CBSE Class 11 Chemistry Important Questions with Solutions for 2025-26 exams

Chemistry in Class 11 can feel like a big jump - suddenly there are new terms, strange-looking orbitals, long organic reactions, and concepts that take a minute to settle. Feeling confused at first is completely normal.

But once you start solving questions instead of only reading, everything becomes clearer. Important questions show you what CBSE actually focuses on and help you understand how each concept is used in real problems.

This blog gives you a clean, student-friendly set of Class 11 Chemistry questions with simple explanations - like someone guiding you step by step.

What’s inside

• 20 most important Class 11 Chemistry questions with solutions
• 10 extra practice questions for self-check
• Why practising these questions actually helps
• Benefits that genuinely improve your preparation
• FAQs based on real student doubts

Chapter-wise most important questions for class 11 Chemistry

In the table given below, we have provided the links to downloadable PDFs of chapter-wise most important questions for class 11 Chemistry and that too for different categories of marks.

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Some Important Questions with Answers

Q1. What is the difference between molecular mass and molar mass?

Answer: Think of it this way - molecular mass tells you the mass of one single molecule, and it’s measured in atomic mass units (u). Molar mass tells you the mass of 1 mole of those molecules (6.022 × 10²³ of them), and it’s measured in g/mol.

Both values are usually the same number, but the meaning changes.

Example: For H₂O → 2(1) + 16 = 18 u, So its molar mass is 18 g/mol (because 1 mole weighs 18 g).

Why it matters:

• Molecular mass → microscopic calculations
• Molar mass → stoichiometry, conversions, reacting mass problems

Q2. Derive the ideal gas equation using Boyle’s, Charles’ and Avogadro’s laws.

Answer: All three gas laws talk about how gases behave:

• Boyle’s Law: V ∝ 1/P
• Charles’ Law: V ∝ T
• Avogadro’s Law: V ∝ n

Combine them → V ∝ nT / P

Bring in the constant R (gas constant) to convert it into an equation: PV = nRT

Why useful: This equation helps you calculate any gas-related value (P, V, T, n) as long as other values are known.

Q3. Explain Heisenberg’s Uncertainty Principle with an example.

Answer: Heisenberg basically said: “You cannot know both the exact position and exact momentum of a particle at the same time.”

Formula: Δx × Δp ≥ h/4π`

Example: To “see” an electron, you use high-energy radiation.
But that radiation hits the electron and changes its momentum.
So, the more precisely you try to locate it, the less you know about its momentum.

Why important: This proved electrons cannot follow fixed orbits (Bohr failed), and instead stay in probability clouds (orbitals).

Q4. State Hund’s Rule and apply it to nitrogen (Z = 7).

Answer: Hund’s Rule: Electrons fill degenerate orbitals singly first, with parallel spins, to reduce repulsion.

Nitrogen (7 electrons):
1s² 2s² 2p³ → the 3 electrons go into three separate p-orbitals: ↑ ↑ ↑

This arrangement increases stability.

Q5. What is hybridisation? Explain sp³ hybridisation.

Answer: Hybridisation = mixing of atomic orbitals to form new, identical hybrid orbitals.

sp³ hybridisation: 1 s-orbital + 3 p-orbitals → 4 sp³ hybrids

• Shape: tetrahedral
• Bond angle: 109.5°

Example: In CH₄, carbon uses sp³ orbitals to make four equal C–H sigma bonds.

Q6. Explain hydrogen bonding in water and its effect on boiling point.

Answer: In water, oxygen is very electronegative, so the O-H bonds become polar. Hydrogen from one water molecule sticks to oxygen of another → hydrogen bond.

These bonds are stronger than normal dipole attractions.

To boil water, you must break these bonds - that needs high energy. That’s why water boils at 100°C, which is unusually high for a small molecule.

Also explains:

• high surface tension
• high specific heat

Q7. Define enthalpy of formation.

Answer: Enthalpy of formation (ΔHf°): Heat changes when 1 mole of a compound forms from its elements in their standard states.

Example: C (graphite) + O₂ → CO₂ ; ΔHf° = –393.5 kJ/mol

Used in:

• Hess’s Law
• Enthalpy calculations
• Born–Haber cycles

Q8. State and explain Hess’s Law.

Answer: Hess’s Law: The total enthalpy change of a reaction is the same, no matter what path you take. Because enthalpy is a state function.

Use: Helps calculate ΔH when direct measurement isn’t possible by adding/subtracting simpler equations.

Q9. What are sigma and pi bonds?

Answer: Sigma (σ) bond:

• Formed by head-on overlap
• Stronger
• Allows rotation

Pi (π) bond:

• Sideways overlap
• Weaker
• Restricts rotation

Double bond = 1 σ + 1 π
Triple = 1 σ + 2 π

Q10. Difference between electrolytes & non-electrolytes.

Answer: Electrolytes:

• Ionise in water
• Conduct electricity
• Eg: NaCl, HCl, CH₃COOH (weak)

Non-electrolytes:

• Don’t ionise
• Don’t conduct
• Eg: glucose, urea

Conductivity depends on the number of ions & degree of ionisation.

Q11. State Raoult’s Law.

Answer: For an ideal solution:

Pa = Xa × P°a (total vapour pressure - sum of all partial pressures)

Raoult’s Law forms the basis of:

• boiling point elevation
• freezing point depression
• osmotic pressure

Q12. What is a redox reaction? Give an example.

Answer: Redox = oxidation + reduction

• Oxidation → loss of electrons
• Reduction → gain of electrons

Example:Zn + Cu²⁺ → Zn²⁺ + Cu
Zn is oxidised; Cu²⁺ is reduced.

Redox reactions occur in batteries, rusting, respiration, etc.

13. Explain electron gain enthalpy across a period.

Answer: Across a period:

• Nuclear charge ↑
• Atomic radius ↓
• Attraction for incoming electron ↑
  → Electron gain enthalpy becomes more negative.

Exceptions:

• Noble gases
• N and Mg (extra stability)

Q14. What is Pauli’s Exclusion Principle?

Answer: No two electrons in an atom can have the same four quantum numbers. So, one orbital can hold a maximum of two electrons, with opposite spins. Basis for the periodic table and electronic configurations.

Q15. Define limiting reagent with an example.

Answer: Limiting reagent = the reactant that gets used up first, deciding how much product forms.

Example: 2H₂ + O₂ → 2H₂O
If H₂ is less than required, it becomes the limiting reagent. Important in stoichiometry and industry calculations.

Q16. What is lattice enthalpy?

Answer: Lattice enthalpy = energy released when gaseous ions form 1 mole of an ionic solid.

Higher lattice enthalpy → stronger ionic bond → higher melting point.

Example: MgO > NaCl (because Mg²⁺ & O²⁻ have higher charges)

Q17. Explain the law of chemical equilibrium.

Answer: At equilibrium: Rate of forward reaction = rate of backward reaction.

Equilibrium constant: Kc = [products]ᶜ / [reactants]ᵃ

• Large Kc → products favoured
• Small Kc → reactants favoured

Foundation of Le Chatelier’s Principle.

Q18. What are isomers?

Answer: here is the explanation in points:

• Same molecular formula → different arrangements.
• Structural isomers: different connectivity.
• Stereoisomers: same connectivity, different spatial arrangement.
• Example: Butane (straight) vs isobutane (branched).

Q19. Explain the concept of solubility product (Ksp).

Answer: Here is the explanation:

• Ksp = equilibrium constant for dissolution of a sparingly soluble salt.
• For AB → A⁺ + B⁻: Ksp = [A⁺][B⁻]
• Higher Ksp → more soluble.
• Used to predict precipitation.
• Common ion effect reduces solubility.

Q20. What is the difference between adsorption and absorption?

Answer: Here is the differences between adsorptions and absorptions:

1. Adsorption: accumulation on surface.

2. Absorption: uniform entry into bulk.

Overall, Adsorption is a surface phenomenon; absorption is a bulk phenomenon.

Q21. What is the difference between order of a reaction and molecularity?

Answer: Order = how the rate depends on concentration. It is found experimentally.
Molecularity = number of molecules colliding in the slowest step. It is theoretical.
Key differences:

• Order can be 0, fractional or even negative.
• Molecularity is always a whole number (1, 2, 3...).
• Order applies to complex reactions; molecularity only to elementary reactions.

Q22. Explain Le Chatelier’s Principle with an example.

Answer: Le Chatelier’s Principle says: “If a system at equilibrium is disturbed, it will shift in a direction that reduces the disturbance.”
Example: For N₂ + 3H₂ ⇌ 2NH₃ (exothermic),

• Increase temperature → equilibrium shifts backward.
• Increase pressure → equilibrium shifts forward (less moles).

Useful in predicting direction of equilibrium shift in chemical reactions.

Q23. Define pH and derive its formula.

Answer: pH tells you how acidic or basic a solution is.
Definition: pH = –log[H⁺]
Meaning:
Lower pH → more acidic
Higher pH → more basic
Example: [H⁺] = 1 × 10⁻³ → pH = 3
pH scale ranges from 0 to 14 for most solutions.

Q24. What are colligative properties? Give examples.

Answer: Colligative properties depend only on the number of solute particles, not their type.
Examples:

• Relative lowering of vapour pressure
• Boiling point elevation
• Freezing point depression
• Osmotic pressure, Used to find molar mass and behaviour of solutions.

Q25. What is the common ion effect?

Answer: The common ion effect is the decrease in solubility of an ionic compound when a solution already contains one of its ions.
Example: Solubility of AgCl decreases in solution containing NaCl (common ion = Cl⁻).
Used in:

• Salt analysis
• Buffer preparation
• Precipitation control

Q26. What is the difference between oxidation number and valency?

Answer: Oxidation number:

• Hypothetical charge an atom would have in a compound.
• Can be positive, negative, fractional, or zero.

Valency:

• Combining capacity of an element.
• Always a whole number.
  Example: In H₂SO₄, oxidation number of S = +6, but its valency = 6.

Q27. Explain the term “activation energy.”

Answer: Activation energy (Ea) is the minimum energy required for a reaction to start.
If molecules collide with energy ≥ Ea → reaction occurs.
If not → no reaction.
Catalysts lower activation energy, making reactions faster without being consumed.

Q28. What is a buffer solution? Give an example.

Answer: A buffer solution resists changes in pH when small amounts of acid or base are added.
Made of:

• Weak acid + its salt (CH₃COOH + CH₃COONa) or
• Weak base + its salt (NH₄OH + NH₄Cl) Buffers are important in:
• Blood pH (7.4)
• Biological systems
• Analytical chemistry

Q29. What is the Arrhenius equation? What does it explain?

Answer: Arrhenius Equation:
k = A e^(–Ea/RT)
Where:
k = rate constant
A = frequency factor
Ea = activation energy
T = temperature

Meaning:
Higher temperature → higher k → faster reaction
Lower Ea → faster reaction
Explains how reactions speed up when heated.

Q30. Explain the concept of resonance.

Answer: Resonance occurs when one structure is not enough to represent a molecule. Actual structure is a hybrid of two or more resonance structures.
Example: O₃, CO₃²⁻, benzene
Resonance increases:

• Stability
• Equal bond lengths
• Charge distribution

10 Extra Questions for Practice!

Q1. State the Aufbau principle with an example.

Q2. Write the postulates of Dalton’s atomic theory.

Q3. What is the difference between crystalline and amorphous solids?

Q4. Explain the concept of ionisation enthalpy.

Q5. Define azeotropes.

Q6. What is the significance of the Schrodinger wave equation?

Q7. Explain the term polymerisation.

Q8. What are amphoteric substances?

Q9. Give the IUPAC name of (CH₃)₂CH–CH₂–OH.

Q10. Why are noble gases chemically inert?

Importance of Practising Important Questions

Practising important questions is honestly one of the smartest things you can do while preparing. Think of them as a shortcut map - they show you exactly what CBSE repeats, what examiners love, and which parts of the chapter actually matter.

When you solve these questions regularly, you slowly start seeing a pattern: which definitions are always asked, which diagrams fetch marks, and which steps you can’t afford to miss.

It also trains the two skills every board exam checks the most:

• Do you actually understand the concept?
• Can you apply it properly in a question?

A lot of students find Chemistry tough simply because they keep reading and rereading the textbook. But the moment you start solving questions, everything becomes more active - you think, you recall, you connect formulas, and you quickly notice where you’re getting stuck.

Even big theory chapters feel lighter because practising questions breaks them into scoring, manageable chunks. Basically, your prep becomes sharper, faster, and much more exam-oriented.

Benefits of Solving These Questions 

This section is made to actually help you study smarter - not harder. Think of it like your quick warm-up before the real prep begins.

• You start feeling way more confident because the important topics don’t look scary anymore.
• You understand how CBSE actually twists concepts in questions instead of just memorising theory.
• Your speed gets better, especially in those Physical Chemistry numericals that usually take forever.
• The exam pattern starts feeling familiar, so the fear slowly disappears.
• Revision becomes super easy because you already know what matters and what doesn’t.

FAQs

Q1. Are these Class 11 Chemistry questions enough for exams?

Ans. Not completely - but they cover all the important concepts you MUST know. If you want full confidence, solve NCERT back questions too. Think of these as your core practice, and NCERT as your final polish.

Q2. What’s the best way to revise Chemistry without forgetting everything?

Ans. Break it into 3 steps:

• Read NCERT once
• Practise numericals + important questions
• Revise every week using short notes 

If you follow this, older chapters won’t fade, and formulas stay fresh.

Q3. Is Class 11 Chemistry actually harder than Class 12?

Ans.  Most students say yes  -not because the syllabus is huge, but because everything is NEW. Once basics like atomic structure, thermodynamics, and equilibrium are clear, Class 12 feels way smoother.

Q4. Do I really need to make handwritten notes?

Ans. Yes, 100%. Writing helps you remember better than just reading. Plus, during revision, your notes save HOURS because everything important is in one place.

Q5. How do I get better at Organic Chemistry?

Ans. Start with the basics - nucleophiles, electrophiles, resonance, inductive effect. Once you understand why a reaction happens, the whole chapter becomes easier. Also, practise predicting products. It builds confidence super fast.