Class 11 Chemistry and all subjects of NCERT Solutions for Class 11 from the Myclass24 page.
The NCERT solutions for chapter also introduces the concept of periodicity — the idea that properties of elements repeat at regular intervals when arranged by atomic number. This is not just useful for board exams; understanding periodicity deeply will help students predict properties of elements and compounds in higher classes and competitive exams like JEE Advanced and NEET. The solutions at Myclass24 are crafted for all learners — whether you are a student in Bangalore, Lucknow, Chandigarh, or Bhopal, you will find them easy to follow and effective for revision.
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NCERT Solutions — Class 11 Chemistry Chapter 3: Classification of Elements and Periodicity in Properties
Trend-based explanations, periodic table data, and exercise solutions | Free PDF on Myclass24
Chapter 3 Explained: Periodic Table, Trends & Key Facts
History of Classification of Elements
The attempt to organise elements is centuries old. Early classification by Döbereiner grouped elements into triads where the middle element's atomic mass was the average of the other two. Newlands then arranged elements by atomic mass and noted that every eighth element had similar properties — his Law of Octaves. However, it was Dmitri Mendeleev (1869) who created the first successful periodic table, arranging elements by atomic mass and leaving gaps for undiscovered elements. He even predicted the properties of three unknown elements: Eka-boron (Sc), Eka-aluminium (Ga), and Eka-silicon (Ge). Henry Moseley later showed that atomic number, not mass, is the true basis of periodicity.
| Scientist | Contribution | Basis of Arrangement |
|---|---|---|
| Döbereiner (1817) | Law of Triads | Atomic mass (groups of 3) |
| Newlands (1864) | Law of Octaves | Atomic mass |
| Mendeleev (1869) | First Periodic Table | Atomic mass |
| Moseley (1913) | Modern Periodic Law | Atomic number |
Modern Periodic Table — Structure
The modern periodic table has 118 elements arranged in 7 periods (horizontal rows) and 18 groups (vertical columns). The period number corresponds to the highest principal quantum number of the element's valence electrons. Groups 1–2 are s-block elements, groups 13–18 are p-block, groups 3–12 are d-block (transition metals), and the lanthanides and actinides are f-block elements.
| Block | Groups | Valence Subshell | Example |
|---|---|---|---|
| s-block | 1, 2 | ns¹⁻² | Na, Mg |
| p-block | 13–18 | np¹⁻⁶ | C, N, Cl, Ne |
| d-block | 3–12 | (n−1)d¹⁻¹⁰ | Fe, Cu, Zn |
| f-block | Inner transition | (n−2)f¹⁻¹⁴ | La, U |
Atomic Radius
Atomic radius generally decreases across a period (left to right) because the nuclear charge increases while electrons are added to the same shell — pulling them closer. Down a group, atomic radius increases because a new electron shell is added with each period, increasing the distance from the nucleus. For metals, atomic radius is typically the metallic radius; for non-metals, the covalent radius is used.
Ionisation Enthalpy (Ionisation Energy)
Ionisation enthalpy is the energy required to remove the outermost electron from an isolated gaseous atom. It increases across a period (stronger nuclear pull) and decreases down a group (electron farther from nucleus, more shielded). Successive ionisation enthalpies always increase — it takes more energy to remove each subsequent electron. Noble gases have the highest ionisation energies in their periods.
| Periodic Trend | Across Period (→) | Down Group (↓) |
|---|---|---|
| Atomic Radius | Decreases | Increases |
| Ionisation Enthalpy | Increases | Decreases |
| Electron Gain Enthalpy | More negative (increases) | Less negative (decreases) |
| Electronegativity | Increases | Decreases |
| Metallic Character | Decreases | Increases |
| Non-Metallic Character | Increases | Decreases |
Electronegativity and Electron Gain Enthalpy
Electronegativity is the tendency of an atom to attract shared electrons in a bond. Fluorine is the most electronegative element (Pauling scale: 4.0). Electron gain enthalpy is the energy change when an electron is added to a neutral gaseous atom. Chlorine, not fluorine, has the most negative electron gain enthalpy because F's small size creates electron-electron repulsion that slightly resists gaining an electron.
Valency and Oxidation States
Valency is the combining capacity of an element. For s- and p-block elements, valency is straightforward — it is often equal to the group number or (18 − group number). Transition metals (d-block) show variable oxidation states because electrons from both the d and s subshells can participate in bonding. For example, iron can be Fe²⁺ or Fe³⁺, and manganese can range from +2 to +7.
Periodicity in Properties — Additional Facts:
- Elements in the same group have similar chemical properties because they have the same number of valence electrons.
- Diagonal relationships exist: Li resembles Mg, Be resembles Al, B resembles Si in certain properties.
- The d-block elements of Period 4 (Sc to Zn) are often called the first transition series.
- Oxides of metals are generally basic; oxides of non-metals are acidic; oxides of metalloids are amphoteric.
- The most electronegative element is F (4.0); the least electronegative is Cs (0.79).
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FAQs for NCERT Solutions Class 11 Chemistry Chapter 3 Periodic Properties
The modern periodic table organizes elements based on increasing atomic number. It groups elements with similar chemical properties into the same columns, making it easier to predict their behavior. The periodic table is divided into periods and groups and helps students understand trends such as atomic size, ionization energy, and electronegativity. It serves as a roadmap for studying chemistry because it reveals relationships among elements. Understanding the modern periodic table allows students to analyze chemical properties efficiently and predict reactions. This topic forms the foundation for advanced studies in inorganic chemistry and chemical bonding.
Periodicity refers to the recurring pattern of physical and chemical properties of elements when arranged by atomic number. As elements are placed in the periodic table, properties such as atomic radius, ionization enthalpy, electron affinity, and electronegativity show regular trends. These trends arise from the electronic configuration of elements. Understanding periodicity helps students compare elements and predict their behavior. It also explains why elements in the same group have similar chemical properties. Questions based on periodic trends are common in examinations because they require both conceptual understanding and analytical thinking.
Atomic size changes due to variations in nuclear charge and the number of electron shells. Across a period, atomic size generally decreases because the increasing nuclear charge pulls electrons closer to the nucleus. Down a group, atomic size increases as new electron shells are added. Understanding these trends helps students explain many other periodic properties, including ionization energy and electronegativity. Atomic radius is a frequently tested topic in board and competitive examinations. Learning the reasons behind these trends improves conceptual clarity and enables students to solve comparative questions accurately.
Ionization enthalpy is the energy required to remove the most loosely bound electron from an isolated gaseous atom. It depends on factors such as atomic size, nuclear charge, shielding effect, and electron configuration. Smaller atoms generally have higher ionization enthalpy because their electrons are held more strongly by the nucleus. Trends in ionization enthalpy help explain the reactivity of elements and their positions in the periodic table. Understanding this concept is essential for studying chemical bonding and periodicity. Examination questions often focus on comparing ionization energies among different elements and explaining observed trends.