50 Short Answer Questions — Atoms and Molecules
Each question below is followed by a clear, exam-focused short answer to aid revision. Questions are organized topic-wise and strictly follow NCERT Class 9 syllabus.
1. Laws of Chemical Combination
Q1. Explain the law of conservation of mass with an example.
Answer: The law states that mass is neither created nor destroyed in a chemical reaction; total mass of reactants equals total mass of products. Example: When hydrogen burns in oxygen to form water, the combined mass of H₂ and O₂ before reaction equals mass of the water produced (ignoring mass loss as gas/vapour).
Q2. What is the law of definite proportions and why is it important?
Answer: It states that a given compound always contains the same elements in the same fixed mass ratio. Importance: It implies compounds have definite composition (e.g., water always has hydrogen and oxygen in mass ratio ≈1:8), supporting the existence of discrete entities (atoms) combining in fixed ratios.
Q3. Describe the law of multiple proportions with a suitable chemical example.
Answer: When two elements form more than one compound, the masses of one element that combine with a fixed mass of the other are in small whole-number ratios. Example: Carbon and oxygen form CO and CO₂. If 12 g C combines with 16 g O in CO, in CO₂ 12 g C combines with 32 g O; ratio 16:32 = 1:2.
2. Dalton's Atomic Theory
Q4. List four main postulates of Dalton's atomic theory.
Answer: (1) Matter is made of small indivisible particles called atoms. (2) Atoms of an element are identical in mass and properties (classical postulate). (3) Compounds form by combination of atoms in simple whole-number ratios. (4) Atoms are not created or destroyed in chemical reactions.
Q5. Which aspect of Dalton’s theory is modified by modern chemistry and why?
Answer: Dalton’s claim that all atoms of an element are identical is modified; isotopes exist — atoms of the same element with different masses due to differing neutron numbers.
Q6. How did Dalton’s theory explain the laws of chemical combination?
Answer: It proposed that compounds are formed by specific combinations of atoms in whole-number ratios; this explains definite proportions and conservation of mass because atoms retain identity and combine in fixed numbers.
3. Atom and Molecule — Definitions & Differences
Q7. Define an atom and give an example.
Answer: An atom is the smallest unit of an element that can participate in chemical reactions. Example: a sodium atom (Na) is an atom of the element sodium.
Q8. Define a molecule and illustrate with an example.
Answer: A molecule is the smallest neutral particle of a substance that exhibits its chemical properties. Example: A water molecule (H₂O) consists of two hydrogen atoms and one oxygen atom bonded together.
Q9. How is a molecule of an element different from a molecule of a compound?
Answer: A molecule of an element consists of only one type of atom (e.g., O₂), while a molecule of a compound contains atoms of different elements chemically bonded (e.g., CO₂ contains C and O).
Q10. What is meant by a formula unit?
Answer: Formula unit refers to the simplest whole-number ratio of ions in an ionic compound and is used because ionic compounds do not exist as discrete molecules. Example: NaCl is a formula unit for sodium chloride.
4. Chemical Formulae & Representation
Q11. What does the chemical formula H₂SO₄ tell you?
Answer: It shows that each molecule (or formula unit) contains 2 hydrogen atoms, 1 sulfur atom and 4 oxygen atoms in that order; it indicates composition but not structure.
Q12. Explain how to interpret the formula Ca(OH)₂ in terms of atom counts.
Answer: Ca(OH)₂ contains one calcium atom, two oxygen atoms and two hydrogen atoms because the OH group appears twice: Ca =1, O =1×2=2, H=1×2=2.
Q13. Why are brackets used in chemical formulae?
Answer: Brackets group atoms that repeat as a unit; a subscript outside the bracket multiplies all atoms inside. Example: Al₂(SO₄)₃ indicates three sulfate groups.
5. Relative Atomic Mass and Relative Molecular Mass
Q14. Define relative atomic mass (Ar).
Answer: Ar is the average mass of atoms of an element compared to 1/12th the mass of carbon-12 atom; Ar values are dimensionless compared to the C-12 scale.
Q15. Define relative molecular mass (Mr) and how to calculate it.
Answer: Mr is the sum of relative atomic masses of atoms in a molecule. Calculate by adding Ar values of all atoms present (e.g., Mr of H₂O = 2×Ar(H)+Ar(O) = 2×1 + 16 =18).
Q16. Calculate Mr of C₆H₆ (Ar C=12, H=1).
Answer: Mr = 6×12 + 6×1 = 72 + 6 = 78.
6. Counting Atoms and Molecules
Q17. How many hydrogen atoms are there in 2 molecules of H₂SO₄?
Answer: Each H₂SO₄ has 2 H atoms, so 2 molecules contain 2×2 = 4 hydrogen atoms.
Q18. In 4 units of Al₂(SO₄)₃, how many sulphur atoms are present?
Answer: Each Al₂(SO₄)₃ contains 3 S atoms, so 4 units contain 4×3 = 12 sulphur atoms.
Q19. In the formula 5NH₃, how many N and H atoms are present?
Answer: NH₃ has 1 N and 3 H atoms; in 5NH₃: N = 5×1 = 5; H = 5×3 = 15.
7. Chemical Equations and Balancing
Q20. What is meant by balancing a chemical equation?
Answer: Balancing ensures equal numbers of each type of atom on both sides of the equation, satisfying the law of conservation of mass.
Q21. Balance the equation: Fe + O₂ → Fe₂O₃ and explain steps briefly.
Answer: Balanced equation is 4Fe + 3O₂ → 2Fe₂O₃. Steps: Count atoms, place coefficients: Fe: 4 on left gives 4 in product (2×2), O: 3×2=6 atoms on right matched by 3O₂ on left.
Q22. Why are coefficients not the same as subscripts?
Answer: Coefficients indicate number of molecules or formula units; subscripts indicate number of atoms in a single molecule. Changing subscripts changes the substance; changing coefficients balances the equation.
8. Classification: Elements, Compounds and Mixtures
Q23. How can you distinguish between an element and a compound?
Answer: An element contains only one type of atom and cannot be chemically broken into simpler substances (e.g., O₂). A compound contains atoms of two or more elements chemically combined in fixed ratios (e.g., H₂O) and can be decomposed chemically.
Q24. Give two differences between compounds and mixtures.
Answer: (1) Compounds have fixed composition; mixtures do not. (2) Compounds formed by chemical bonds and need chemical methods to separate; mixtures can often be separated by physical means like filtration or distillation.
Q25. Is air a compound? Explain.
Answer: No. Air is a homogeneous mixture of gases (mainly N₂ and O₂) and not a compound because its composition can vary and components are not chemically bonded.
9. Bonding — Ionic and Covalent (Introductory)
Q26. What type of bond forms between sodium and chlorine in NaCl and why?
Answer: Ionic bond forms because sodium donates an electron to chlorine, producing Na⁺ and Cl⁻ ions; electrostatic attraction between opposite charges holds them together.
Q27. Why do two hydrogen atoms form a covalent bond in H₂?
Answer: Each hydrogen atom requires one electron to attain the noble gas configuration; by sharing their electrons they form a covalent bond and achieve stability.
10. Atomic Structure Basics
Q28. Name the three main subatomic particles and state one property of each.
Answer: Proton (positive charge, in nucleus, relative mass ≈1), Neutron (no charge, in nucleus, relative mass ≈1), Electron (negative charge, orbits nucleus, negligible mass relative to proton).
Q29. Where are protons and neutrons located in an atom?
Answer: Both protons and neutrons are located in the nucleus at the center of the atom.
Q30. Which particle determines the identity of an element?
Answer: The number of protons (atomic number) determines the identity of an element.
11. Isotopes and Isobars — Simple Concepts
Q31. Define isotopes with an example.
Answer: Isotopes are atoms of the same element with same atomic number but different mass numbers due to differing neutron counts. Example: Carbon-12 and Carbon-14.
Q32. Do isotopes have identical chemical properties? Explain briefly.
Answer: Isotopes have nearly identical chemical properties because chemical behaviour depends on electron configuration, which is same for isotopes; small differences in reaction rates can occur due to mass differences.
12. Formula Calculations & Short Numerical Ideas
Q33. If Ar of Na = 23 and Cl = 35.5, calculate Mr of NaCl and explain significance.
Answer: Mr = 23 + 35.5 = 58.5. Significance: Mr helps compare mass of one formula unit of NaCl to another substance and is used for stoichiometric calculations.
Q34. How would you compute Mr of CaCO₃ (Ar: Ca=40, C=12, O=16)?
Answer: Mr = Ar(Ca) + Ar(C) + 3×Ar(O) = 40 +12 + 3×16 = 100.
Q35. Explain briefly why Mr is a dimensionless quantity.
Answer: Mr is a ratio comparing masses relative to 1/12 of carbon-12 mass, so it has no units; it is a relative (dimensionless) number.
13. Common Exam Questions and Tips
Q36. What should you always state when calculating Mr in an exam?
Answer: State the Ar values used for each element and show the addition steps to reach Mr; this helps examiners follow your working.
Q37. Give one memory tip to remember Ar values of common elements.
Answer: Use mnemonics or small flashcards for common Ar values (H=1, C=12, N=14, O=16, Na=23, Cl=35.5, S=32) and revise frequently.
Q38. How can you avoid mistakes while balancing equations?
Answer: Count atoms of each element on both sides systematically and adjust coefficients starting with the most complex molecule; recheck atom counts after placing coefficients.
14. Conceptual Short Questions
Q39. Why is understanding atomic theory important for chemistry?
Answer: Atomic theory explains why elements combine in fixed ratios, how chemical reactions conserve mass, and provides a foundation for stoichiometry and modern chemistry concepts.
Q40. Can atoms of different elements have the same mass? Give reason.
Answer: Different elements may have very similar Ar values but not identical atomic numbers; isotopes lead to variation; atomic mass depends on protons+neutrons, so different combinations can give similar masses.
Q41. Explain briefly how empirical and molecular formula differ (intro-level).
Answer: Empirical formula gives simplest whole-number ratio of atoms in a compound (e.g., CH₂ for C₂H₄), molecular formula gives actual number of atoms in a molecule (e.g., C₂H₄).
15. Final Practice & Recall
Q42. What is the atomic number and mass number?
Answer: Atomic number (Z) is number of protons in nucleus. Mass number (A) is total number of protons and neutrons in nucleus.
Q43. How would you represent an isotope of chlorine with mass number 37?
Answer: As ³⁷Cl (or using standard notation: ³⁷Cl with 17 as atomic number subscript if needed). Mass number 37 indicates total protons+neutrons=37.
Q44. Give one example of a monatomic molecule.
Answer: Noble gases like helium (He) exist as monatomic species; He is a monatomic 'molecule' in gaseous state.
Q45. Explain briefly why ionic compounds are represented by formula units rather than molecules.
Answer: Ionic compounds form extended lattices of ions rather than discrete molecules; formula unit represents the simplest ratio of ions in that lattice (e.g., NaCl represents 1:1 ratio).
Q46. If a compound has empirical formula CH₂ and Mr = 56, what is its molecular formula? (Show reasoning)
Answer: Empirical mass = 12 + 2×1 = 14. Mr / empirical mass = 56 / 14 = 4 ⇒ molecular formula = 4× empirical = C₄H₈.
Q47. Why is the relative atomic mass of chlorine listed as 35.5 in the periodic table?
Answer: Because chlorine occurs naturally as a mixture of isotopes (mainly Cl-35 and Cl-37); the Ar value is a weighted average of isotope abundances, resulting in ~35.5.
Q48. What happens to the number of atoms when a chemical reaction occurs?
Answer: The total number of each type of atom remains the same before and after reaction; atoms are rearranged to form new substances but not created or destroyed.
Q49. Give one experimental technique that supports the idea of atoms (historical/modern).
Answer: Brownian motion observation (by Einstein and Perrin) provided evidence for existence of atoms; modern techniques like scanning tunnelling microscopy also image surfaces at atomic scale.
Q50. Provide one practical exam tip for answering questions from this chapter.
Answer: Always define laws/theories clearly, show workings for Mr calculations with Ar values, and practice balancing equations and counting atoms — show steps for full marks.