Class 10 NBF Chemistry(New Book) Notes – Chapters Exercises, MCQs, Short & Long Questions

Chapter 1: History of Chemistry

1.1 Multiple Choice Questions

1. C. Mass remains constant during a chemical reaction
2. B. Accuracy and validity of findings
3. B. Phlogiston theory
4. B. Elements by their properties and atomic number
5. B. There is a 5% chance the results are incorrect
6. B. Rutherford model
7. B. Same results under the same conditions
8. B. Same results using different methods
9. B. Theory of combustion
10. B. Properties of elements

1.2 Short Response Questions

1. Mass is neither created nor destroyed in chemical reactions. Total mass of reactants equals total mass of products because atoms are rearranged, not changed in number (Lavoisier).
2. Empirical evidence comes from direct observation and experimentation. It forms the basis for hypotheses, testing, and conclusions, ensuring science is grounded in measurable reality.
3. Independent experts evaluate research for methodology, validity, and significance before publication. It ensures quality, detects errors, and maintains scientific integrity.
4. Combustible materials contain “phlogiston” released during burning. Air absorbs phlogiston; combustion stops when air is saturated.
5. Rutherford discovered the nucleus — tiny, dense, positive center with most mass, electrons orbiting at large distances. Explained alpha scattering and revolutionized atomic theory.
6. Elements are arranged by increasing atomic number in periods and groups. Same group elements have similar properties due to same valence electrons.
7. A paradigm is a shared framework of concepts guiding research. Example: phlogiston theory dominated 18th-century chemistry until replaced by oxygen theory.
8. It shows the probability that results are statistically significant (e.g., 95% confidence means only 5% chance results are due to random variation).
9. Repeatability: same scientist, same conditions → same results. Reproducibility: different scientists, different methods → same results.
10. Skepticism demands evidence, questions claims, and prevents flawed ideas. It promotes rigorous testing and self-correction in science.

1.3 Long Response Questions

1. The shift from phlogiston to oxygen theory was a major paradigm change in chemistry. Phlogiston theory (17th–18th century) claimed combustible materials contain “phlogiston” released during burning. Metals were compounds of phlogiston and calx (oxide); burning released phlogiston, leaving calx. Air absorbed phlogiston; combustion stopped when air was saturated. Problems: metals gained mass during calcination (contradicting phlogiston loss), and phlogiston was given negative mass to explain this — increasingly contrived. Antoine Lavoisier (late 18th century) used quantitative experiments in sealed containers. He showed combustion involves combination with air’s active part (oxygen). Burning phosphorus/sulfur gained mass equal to air absorbed. Metals gained mass from oxygen forming oxides. Respiration consumed oxygen. Lavoisier named oxygen (“acid former”) and stated combustion is oxidation. His 1789 textbook established modern chemistry. Impact: introduced mass conservation as fundamental law; made chemistry quantitative; explained oxidation-reduction; enabled stoichiometry; replaced qualitative phlogiston with precise oxygen theory; laid foundation for atomic theory and modern chemical nomenclature.
2. Atomic models evolved with new evidence. Thomson’s 1904 plum pudding model: atom as uniform positive sphere with embedded electrons (neutrality explained). Failed to explain alpha scattering. Rutherford’s 1911 gold foil experiment: most alpha particles passed through, some deflected/bounced back → atom mostly empty with tiny, dense, positive nucleus (most mass) and orbiting electrons. Explained scattering via Coulomb repulsion but couldn’t explain electron stability (classical physics predicted spiraling into nucleus). Bohr’s 1913 model: electrons in fixed energy levels/shells; jumps emit/absorb photons. Explained hydrogen spectrum but failed for multi-electron atoms. Quantum mechanical model (1920s, Schrödinger, Heisenberg): electrons as probability clouds/orbitals (wave-particle duality). Heisenberg uncertainty principle: position and momentum cannot be known precisely. Quantum numbers describe electron states. Model predicts behavior of all elements accurately. Each model incorporated new evidence, progressing from uniform sphere to nuclear structure to probabilistic quantum description.
3. Mendeleev’s 1869 periodic table arranged elements by atomic weight, grouping similar properties and predicting undiscovered elements (e.g., gallium, scandium, germanium matched predictions). Moseley’s 1913 atomic number refinement solidified it. As paradigm, it organizes knowledge, predicts properties (group trends: alkali metals reactive, halogens form salts), reveals periodic trends (radius, ionization energy, electronegativity). Guides synthesis, bonding understanding, and new element searches (transuranium elements followed table predictions). Reveals electron configuration link to properties (valence electrons determine reactivity). Remains chemistry’s central organizing tool, validated by discoveries and guiding research.
4. Repeatability (same scientist, identical conditions → consistent results) shows precision and reliability. Reproducibility (different scientists, methods → same results) confirms robustness and generality. Both prevent errors, bias, fraud. Repeatability detects experimental variation; reproducibility validates broadly. Replication crises (psychology, some chemistry) led to preregistration, data sharing, negative result publication. In chemistry, reproducible syntheses, analyses, and theories are essential for trust and application (pharma, industry). Ensures cumulative knowledge, self-correction, and progress.
5. Uncertainty arises from instrument limits, sample variation, etc. Confidence intervals show range likely containing true value (95% confidence: 95% of intervals from repeated experiments contain true value). Example: calcium in water 45.2 ± 0.3 mg/L (95% CI) → true value between 44.9–45.5 mg/L. Standard deviation measures data spread; small → high precision. Standard error of mean decreases with more measurements. Calibration curves have error bars/confidence bands. Uncertainty propagates in calculations. t-tests assess significance at confidence levels. Expressing uncertainty shows rigor, allows proper interpretation, and prevents over-confidence in results.

1.4 Summary of Key Concepts

Concept	Description
Conservation of Mass	Mass constant in reactions (Lavoisier)
Phlogiston Theory	18th-century combustion paradigm
Oxygen Theory	Lavoisier’s replacement: combustion with oxygen
Rutherford Model	Nucleus with orbiting electrons
Periodic Table	Organizes elements by atomic number, properties
Peer Review	Expert evaluation for quality/validity
Repeatability	Same scientist, same conditions, same results
Reproducibility	Different scientists, different methods, same results
Confidence Level	Statistical reliability (e.g., 95%)
Empirical Evidence	Knowledge from observation/experiment

Chapter 2: Matter

1.1 Multiple Choice Questions

1. C. It increases
2. C. Sublimation
3. B. A fixed temperature
4. B. Lose energy and move slower
5. B. Charles’s Law
6. A. Evaporation occurs at the surface of a liquid at any temperature
7. D. The temperature increases
8. D. Only temperature and molecular mass
9. C. Sublimation
10. B. It ensures the medicine reaches all parts of the body

1.2 Short Response Questions

1. Heat energy is used to break intermolecular forces, allowing molecules to escape as gas without raising temperature.
2. Particles in constant motion; kinetic energy determines state.
3. Cooling curve plots temperature vs time during cooling.
4. Diffusion is spontaneous spreading of particles from high to low concentration.
5. Evaporation at surface, any temperature; boiling throughout, specific temperature.
6. Graham’s law: rate inversely proportional to square root of molar mass.
7. Sublimation: solid to gas without liquid phase.
8. Charles’s law: volume proportional to temperature at constant pressure.
9. Boyle’s law: pressure inversely proportional to volume at constant temperature.
10. Kinetic theory explains states by particle motion and forces.

1.3 Long Response Questions

1. Kinetic particle theory states particles in constant motion; energy determines state.
2. Cooling curve plots temperature vs time; plateaus at phase changes.
3. Diffusion is spontaneous spreading; factors: temperature, mass, medium.

1.4 Think Tank

1. Boiling occurs when vapor pressure equals atmospheric; factors affect rate.
2. Diffusion rates critical in drug development for absorption.
3. Sublimation offers controlled phase change in pharma.

1.5 Summary of Key Concepts

Concept	Description
Kinetic Particle Theory	Particles in constant motion; energy determines state
Evaporation	Liquid to gas at surface
Boiling	Liquid to gas throughout
Sublimation	Solid to gas
Diffusion	Spreading from high to low concentration
Graham’s Law	Rate ∝ 1/√molar mass
Charles’s Law	V ∝ T
Boyle’s Law	P ∝ 1/V

Chapter 3: Stoichiometry

1.1 Multiple Choice Questions

1. B. 1.0
2. B. 24 dm³
3. D. 7 moles of atoms in total
4. A. 106 g/mol
5. C. Limiting reactant
6. A. NH₃
7. D. 3.01 × 10²³
8. B. Diamond
9. D. Number of molecules
10. B. 13 mol

1.2 Short Questions

1. Assume 100 g sample: C: 65.45 g / 12 = 5.454 mol; H: 5.45 g / 1 = 5.45 mol; N: 29.1 g / 14 = 2.078 mol. Divide by smallest: C: 2.626, H: 2.623, N: 1 → empirical C₅H₅N₂ (approx).
2. Molar mass = 12n + 1n = 13n = 65 → n = 5, molecular C₅H₅.
3. Moles = mass / molar mass.
4. Limiting reactant limits product amount.
5. Percentage yield = (actual / theoretical) × 100.
6. Avogadro’s number = 6.02 × 10²³.
7. 1 dm³ = 1000 cm³.
8. STP: 0°C, 1 atm; molar volume 22.4 dm³.
9. Allotropes: different forms same element (diamond/graphite).
10. Isotopes: same protons, different neutrons.
11. Empirical from percentages; molecular from molar mass.
12. Stoichiometry: quantitative relations in reactions.
13. Theoretical yield from balanced equation.
14. Actual yield < theoretical due to losses.

1.3 Think Tank

1. Percentage yield rarely 100% due to side reactions, losses.
2. Add additional solute to reach 1.0 mol/dm³.
3. Na₂CO₃ + 2HCl → products.
4. M(H₂SO₄) = 98 g/mol.
5. Acid-base titration with standardized NaOH.

1.4 Summary of Key Formulas

Formula	Description
Moles = mass / molar mass	Calculate moles from mass
Percentage composition	(Mass element / molar mass) × 100
Empirical formula	Simplest ratio from percentages
Molecular formula	Empirical × n (n = molar mass / empirical mass)
Limiting reactant	Reactant with least product moles
Percentage yield	(Actual / theoretical) × 100
Avogadro’s number	6.02 × 10²³ particles/mol
Molar volume STP	22.4 dm³/mol

Chapter 4: Electrochemistry

1.1 Multiple Choice Questions

1. C. Electrons flow from cathode to anode
2. A. Half-cell of an active metal acts as a cathode
3. D. Reduction of metal oxide by a reducing agent
4. B. Coating with Zn
5. B. Chlorine gas is produced at anode
6. B. Coating a metal object with a thin layer of metal through electrolysis
7. B. Maintaining electrical neutrality by allowing ion movement
8. B. Metal B is more reactive than metal A

1.2 Short Questions

1. A fuel cell is an electrochemical cell converting chemical energy to electricity.
2. Oxidation: loss of electrons; reduction: gain.
3. Redox: oxidation + reduction.
4. Galvanization: zinc coating to prevent rust.
5. Electroplating: depositing metal layer via electrolysis.
6. Salt bridge: maintains neutrality.
7. Cathode: reduction; anode: oxidation.
8. Electrolytic cell: non-spontaneous, needs energy.

1.3 Long Questions

1. Oxidation: gain oxygen/loss electrons; examples.
2. (sketch description)
3. Lead from PbCl₂ via electrolysis.
4. With copper electrodes: reactions.
5. (a) (sketch) (b) Electrolysis non-spontaneous.
6. a. Anode: Cu → Cu²⁺ + 2e⁻.
7. Cathode: water reduced.
8. Molten PbCl₂: Pb cathode, Cl₂ anode.
9. Molten PbCl₂ electrolysis.
10. (i) H₂, O₂ (ii) Cu, Cl₂ (iii) Pb, Br₂
11. Conc NaCl (inert): NaOH, Cl₂, H₂.

1.4 Page 62 Questions

1. Fe oxidized, O reduced.
2. (i) Br₂ (ii) I₂ vapors.
3. (i) H₂, O₂ (ii) Cu, Cl₂ (iii) Pb, Br₂

1.5 Summary of Key Concepts

Concept	Description
Oxidation	Loss of electrons/gain oxygen
Reduction	Gain electrons/loss oxygen
Redox	Oxidation + reduction
Galvanic Cell	Spontaneous redox
Electrolytic Cell	Non-spontaneous, driven by electricity
Galvanization	Zn coating for corrosion protection
Electroplating	Metal deposition via electrolysis
Salt Bridge	Ion flow for neutrality

Chapter 5: Chemical Kinetics

1.1 Multiple Choice Questions

1. B. Decreases
2. D. None of these
3. B. Slow
4. C. Forward as well as reverse reaction
5. D. Molar mass of reactants
6. D. By providing an alternate pathway with lower activation energy
7. D. Colour of the reactants
8. B. Change in pressure
9. B. Increases the rate of reaction
10. D. Higher temperature increases the frequency of collisions and kinetic energy of particles

1.2 Short Questions

1. (energy diagram)
2. Collision theory: reactions from collisions with energy.
3. Activation energy: minimum for reaction.
4. Catalyst: lowers activation energy.
5. Temperature increases rate by more collisions/energy.
6. Concentration increases rate by more collisions.
7. Surface area increases rate for solids.
8. Pressure increases rate for gases.
9. Light increases rate in photochemical.
10. Reversible: forward/reverse.
11. Forward: reactants to products.
12. Reverse: products to reactants.
13. Rate law: rate = k [reactants]^n
14. Order: exponent in rate law.
15. Half-life: time for half concentration.

1.3 Long Questions

1. Concentration vs time graph.
2. Supported by collision theory.
3. Catalyst increases rate, alternate path.
4. Temperature profound effect.
5. Kinetics in food industry.
6. Hypothesis: temp/surface area effect.
7. (not in query)
8. Catalyst in pharma.
9. Objective: compare pharma rate.
10. Maximize rate per collision theory.
11. Similarities: increase rate.
12. Activation energy importance.

1.4 Think Tank

(a) Catalyst ineffective scenarios.
(b) Kinetics in food production.
(c) Double surface double rate.
(d) Gas formation associated.

1.5 Summary of Key Concepts

Concept	Description
Collision Theory	Reactions occur when particles collide with sufficient energy
Activation Energy	Minimum energy for reaction
Catalyst	Lowers activation energy
Rate Law	Rate = k [A]^m [B]^n
Order of Reaction	Sum of exponents in rate law
Half-Life	Time for concentration to halve

Chapter 6: Salts

1.1 Multiple Choice Questions

1. C. Sodium nitrate (NaNO₃)
2. A. Negative ions
3. B. Strong electrostatic forces
4. C. Molten
5. C. Chlorides are soluble except lead and silver chlorides
6. C. They become mobile
7. D. Acid + Alkali
8. D. Lead chloride (PbCl₂)
9. B. Salt and water
10. B. NO₃⁻

1.2 Short Questions

1. Salt: ionic compound from acid H⁺ replaced by metal.
2. Normal: complete H⁺ replacement.
3. Acidic: partial H⁺ replacement.
4. Basic: from weak acid/strong base.
5. Double: two cations/anions.
6. Mixed: two acids/bases.
7. Soluble: nitrates, Na/K salts.
8. Insoluble: Ag/Pb chlorides.
9. Preparation methods: titration, precipitation.

1.3 Long Questions

1. Titration for soluble salts.
2. Salts solid at STP.
3. Solubility rules predict dissolution.
4. Ionic lattice from bonding.
5. Method for insoluble carbonate.

1.4 Think Tank

1. Prepare MgSO₄ crystals.
2. KOH + H₂SO₄ titration.
3. Always soluble: nitrates, Na salts.
4. Prepare CuSO₄ from CuCO₃.
5. CuSO₄ from H₂SO₄.

1.5 Summary of Key Concepts

Concept	Description
Salt	Ionic compound from acid H⁺ replaced by metal
Normal Salt	All H⁺ replaced
Acid Salt	Partial H⁺ replacement
Basic Salt	Weak acid/strong base
Solubility Rules	Predict salt dissolution
Titration	For soluble salts
Precipitation	For insoluble salts

Chapter 7: Nitrogen, Sulphur and Metals

1.1 Multiple Choice Questions

1. B. Peroxyacetyl nitrate (PAN)
2. D. 2N₂(g) + 3H₂(g) ⇌ 2NH₃(g)
3. B. Methane and air
4. C. Iron
5. B. 450°C
6. B. 2SO₂(g) + O₂(g) ⇌ 2SO₃(g)
7. B. Vanadium(V) oxide
8. A. Burning sulphur or roasting sulphide ores
9. A. Forming a salt and water
10. D. Al₂O₃

1.2 Short Questions

1. PAN: secondary pollutant from NOₓ/VOCs.
2. NOₓ: air pollutants from combustion.
3. Haber process: ammonia production.
4. Contact process: H₂SO₄ production.
5. Acidic oxide: reacts with base.
6. Basic oxide: reacts with acid.
7. Amphoteric oxide: both acid/base.
8. Neutral oxide: no reaction.
9. Reactivity series: metals by reactivity.
10. Extraction: from ores.
11. Corrosion: metal deterioration.
12. Galvanization: Zn coating.

1.3 Long Questions

1. NOₓ pollutants: sources, effects.
2. Haber process details.
3. Contact process details.
4. Oxides classification.
5. Haber/Contact environmental impact.
6. NO/N₂O acid rain contribution.
7. Oxide acidity by position.

1.4 Think Tank

1. Gold/platinum for jewellery.
2. New metal in reactivity series.
3. Reactivity series A>B>C.
4. Basic oxides uses.
5. Haber 450°C/200 atm choice.
6. Acidic oxides environmental impact.
7. Controlling NOₓ/SO₂ emissions.

1.5 Summary of Key Concepts

Concept	Description
Haber Process	N₂ + 3H₂ ⇌ 2NH₃, Fe catalyst, 450°C, 200 atm
Contact Process	SO₂ to SO₃, V₂O₅ catalyst
Acidic Oxide	Reacts with base
Basic Oxide	Reacts with acid
Amphoteric Oxide	Reacts with both
Reactivity Series	K > Na > Ca > Mg > Al > Zn > Fe > Sn > Pb > H > Cu > Ag > Au
Corrosion	Metal oxidation

Chapter 8: Organic Chemistry

1.1 Multiple Choice Questions

1. A. CH₃CH=CHCH₃
2. B. Butan-1-ol
3. A. Butanoic acid
4. B. They have the same molecular formula
5. A. Methyl propanoate
6. C. Butan-1-ol
7. B. -COOH
8. B. CH₃COOCH₂CH₂CH₂CH₃
9. A. CH₃CH₂CH₂CH₂OH
10. B. Propan-2-ol
11. A. But-2-ene: CH₃CH=CHCH₃

1.2 Short Questions

1. Displayed structural formula shows arrangement.
2. IUPAC naming rules.
3. Alkene: C=C.
4. Alcohol: -OH.
5. Carboxylic acid: -COOH.
6. Ester: -COO-.
7. Structural isomers: same formula, different structure.
8. Functional isomers: same formula, different group.
9. Position isomers: group different position.
10. Chain isomers: carbon chain different.
11. Esterification: alcohol + acid → ester.
12. Saponification: ester hydrolysis.
13. Homologous series: similar properties.
14. Organic compounds classification.

1.3 Long Questions

1. Functional groups determine properties.
2. Distinguish alkenes/alcohols/acids.
3. Alcohol + carboxylic acid reaction.
4. Ethanol + propanoic acid esterification.
5. Methyl methanoate etc.
6. Suffixes indicate type.
7. Esters/carboxylic acids differences.
8. (table)

1.4 Think Tank

1. Decision tree for identification.
2. C₄H₁₀O isomers.

1.5 Summary of Key Concepts

Compound Type	Functional Group	Suffix	Example
Alkane	None	-ane	Butane (C₄H₁₀)
Alkene	C=C	-ene	Butene (C₄H₈)
Alcohol	-OH	-ol	Butanol (C₄H₉OH)
Carboxylic Acid	-COOH	-oic acid	Butanoic acid (C₃H₇COOH)
Ester	-COO-	-oate	Methyl butanoate (C₃H₇COOCH₃)

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