NCERT Solutions
Class 11 Chemistry
9 chapters · 47 important questions · Updated 2025-26
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Some Basic Concepts of Chemistry
This chapter lays the quantitative foundation of chemistry by introducing the mole concept, atomic and molecular masses, and stoichiometry. Students learn to calculate empirical and molecular formulae from percentage composition data. Laws of chemical combination — conservation of mass, definite proportions, multiple proportions, and Gay-Lussac's law — are thoroughly covered. Molar volume, Avogadro's number, and limiting reagent calculations form the core of this chapter.
Key Topics
Important Questions
Calculate the empirical and molecular formula of a compound containing 40% carbon, 6.67% hydrogen, and 53.33% oxygen. Its molar mass is 60 g/mol.
State the Law of Conservation of Mass. In a reaction, 1.6 g of CH₄ burns completely in oxygen. Calculate the mass of CO₂ and H₂O formed.
Calculate the number of molecules and atoms in 5.6 L of CO₂ at STP.
0.5 mol of H₂ reacts with 0.5 mol of O₂ to form water. Identify the limiting reagent and calculate the mass of water formed.
What is the molarity of a solution prepared by dissolving 5.85 g of NaCl (Molar mass = 58.5 g/mol) in 500 mL of water?
Structure of Atom
This chapter traces the evolution of atomic models from Thomson's plum-pudding model through Rutherford's nuclear model to Bohr's model and the quantum mechanical model. Quantum numbers, the aufbau principle, Pauli exclusion principle, and Hund's rule govern electronic configurations. The dual nature of matter (de Broglie relation) and Heisenberg's uncertainty principle are key modern concepts. Students must also understand hydrogen spectrum line series (Lyman, Balmer, Paschen).
Key Topics
Important Questions
State Heisenberg's Uncertainty Principle. If the velocity of an electron is 2.2 × 10⁶ m/s with uncertainty of 0.001%, calculate the uncertainty in its position.
Write the electronic configuration of Cr (Z=24) and Cu (Z=29) and explain why they are exceptions to the Aufbau principle.
Calculate the wavelength of the electron moving with a velocity of 2.05 × 10⁷ m/s. (h = 6.626 × 10⁻³⁴ J·s; mₑ = 9.1 × 10⁻³¹ kg)
Explain Bohr's model of the hydrogen atom. What are its two main limitations?
How many electrons in an atom can have n = 3? Write all the subshells and the number of orbitals in each.
Classification of Elements and Periodicity in Properties
This chapter covers the historical development of periodic classification from Döbereiner's triads to Mendeleev's table to the modern long form of the periodic table. Students study periodic trends in atomic radius, ionic radius, ionisation enthalpy, electron gain enthalpy, and electronegativity. The concept of effective nuclear charge (Zeff) and shielding is central to explaining all periodic trends. Anomalous properties of second-period elements and diagonal relationships are important for board exams.
Key Topics
Important Questions
The first ionisation enthalpies of B, C, N, O are 800, 1086, 1402, 1314 kJ/mol respectively. Why is the IE₁ of O less than that of N?
Arrange F, Cl, Br, I in increasing order of (i) atomic radius and (ii) electron gain enthalpy, with justification.
What is the diagonal relationship? Explain with the example of Li and Mg.
Define ionisation enthalpy. How does it vary along a period and down a group? Give reasons.
Write the number of elements in (i) 1st period, (ii) 2nd period, (iii) 4th period of the modern periodic table.
Chemical Bonding and Molecular Structure
This chapter explains why and how atoms combine to form molecules using the octet rule, ionic bonding, covalent bonding, and VSEPR theory. Valence Bond Theory (VBT) introduces orbital overlap and hybridisation (sp, sp², sp³, sp³d, sp³d²), while Molecular Orbital Theory (MOT) describes bonding in terms of BMOs and ABMOs. Resonance, formal charge, dipole moment, and the properties of ionic and covalent compounds are important exam topics.
Key Topics
Important Questions
Predict the shapes of PCl₅ and SF₆ using VSEPR theory. State the hybridisation of the central atom in each.
Draw Lewis structures of CO₂, SO₃, and NO₃⁻. Indicate formal charges wherever applicable.
Explain why NH₃ has a higher boiling point than PH₃.
On the basis of Molecular Orbital Theory, predict the bond order and magnetic nature of O₂ and O₂²⁻.
Why is the H–O–H bond angle in water 104.5° and not 109.5°?
What is resonance? Explain with the structure of SO₂.
Thermodynamics
This chapter applies the laws of thermodynamics to chemical reactions, defining internal energy, enthalpy, and entropy as state functions. The first law (ΔU = q + w) and Hess's law of constant heat summation are foundational. Standard enthalpies of formation, combustion, atomisation, bond dissociation, and solution/hydration are covered in depth. The second law introduces entropy and spontaneity; the Gibbs free energy equation (ΔG = ΔH − TΔS) determines whether a reaction is spontaneous.
Key Topics
Important Questions
Using Hess's law, calculate the standard enthalpy of formation of CH₄(g) given that: C(s) + O₂(g) → CO₂(g), ΔH° = −393.5 kJ/mol; H₂(g) + ½O₂(g) → H₂O(l), ΔH° = −285.8 kJ/mol; CH₄(g) + 2O₂(g) → CO₂(g) + 2H₂O(l), ΔH° = −890.4 kJ/mol.
For the reaction N₂(g) + 3H₂(g) → 2NH₃(g), ΔH° = −92 kJ/mol. Calculate ΔU° at 27°C. (R = 8.314 J mol⁻¹ K⁻¹)
Predict the sign of ΔS for the following reactions: (i) CaCO₃(s) → CaO(s) + CO₂(g); (ii) 2SO₂(g) + O₂(g) → 2SO₃(g). Give reasons.
For a reaction ΔH = +ve and ΔS = +ve. At what temperature will the reaction be spontaneous? Justify using ΔG = ΔH − TΔS.
What is the relationship between standard Gibbs free energy change (ΔG°) and the equilibrium constant K? If ΔG° < 0, what does this imply about K?
Equilibrium
This chapter examines both chemical and ionic equilibrium. For chemical equilibrium, the law of mass action gives the equilibrium constant expression (Kc and Kp), and Le Chatelier's principle predicts how equilibrium shifts on changing concentration, pressure, or temperature. Ionic equilibrium covers Arrhenius, Brønsted–Lowry, and Lewis definitions of acids and bases, pH, the ionic product of water (Kw), Ka, Kb, and their relationship pKa + pKb = pKw, as well as buffer solutions and the Henderson–Hasselbalch equation.
Key Topics
Important Questions
For the equilibrium H₂(g) + I₂(g) ⇌ 2HI(g), Kc = 57.0 at 700 K. If 0.5 mol H₂ and 0.5 mol I₂ are placed in a 1 L container, find the equilibrium concentrations.
State Le Chatelier's Principle. How will increasing pressure affect the equilibrium N₂(g) + 3H₂(g) ⇌ 2NH₃(g)?
Calculate the pH of a 0.01 M HCl solution and a 0.001 M NaOH solution.
What is a buffer solution? Calculate the pH of a buffer containing 0.1 M acetic acid and 0.1 M sodium acetate. (pKa of acetic acid = 4.74)
Derive the relationship Kp = Kc(RT)^Δng. For which reactions is Kp = Kc?
Redox Reactions
This chapter develops the concept of oxidation and reduction in terms of electron transfer and changes in oxidation number. Students learn to assign oxidation states systematically and to balance redox equations by both the oxidation number method and the ion-electron (half-reaction) method. The electrochemical series and the concept of reducing and oxidising agents, disproportionation reactions, and comproportionation are also covered.
Key Topics
Important Questions
Balance the following redox equation by the ion-electron method in acidic medium: MnO₄⁻ + Fe²⁺ → Mn²⁺ + Fe³⁺
Assign oxidation numbers to each element in the following: (i) K₂Cr₂O₇, (ii) Na₂S₂O₃, (iii) H₂SO₅.
Identify the oxidising agent and reducing agent in the reaction: Zn + CuSO₄ → ZnSO₄ + Cu. Justify.
What is a disproportionation reaction? Give one example and identify the element that is simultaneously oxidised and reduced.
Balance by oxidation number method: HNO₃ (dil.) + Cu → Cu(NO₃)₂ + NO + H₂O.
Organic Chemistry – Some Basic Principles and Techniques
This chapter introduces the principles of organic chemistry: classification of organic compounds, IUPAC nomenclature, types of organic reactions, reaction intermediates (carbocations, carbanions, free radicals, carbenes), electronic effects (inductive, mesomeric/resonance, hyperconjugation), and methods of purification. Qualitative analysis — detection of C, H, N, S, and halogens (Lassaigne's test) — and quantitative analysis (Dumas, Kjeldahl methods) are also covered.
Key Topics
Important Questions
Name the following compounds using IUPAC rules: (i) CH₃CH(OH)CH₂CHO; (ii) (CH₃)₃CCl; (iii) CH₂=CHCH₂CH₃.
Explain the inductive effect. Using the inductive effect, compare the acidity of formic acid and acetic acid.
What is hyperconjugation? How does it explain the stability of alkenes with greater substitution?
Describe Lassaigne's test for the detection of nitrogen in an organic compound. Write the relevant chemical equations.
Arrange the following carbocations in decreasing order of stability: (CH₃)₃C⁺, CH₃CH⁺CH₃, CH₃CH₂CH₂⁺.
Hydrocarbons
This chapter systematically covers the chemistry of alkanes, alkenes, alkynes, and aromatic hydrocarbons. Alkane reactions include free-radical halogenation (mechanism in steps); alkene reactions include electrophilic addition (Markovnikov's rule, anti-Markovnikov addition via HBr/peroxide), ozonolysis, and oxidation. Alkynes form acidic H due to sp hybridisation. Benzene's aromaticity (Hückel's 4n+2 rule), electrophilic aromatic substitution (EAS — halogenation, nitration, sulphonation, Friedel-Crafts), and conformations (Sawhorse and Newman) of alkanes are key topics.
Key Topics
Important Questions
Explain the mechanism of electrophilic addition of HBr to propene. What product is predominantly formed and why (Markovnikov's rule)?
Give the mechanism of free-radical chlorination of methane. Name the three steps.
What is Markovnikov's rule? How is it violated in the presence of peroxides? Give the product of HBr addition to propene with and without peroxide.
Explain the mechanism of nitration of benzene. What is the electrophile involved?
Draw the Newman projection of the staggered and eclipsed conformations of ethane and compare their stability.
Why are terminal alkynes more acidic than alkenes? Explain in terms of hybridisation.
Frequently Asked Questions
Is NCERT enough for CBSE Class 11 Chemistry board exam?
Yes. CBSE board exams are designed entirely around NCERT. 80–90% of questions in the Class 11 Chemistry paper are directly based on NCERT concepts — sometimes reworded, never from outside NCERT. Completing all 47 important questions listed here is the minimum you need.
How to study NCERT Class 11 Chemistry for board exams?
Go chapter-by-chapter. Read the chapter first, then attempt the important questions without looking at answers. Check your answer structure — CBSE gives marks per point, so structure matters as much as content. For 5-mark answers: brief intro + 4–5 numbered points with keywords + conclusion.
Which chapters are most important in Class 11 Chemistry for CBSE boards?
Chapters with the most long-answer (5-mark) important questions carry the most marks in the board paper. Look at which chapters here have the highest question count — those are the ones CBSE has historically focused on. Don't skip any chapter, but spend extra time on these.
How many questions come from NCERT in CBSE Class 11 Chemistry exam?
Nearly all 47 questions in the Class 11 Chemistry board paper are rooted in NCERT. The exact question may be reworded, but the concept, definition, or formula always comes from the NCERT textbook. Practicing these 47 important questions covers the vast majority of what can be asked.