Short Q&A for Some Basic Concepts of Chemistry (Class 11 NCERT)

 Over 110 Short Q&A for Some Basic Concepts of Chemistry (Class 11 NCERT)

Part 1: Nature of Matter and Laws of Chemical Combination

Q. No.	Question	Answer
1	What is the scientific study of matter's composition, structure, properties, and reactions?	Chemistry.
2	What are the three common states of matter?	Solid, Liquid, and Gas.
3	How is matter classified at the macroscopic level?	Mixtures and Pure Substances.
4	Define a Pure Substance.	Substance having a fixed composition and single set of properties.
5	What is the difference between an element and a compound?	Element cannot be broken down; Compound can be broken down chemically into elements.
6	Give an example of a Homogeneous Mixture.	Salt solution or Air.
7	Give an example of a Heterogeneous Mixture.	Sand and water or Gunpowder.
8	State the Law of Conservation of Mass.	Mass can neither be created nor destroyed in a chemical reaction.
9	Who proposed the Law of Conservation of Mass?	Antoine Lavoisier.
10	State the Law of Definite Proportions.	A pure chemical compound always contains the same elements combined in a fixed proportion by mass.
11	Who stated the Law of Definite Proportions?	Joseph Proust.
12	State the Law of Multiple Proportions.	If two elements combine to form more than one compound, the masses of one element that combine with a fixed mass of the other are in a simple whole-number ratio.
13	Who proposed the Law of Multiple Proportions?	John Dalton.
14	What is the basis of Gay-Lussac's Law of Gaseous Volumes?	Volumes of reacting gases and products are in a simple whole-number ratio (at constant T and P).
15	State Avogadro's Law.	Equal volumes of all gases (at the same T and P) contain an equal number of molecules.
16	What is the fundamental idea of Dalton's Atomic Theory?	Matter consists of indivisible atoms.
17	Which of Dalton's postulates was proved wrong by later discoveries?	Atoms are indivisible (we now know of subatomic particles).

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Part 2: Concepts of Mass, Moles, and Stoichiometry

Q. No.	Question	Answer
18	What is the SI unit of mass?	Kilogram (kg).
19	What is the SI unit of amount of substance?	Mole (mol).
20	Define Atomic Mass Unit (amu) or u.	1/12th of the mass of a C−12 atom.
21	Define Average Atomic Mass.	The weighted average of the atomic masses of all naturally occurring isotopes of an element.
22	Define Molecular Mass.	The sum of the atomic masses of all atoms in a molecule.
23	Define Formula Mass.	The sum of the atomic masses of ions in a formula unit of an ionic compound (e.g., NaCl).
24	What is the mass of one mole of a substance called?	Molar Mass (g mol−1).
25	Define a Mole.	The amount of substance that contains as many particles as there are atoms in 12 g of C−12 isotope.
26	What is the value of Avogadro's constant (NA​)?	6.022×1023 particles/mol.
27	How many molecules are in 1 mole of water (H2​O)?	6.022×1023 molecules.
28	What is the volume occupied by 1 mole of an ideal gas at STP?	22.7 L (Standard Temperature and Pressure, 273.15 K,1 bar).
29	What is the molar mass of CO2​?	12.01+2(16.00)=44.01 g/mol.
30	Define Percentage Composition.	The percentage of the mass of each element in a compound.
31	Define Empirical Formula.	The simplest whole-number ratio of atoms of different elements present in a compound.
32	Define Molecular Formula.	The actual number of atoms of different elements present in a molecule of a compound.
33	What is the empirical formula of glucose (C6​H12​O6​)?	CH2​O.
34	How are molecular mass and empirical formula mass related?	Molecular Mass =n×Empirical Formula Mass (n is a simple integer).
35	What is Stoichiometry?	The calculation of the mass/moles/volumes of reactants and products in a balanced chemical reaction.
36	What is the Limiting Reagent (or Reactant)?	The reactant that is completely consumed during the reaction and limits the amount of product formed.
37	What is the Excess Reagent (or Reactant)?	The reactant present in an amount greater than required by the limiting reagent.

Part 3: Solutions and Concentration Terms

Q. No.	Question	Answer
38	Define Mass Percentage of a component in a solution.	(Mass of component/Total mass of solution)×100.
39	Define Molarity (M).	Moles of solute dissolved per liter of solution (mol L−1).
40	Write the formula for Molarity.	M=n/V (where n is moles, V is volume in L).
41	Define Molality (m).	Moles of solute dissolved per kilogram of solvent (mol kg−1).
42	Which concentration term is independent of temperature?	Molality (since it involves only masses, which don't change with T).
43	Define Mole Fraction (χ).	The ratio of the number of moles of one component to the total number of moles of all components.
44	What is the sum of the mole fractions of all components in a solution?	Always 1 (χA​+χB​+⋯=1).
45	What is the relationship between Molarity and Volume of solution upon dilution?	M1​V1​=M2​V2​ (Dilution Equation).
46	If you dissolve 10 g of NaOH (molar mass ≈40) in 1 L of water, what is its Molarity?	0.25 M (0.25 moles/1 L).
47	If you dissolve 1 mole of glucose in 1 kg of water, what is its Molality?	1 mol/kg or 1 m.

Part 4: Scientific Notation, Units, and Measurements

Q. No.	Question	Answer
48	What is the standard notation for expressing very large or very small numbers?	Scientific Notation (N×10n).
49	Convert 0.00045 into scientific notation.	4.5×10−4.
50	Convert 602200000000000000000000 into scientific notation.	6.022×1023.
51	What does Precision refer to in measurements?	The closeness of a set of measurements to each other.
52	What does Accuracy refer to in measurements?	The closeness of a single measurement to the true value.
53	What are Significant Figures?	All certain digits plus one uncertain digit.
54	How many significant figures are in 0.0025?	Two (leading zeros are not significant).
55	How many significant figures are in 200.0?	Four (trailing zeros after the decimal point are significant).
56	How many significant figures are in 500?	One (unless specified with a decimal point).
57	What is the result of 4.5+2.15 to the correct number of significant figures?	6.7 (result limited by 4.5 which has one digit after decimal).
58	What are the seven Base SI Units?	metre, kilogram, second, ampere, kelvin, mole, candela.
59	What is the SI unit of temperature?	Kelvin (K).
60	What is the relationship between Celsius (tC​) and Kelvin (TK​)?	TK​=tC​+273.15.
61	What is the SI unit of density?	kg m−3 (or g cm−3).
62	What is the conversion factor between 1 L and cm3?	1 L=1000 cm3 (1 dm3).
63	What method is used to convert units by tracking dimensions?	Dimensional Analysis (or Factor Label Method).

Part 5: Stoichiometry Calculations (Numerical Focus)

Q. No.	Question	Answer
64	What is the formula to calculate the number of moles (n) from mass (m)?	n=m/Molar Mass (MM).
65	How many moles are in 4.0 g of NaOH (MM=40 g/mol)?	0.1 mole.
66	What is the mass of 0.5 mole of H2​O (MM=18 g/mol)?	9 g.
67	What is the volume of 2 moles of O2​ gas at STP?	2×22.7=45.4 L.
68	Calculate the number of atoms in 1 mole of He.	6.022×1023 atoms.
69	Calculate the number of O atoms in 1 molecule of H2​SO4​.	4 atoms.
70	Calculate the number of H atoms in 1 mole of H2​O.	2×NA​=1.2044×1024 atoms.
71	If 1 g of H reacts with O2​ to form H2​O, how much H2​O is formed? (Balanced Eqn:2H2​+O2​→2H2​O).	9 g (H2​ MM=2; 1 g H2​=0.5 mole. 0.5×18 g/mol=9 g).
72	What does 2 L of N2​ reacting with 6 L of H2​ yield (volume basis)? (N2​+3H2​→2NH3​).	4 L of NH3​.
73	If H2​ and O2​ react (2H2​+O2​→2H2​O) and H2​ is the limiting reactant, which reactant is in excess?	Oxygen (O2​).

Part 6: Important Definitions and Miscellaneous Concepts

Q. No.	Question	Answer
74	What is the property of CO2​ being 44.01 g/mol?	Its Molar Mass.
75	What is the freezing point of water in Kelvin?	273.15 K (0∘C).
76	What is the boiling point of water in Kelvin?	373.15 K (100∘C).
77	Is the number of atoms in 12 g of C−12 exactly 6.022×1023?	Yes, by definition of the mole.
78	What is the SI unit of volume?	m3 (cubic meter).
79	Which unit of pressure is equal to 105 Pa?	1 bar.
80	What is the standard temperature for STP in Celsius?	0∘C.
81	What is the term for the number of valence shell electrons?	Valency (for main group elements).
82	How is the mass of a single atom of H calculated?	Molar Mass/NA​ (1.008 g/mol/6.022×1023).
83	What are the two types of uncertainties in measurements?	Random and Systematic errors.
84	What is the term for matter that is uniform in composition?	Homogeneous (can be a solution or a pure substance).
85	What is the term for matter that has components visually separable?	Heterogeneous (a mixture).
86	What did Dalton say about atoms of the same element?	They are identical in all respects (later proved incorrect due to isotopes).
87	What is the percentage of O in H2​O? (MM=18.02 g/mol).	88.81% ((16.00/18.02)×100).
88	What is the relationship between the empirical formula and the simplest ratio of moles in a compound?	They are the same; the EF is derived from the mole ratio.
89	How does Molarity change if the solution volume is doubled by adding water?	Molarity is halved.
90	What is the minimum number of significant figures in the result of any calculation?	It is limited by the term with the fewest significant figures used in the calculation.
91	Define Standard Conditions of Temperature and Pressure (STP).	273.15 K and 1 bar.
92	What is the maximum number of decimal places allowed in the final answer of an addition problem?	It is limited by the term with the fewest decimal places.
93	Name the chemical principle behind balancing a chemical equation.	Law of Conservation of Mass.
94	What is the definition of 1 Pascal (Pa)?	**1 Newton per meter2 (1 N/m2).
95	What is the mass of one electron?	9.109×10−31 kg (negligible in chemical mass calculations).
96	What is the SI unit for luminous intensity?	Candela (cd).
97	What is the SI unit for electric current?	Ampere (A).
98	Give an example of a derived SI unit.	Density (kg/m3), Volume (m3), or Force (Newton).
99	What is the conversion factor for 1 L to m3?	1 L=10−3 m3.
100	What is the difference between 1 bar and 1 atm?	1 atm (≈1.01325 bar) is slightly greater than 1 bar.
101	What is the term for the degree of exactness of a measurement?	Precision.
102	When do the empirical and molecular formulas become the same?	When the simplest whole-number ratio of atoms is also the actual number of atoms (e.g., CH4​).
103	What is the factor 10−6 represented by?	Micro (μ).
104	What is the factor 109 represented by?	Giga (G).
105	Which Law explains why 1 g of H2​ reacts completely with 8 g of O2​ to form 9 g of H2​O?	Law of Definite Proportions (fixed mass ratio 1:8).
106	If 10 g of CaCO3​ produces 4.4 g of CO2​, how much CaO is produced?	10−4.4=5.6 g (Law of Conservation of Mass).
107	Give the scientific notation for the speed of light (300,000,000 m/s).	3.0×108 m/s.
108	Is Molarity preferred over Molality in technical applications?	Yes, Molarity is easier to measure (volume-based).
109	What is the percentage purity of a sample that contains 8 g of pure NaCl in a 10 g total sample?	80%.
110	What is the mass of 1 molecule of O2​?	32/6.022×1023 grams (Molar Mass divided by NA​).
111	What is the number of significant figures in the constant 2 in the formula Circumference=2πr?	Infinite (it is an exact number).
112	What is the density of water at 4∘C?	1.0 g/cm3 (or 1000 kg/m3).