Class 11 ChemistryImportant Questions for Board Exams 2025-26

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?

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.

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.

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.

A sample of gas occupies 2.0 L at 27°C and 1 atm. Calculate its volume at 127°C and 2 atm.

State the postulates of Kinetic Molecular Theory of gases. How does it explain Boyle's law?

Write the van der Waals equation for n moles of a real gas. What do the constants 'a' and 'b' signify?

Calculate the root mean square speed of O₂ molecules at 27°C. (R = 8.314 J mol⁻¹ K⁻¹, M = 32 × 10⁻³ kg/mol)

What is critical temperature? Why can CO₂ be liquefied at room temperature but not N₂?

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?

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?

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.

Why is hydrogen considered anomalous in the periodic table? Justify its placement above both Group 1 and Group 17.

Describe the preparation of hydrogen peroxide from barium peroxide. Why is it stored in dark-coloured bottles?

Distinguish between temporary and permanent hardness of water. How is permanent hardness removed by the ion-exchange method?

Draw the structure of H₂O₂. Is it planar or non-planar? Give one oxidising and one reducing property of H₂O₂.

Classify hydrides into ionic, covalent, and metallic with one example each.

Explain the Solvay process for the manufacture of sodium carbonate (Na₂CO₃). Write the balanced equations for all key steps.

Why does lithium show anomalous behaviour compared to other alkali metals? Give four points of similarity between Li and Mg.

What happens when Na₂O₂ reacts with water? Write the equation. Why is Na₂O₂ used in breathing masks?

Write the reactions for: (i) action of excess CO₂ on Ca(OH)₂; (ii) heating of CaCO₃.

Compare the solubility of Group 2 hydroxides and sulphates down the group and explain the trend.

Draw the structure of diborane (B₂H₆). Explain the 3-centre 2-electron bonds.

Compare the properties of diamond and graphite in terms of structure, hybridisation, electrical conductivity, and hardness.

What is the borax bead test? Explain with the example of copper sulphate.

What are silicones? Write their general formula and list two important uses.

Why does the stability of +2 oxidation state increase down Group 14? Illustrate the inert pair effect.

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⁺.

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.

What is acid rain? How are SO₂ and NOₓ responsible for acid rain? What are its harmful effects?

Explain the mechanism of ozone layer depletion by chlorofluorocarbons (CFCs). Write the relevant equations.

What is photochemical smog? How is PAN (peroxyacetyl nitrate) formed? What are its effects?

Define BOD. Why is the BOD level of a water sample a measure of its pollution?

What is eutrophication? How does it lead to the death of aquatic life? Give two ways to control it.