Chapter 1: Some Basic Concepts of Chemistry Long Questions | chemistry Class 11 CBSE

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Chapter 1: Some Basic Concepts of Chemistry Long Questions | chemistry Class 11 CBSE | ByteNotes.in

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Question 1

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Discuss the four postulates of Dalton's Atomic Theory. Which laws of chemical combination does it explain, and what are its limitations?

Dalton's four postulates (1808): matter consists of indivisible atoms; all atoms of a given element are identical in mass and properties; atoms of different elements differ in mass; compounds form when atoms of different elements combine in fixed ratios; and chemical reactions involve reorganisation of atoms that are neither created nor destroyed.
The theory successfully explains the Law of Conservation of Mass because atoms are neither created nor destroyed in reactions, so total mass remains constant.
It explains the Law of Definite Proportions because atoms combine in fixed, whole-number ratios to form compounds, giving constant elemental mass ratios.
It explains the Law of Multiple Proportions because when two elements form more than one compound, atoms combine in different but simple fixed ratios, giving small whole-number mass ratios.
Dalton's theory could not explain Gay Lussac's Law of Gaseous Volumes, as it failed to account for the simple volume relationships observed when gases react.
Additionally, it could not explain why atoms combine at all; it also incorrectly assumed atoms are indivisible (later disproved by discovery of subatomic particles) and identical (isotopes show otherwise).

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Question 2

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Explain the mole concept in detail. How is molar mass related to molecular mass? Calculate the moles and number of molecules in 36 g of water.

The mole is the SI unit of amount of substance; one mole contains exactly 6.02214076 × 1023 elementary entities (atoms, molecules, ions, electrons, etc.), known as the Avogadro constant (NA).
Just as a dozen represents 12 items, a mole represents 6.022 × 1023 particles; this unit was introduced to handle the extremely large numbers of atoms and molecules involved in chemical reactions.
Molar mass is the mass of one mole of a substance in grams and is numerically equal to the atomic, molecular, or formula mass in unified mass units (u); e.g., molar mass of H2O = 18.02 g mol-1.
The relationship links macroscopic mass (grams) to microscopic entities: moles = mass (g) / molar mass (g mol-1); number of particles = moles × NA.
For 36 g of water: moles of H2O = 36 g / 18.02 g mol-1 ≈ 2 mol; number of molecules = 2 × 6.022 × 1023 = 1.204 × 1024 molecules.
The Avogadro number (6.0221367 × 1023) was determined from the mass of a carbon-12 atom: 12 g/mol divided by 1.992648 × 10-23 g/atom equals 6.022 × 1023 atoms/mol.

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Question 3

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Compare and contrast the five laws of chemical combination, explaining how each one contributed to the development of Dalton's atomic theory.

Law of Conservation of Mass (Lavoisier, 1789): matter is neither created nor destroyed. This was foundational — Dalton incorporated it by postulating that atoms are neither created nor destroyed in reactions.
Law of Definite Proportions (Proust): a compound always contains elements in a fixed mass ratio. This led Dalton to propose that atoms of different elements combine in fixed, definite ratios to form compounds.
Law of Multiple Proportions (Dalton, 1803): when two elements form multiple compounds, the mass ratios are simple integers. Dalton used this directly to demonstrate that atoms combine in small, discrete, whole-number ratios.
Gay Lussac's Law of Gaseous Volumes (1808): gases combine in simple volume ratios at same temperature and pressure. Though Dalton's theory could not fully explain this law, it reinforced the idea of discrete atomic/molecular combinations.
Avogadro's Law (1811): equal volumes of gases at the same T and P contain equal numbers of molecules. This resolved conflicts in Dalton's theory by distinguishing atoms from molecules and explaining gaseous volume ratios.
Together, these five laws demonstrate a consistent pattern of discrete, whole-number atomic combinations that underpins all of Dalton's postulates, establishing the atomic model as a rational explanation for the quantitative behaviour of matter.

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Question 4

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Describe the steps involved in determining the empirical formula and molecular formula of a compound from percentage composition data. Apply these steps to a compound containing 4.07% H, 24.27% C, and 71.65% Cl with molar mass 98.96 g.

Step 1: Convert mass percentages to grams by assuming 100 g of the compound as the starting material; thus, 4.07 g H, 24.27 g C, and 71.65 g Cl are present.
Step 2: Convert grams to moles by dividing by the atomic mass of each element: moles of H = 4.07/1.008 = 4.04; moles of C = 24.27/12.01 = 2.021; moles of Cl = 71.65/35.453 = 2.021.
Step 3: Divide all mole values by the smallest value (2.021) to get the simplest ratio: H = 4.04/2.021 ≈ 2; C = 2.021/2.021 = 1; Cl = 2.021/2.021 = 1. Ratio H:C:Cl = 2:1:1.
Step 4: Write the empirical formula: CH2Cl (carbon, two hydrogens, one chlorine).
Step 5: Calculate the empirical formula mass = 12.01 + 2(1.008) + 35.453 = 49.48 g. Divide molar mass by empirical formula mass: n = 98.96/49.48 ≈ 2.
Step 6: Multiply empirical formula by n = 2 to get the molecular formula: C2H4Cl2.

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Question 5

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Explain the classification of matter into elements, compounds, and mixtures with suitable examples. How can the components of a mixture be separated?

Matter at the macroscopic level is classified into pure substances and mixtures. Pure substances have fixed composition and characteristic properties; mixtures have variable composition.
Elements are pure substances whose particles consist of only one type of atom; examples include sodium (Na), copper (Cu), oxygen (O2), and hydrogen (H2). Their atoms may exist as individual atoms or molecules.
Compounds are pure substances formed when atoms of two or more elements combine in a fixed and definite ratio; the properties of a compound differ from its constituent elements. Examples: water (H2O), ammonia (NH3), carbon dioxide (CO2).
Mixtures contain two or more pure substances in any ratio; their components retain their individual properties. Homogeneous mixtures (e.g., sugar solution, air) have uniform composition; heterogeneous mixtures (e.g., salt and sugar, grains and dirt) have non-uniform composition.
Components of a mixture are separated by physical methods appropriate to the type of mixture: hand-picking (solid mixtures), filtration (solid-liquid), crystallisation (dissolved solids), distillation (liquids with different boiling points), and evaporation.
Constituents of compounds cannot be separated by physical methods — chemical methods are required; this is a fundamental distinction between compounds and mixtures.

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Question 6

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Analyse the contribution of ancient Indian scientists and texts to the development of chemistry, with specific reference to Acharya Kanda, Nagarjuna, and Charaka Samhita.

Ancient India called chemistry Rasayan Shastra, Rastantra, Ras Kriya, or Rasvidya, which encompassed metallurgy, medicine, cosmetics, glass, and dye manufacture — evidence from Mohenjodaro and Harappa demonstrates this antiquity.
Acharya Kanda (born 600 BCE, originally named Kashyap) was the first proponent of atomic theory, formulating the concept of 'Paramanu' — indivisible, indestructible, spherical particles in motion — 2500 years before Dalton. He authored Vaiseshika Sutras and proposed that Paramanu could form pairs and triplets.
Nagarjuna was a reputed chemist, alchemist, and metallurgist; his work Rasratnakar dealt with mercury compound formulations and methods for extracting gold, silver, tin, and copper. The book Rsarnavam (800 CE) described furnaces, ovens, crucibles, and methods of identifying metals by flame colour.
Charaka Samhita, the oldest Ayurvedic text, described preparation of sulphuric acid, nitric acid, oxides and sulphates of copper, tin, zinc, iron, and carbonates of lead and iron. It also described the concept of reducing particle size of metals — an early reference to what is now nanotechnology via bhasmas (shown to contain nanoparticles).
Other contributions: Chakrapani discovered mercury sulphide and invented soap using mustard oil and alkalies. Tamil texts described fireworks using sulphur, charcoal, saltpetre, mercury, and camphor. Rasopanishada described gunpowder preparation.
The Rigveda (1000–400 BCE) mentioned tanning and dyeing; Varāhmihir's Brihat Samhita described perfumes, cosmetics, and construction materials; these indicate a high scientific tradition predating modern Western chemistry.

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Question 7

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Discuss the SI system of units in detail. Why was a common system needed, and what are the seven base units with their definitions?

Before SI, different systems (English and Metric) were used in different parts of the world, leading to inconsistency in scientific communication; a common standard was required to ensure reproducibility and comparability of measurements globally.
The SI system (Système International d'Unités) was established in 1960 by the 11th General Conference on Weights and Measures (CGPM); it is maintained by the Metre Convention treaty signed in Paris in 1875, with India's standards maintained by the National Physical Laboratory (NPL), New Delhi.
Seven base units: metre (m) for length — defined via the speed of light in vacuum; kilogram (kg) for mass — defined via Planck's constant; second (s) for time — defined via caesium-133 atom's hyperfine transition frequency.
Ampere (A) for electric current — defined via elementary charge; Kelvin (K) for thermodynamic temperature — defined via Boltzmann constant; mole (mol) for amount of substance — defined as containing 6.02214076 × 1023 entities; candela (cd) for luminous intensity.
The SI system uses prefixes (nano, micro, milli, kilo, mega, etc.) to express multiples and submultiples of units, from yocto (10-24) to yotta (1024), enabling expression of both atomic-scale and astronomical-scale quantities.
All other physical quantities (speed, volume, density, pressure, energy) are derived from these seven base units; the system ensures universal scientific communication and underpins all quantitative chemistry.

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Question 8

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Explain how significant figures are used to indicate uncertainty in measurements. Describe all rules for counting significant figures with examples.

Every experimental measurement has some uncertainty due to limitations of measuring instruments and the observer's skill; significant figures (s.f.) are used to communicate this uncertainty — they include all certain digits plus one uncertain (estimated) digit.
Rule 1: All non-zero digits are significant. Example: 285 cm has 3 s.f.; 25 mL has 2 s.f. Rule 2: Zeros preceding the first non-zero digit are not significant (position indicators). Example: 0.03 has 1 s.f.; 0.0052 has 2 s.f.
Rule 3: Zeros between two non-zero digits are significant. Example: 2.005 has 4 s.f. Rule 4: Zeros at the end of a number are significant only if they are to the right of a decimal point. Example: 0.200 g has 3 s.f.; 100 without decimal has 1 s.f. but 100.0 has 4 s.f.
Rule 5: Exact counting numbers (e.g., 2 balls, 20 eggs) have infinite significant figures. Numbers in scientific notation: all digits are significant (e.g., 4.01 × 102 has 3 s.f.).
For addition/subtraction: the result has the same number of decimal places as the measurement with fewest decimal places. Example: 12.11 + 18.0 + 1.012 = 31.122 → reported as 31.1.
For multiplication/division: the result has the same number of significant figures as the measurement with fewest s.f. Example: 2.5 × 1.25 = 3.125 → reported as 3.1 (since 2.5 has only 2 s.f.). Rounding rules: if the digit removed is >5, increase the preceding digit; if <5, leave it unchanged; if exactly 5, round to the nearest even number.

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Question 9

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Explain stoichiometry using the combustion of methane as an example. How can a balanced equation be used to calculate masses of reactants and products?

Stoichiometry (from stoicheion = element, metron = measure) is the quantitative study of reactants and products in chemical reactions; it relies on balanced equations and molar relationships.
Balanced equation for combustion of methane: CH4(g) + 2O2(g) → CO2(g) + 2H2O(g). Stoichiometric coefficients indicate molar ratios: 1 mol CH4 reacts with 2 mol O2 to give 1 mol CO2 and 2 mol H2O.
Mass relationships: 16 g CH4 reacts with 64 g O2 to produce 44 g CO2 and 36 g H2O; total mass is conserved (16 + 64 = 44 + 36 = 80 g), consistent with the Law of Conservation of Mass.
Calculation example: 16 g CH4 = 1 mol; 1 mol CH4 produces 2 mol H2O = 2 × 18 g = 36 g water. Using unit factors: 2 mol H2O × (18 g H2O / 1 mol H2O) = 36 g H2O.
The conversion pathway is: mass ⇌ moles ⇌ number of molecules; the molar mass connects grams to moles, and Avogadro's constant connects moles to number of particles.
When reactions involve unequal amounts of reactants, the limiting reagent must be identified first; the amount of product is calculated based on the limiting reagent, not the excess reactant.

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Question 10

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Differentiate between atomic mass, average atomic mass, molecular mass, and formula mass. Explain why formula mass is used for NaCl instead of molecular mass.

Atomic mass is the mass of one atom of an element expressed relative to one-twelfth of the mass of a 12C atom; hydrogen has atomic mass ≈ 1.008 u and oxygen ≈ 16.00 u.
Average atomic mass is the weighted mean of all naturally occurring isotopes of an element based on their relative abundances; for carbon: (0.98892 × 12) + (0.01108 × 13.00335) + (~0 × 14.00317) = 12.011 u.
Molecular mass is the sum of atomic masses of all atoms in a molecule of a covalent compound; for methane (CH4): 12.011 + 4(1.008) = 16.043 u; for water (H2O): 2(1.008) + 16.00 = 18.02 u.
Formula mass is the sum of atomic masses of all atoms in the empirical formula of an ionic compound; it is used because ionic compounds do not exist as discrete molecules but as extended lattice structures.
In sodium chloride (NaCl), one Na+ ion is surrounded by six Cl- ions and vice versa in a three-dimensional lattice; there is no individual 'NaCl molecule'. Formula mass of NaCl = 23.0 + 35.5 = 58.5 u.
Using 'molecular mass' for NaCl would be chemically incorrect because it implies discrete molecules; 'formula mass' correctly refers to the simplest repeating formula unit of the ionic lattice, making it the appropriate term.

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Chapter 1: Some Basic Concepts of Chemistry | Long Questions

Discuss the four postulates of Dalton's Atomic Theory. Which laws of chemical combination does it explain, and what are its limitations?

Explain the mole concept in detail. How is molar mass related to molecular mass? Calculate the moles and number of molecules in 36 g of water.

Compare and contrast the five laws of chemical combination, explaining how each one contributed to the development of Dalton's atomic theory.

Describe the steps involved in determining the empirical formula and molecular formula of a compound from percentage composition data. Apply these steps to a compound containing 4.07% H, 24.27% C, and 71.65% Cl with molar mass 98.96 g.

Explain the classification of matter into elements, compounds, and mixtures with suitable examples. How can the components of a mixture be separated?

Analyse the contribution of ancient Indian scientists and texts to the development of chemistry, with specific reference to Acharya Kanda, Nagarjuna, and Charaka Samhita.

Discuss the SI system of units in detail. Why was a common system needed, and what are the seven base units with their definitions?

Explain how significant figures are used to indicate uncertainty in measurements. Describe all rules for counting significant figures with examples.

Explain stoichiometry using the combustion of methane as an example. How can a balanced equation be used to calculate masses of reactants and products?

Differentiate between atomic mass, average atomic mass, molecular mass, and formula mass. Explain why formula mass is used for NaCl instead of molecular mass.

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