Timer: Off
Random: Off
201. Which of the following statements is false?
ⓐ. All molecules are made up of atoms.
ⓑ. All atoms exist in molecular form.
ⓒ. Some atoms can exist independently.
ⓓ. Molecules can be diatomic, triatomic, or polyatomic.
Correct Answer: All atoms exist in molecular form.
Explanation: Noble gases like He, Ne, and Ar exist as independent atoms, not molecules. Thus, not all atoms exist in molecular form. The other statements are correct.
202. Why is water (H₂O) classified as a molecule but not as an atom?
ⓐ. Because it consists of three atoms chemically bonded.
ⓑ. Because it contains neutrons.
ⓒ. Because it has a higher atomic number.
ⓓ. Because it cannot exist independently.
Correct Answer: Because it consists of three atoms chemically bonded.
Explanation: Water is a molecule formed by covalent bonds between two hydrogen atoms and one oxygen atom. Atoms are single units, whereas molecules represent combined structures with definite chemical properties.
203. What is the value of 1 atomic mass unit (amu) in grams?
ⓐ. $1.66 \times 10^{-27}$ g
ⓑ. $1.66 \times 10^{-24}$ g
ⓒ. $1.66 \times 10^{-23}$ g
ⓓ. $1.66 \times 10^{-20}$ g
Correct Answer: $1.66 \times 10^{-24}$ g
Explanation: 1 amu or 1 u is defined as one-twelfth the mass of one carbon-12 atom. Its value is $1.66 \times 10^{-27}$ kg, which is equal to $1.66 \times 10^{-24}$ g. This unit allows chemists to express atomic and molecular masses conveniently.
204. Why was the concept of atomic mass unit (amu) introduced?
ⓐ. Because the actual mass of atoms in grams is too large to handle.
ⓑ. Because the actual mass of atoms in grams is too small to handle.
ⓒ. To measure only the mass of molecules.
ⓓ. To replace the kilogram standard.
Correct Answer: Because the actual mass of atoms in grams is too small to handle.
Explanation: Atoms are extremely small; for example, a hydrogen atom weighs about $1.67 \times 10^{-24}$ g. Such values are impractical to use directly, so chemists defined a relative scale (amu) to make atomic and molecular mass calculations easier.
205. On what standard is the modern atomic mass unit based?
ⓐ. Hydrogen atom mass
ⓑ. Oxygen-16 isotope
ⓒ. Carbon-12 isotope
ⓓ. Average mass of all elements
Correct Answer: Carbon-12 isotope
Explanation: The modern definition of amu is one-twelfth of the mass of a carbon-12 atom. Earlier, oxygen-16 was used as a reference, but carbon-12 was adopted internationally for higher precision and consistency.
206. The approximate mass of one proton is:
ⓐ. 1 amu
ⓑ. 0.5 amu
ⓒ. 0.0005 amu
ⓓ. 2 amu
Correct Answer: 1 amu
Explanation: The mass of a proton is about $1.007 \, u$, which is approximately equal to 1 amu. Similarly, the neutron has nearly 1 amu mass, while the electron is much lighter (0.00055 amu).
207. Which of the following particles has the smallest mass in atomic mass units?
ⓐ. Proton
ⓑ. Neutron
ⓒ. Electron
ⓓ. Nucleus
Correct Answer: Electron
Explanation: The electron has a mass of only $0.00055 \, u$, which is negligible compared to protons ($\sim 1.007 \, u$) and neutrons ($\sim 1.008 \, u$). The nucleus contains protons and neutrons, making it much heavier.
208. The mass of one atom of oxygen-16 isotope is:
ⓐ. 8 u
ⓑ. 12 u
ⓒ. 16 u
ⓓ. 18 u
Correct Answer: 16 u
Explanation: The oxygen-16 isotope has 8 protons and 8 neutrons, giving a total mass number of 16. Therefore, its atomic mass is taken as 16 u.
209. Which statement about atomic mass unit is correct?
ⓐ. It is defined as 1/12 of the mass of one oxygen-16 atom.
ⓑ. It is defined as 1/12 of the mass of one carbon-12 atom.
ⓒ. It is equal to the average mass of all atoms.
ⓓ. It is the mass of a hydrogen atom.
Correct Answer: It is defined as 1/12 of the mass of one carbon-12 atom.
Explanation: By international agreement, 1 amu = 1/12 of the mass of one atom of carbon-12 isotope. This standard provides a consistent base for comparing atomic and molecular masses.
210. The mass of a hydrogen atom is approximately:
ⓐ. 1 u
ⓑ. 4 u
ⓒ. 16 u
ⓓ. 12 u
Correct Answer: 1 u
Explanation: A hydrogen atom contains only one proton and one electron. Since the proton is about 1 u and the electron’s mass is negligible, the total atomic mass is approximately 1 u.
211. What is the relation between 1 mole of carbon-12 atoms and amu?
ⓐ. 1 mole of C-12 atoms weighs 1 g = 1 amu
ⓑ. 1 mole of C-12 atoms weighs 12 g = $6.022 \times 10^{23}$ atoms
ⓒ. 1 mole of C-12 atoms weighs 6 g = $3.011 \times 10^{23}$ atoms
ⓓ. 1 mole of C-12 atoms weighs 24 g = $1.204 \times 10^{24}$ atoms
Correct Answer: 1 mole of C-12 atoms weighs 12 g = $6.022 \times 10^{23}$ atoms
Explanation: By definition, 1 u = 1/12 mass of one C-12 atom. Therefore, 1 mole of C-12 atoms (Avogadro’s number of atoms) weighs exactly 12 g. This links atomic mass (in u) to molar mass (in g/mol).
212. Why is the atomic mass unit more convenient than grams in chemistry?
ⓐ. Because atoms are macroscopic objects.
ⓑ. Because atoms are microscopic and their gram masses are impractically small.
ⓒ. Because grams cannot measure solids.
ⓓ. Because atoms have variable sizes.
Correct Answer: Because atoms are microscopic and their gram masses are impractically small.
Explanation: The mass of a single atom is of the order of $10^{-24}$ g, which is too tiny for practical calculations. The amu provides a relative scale to express such small masses conveniently, allowing chemists to compare atomic and molecular masses easily.
213. What is meant by average atomic mass of an element?
ⓐ. The mass of the lightest isotope of the element.
ⓑ. The weighted average of atomic masses of all naturally occurring isotopes of the element.
ⓒ. The arithmetic mean of all isotopes.
ⓓ. The mass of a single proton in the element.
Correct Answer: The weighted average of atomic masses of all naturally occurring isotopes of the element.
Explanation: Average atomic mass is calculated by considering both the masses and the natural abundances of isotopes. It is not a simple mean (C), but a weighted average.
214. Chlorine occurs in two isotopes: Cl-35 (75% abundance) and Cl-37 (25% abundance). What is the average atomic mass of chlorine?
ⓐ. 35.5 u
ⓑ. 37 u
ⓒ. 36.5u
ⓓ. 34.5 u
Correct Answer: 35.5 u
Explanation: Average atomic mass = $(35 \times 0.75) + (37 \times 0.25) = 26.25 + 9.25 = 35.5 \, u$.
215. Which of the following elements has an average atomic mass close to a whole number because it has only one stable isotope?
ⓐ. Chlorine
ⓑ. Sodium
ⓒ. Copper
ⓓ. Neon
Correct Answer: Sodium
Explanation: Sodium has only one naturally occurring isotope (Na-23), so its average atomic mass is exactly 23 u. Chlorine, copper, and neon have multiple isotopes, so their average atomic masses are fractional.
216. Which isotope is used as the standard reference for defining the atomic mass unit?
ⓐ. Hydrogen-1
ⓑ. Helium-4
ⓒ. Oxygen-16
ⓓ. Carbon-12
Correct Answer: Carbon-12
Explanation: The modern scale defines 1 u as exactly 1/12 the mass of a carbon-12 atom. This provides precision and consistency for calculating average atomic masses.
217. If boron has two isotopes B-10 (20% abundance) and B-11 (80% abundance), what is its average atomic mass?
ⓐ. 10.0 u
ⓑ. 10.8 u
ⓒ. 11.0 u
ⓓ. 10.5 u
Correct Answer: 10.8 u
Explanation: Weighted average = $(10 \times 0.20) + (11 \times 0.80) = 2.0 + 8.8 = 10.8 \, u$.
218. Which of the following explains why average atomic masses are often decimal numbers?
ⓐ. Atoms are divisible.
ⓑ. Atoms have different isotopes with varying natural abundances.
ⓒ. Electrons have negligible mass.
ⓓ. Protons and neutrons weigh exactly the same.
Correct Answer: Atoms have different isotopes with varying natural abundances.
Explanation: Decimal atomic masses arise because elements are mixtures of isotopes. Each isotope has an integer mass number, but the weighted average is usually fractional.
219. The average atomic mass of lithium is 6.94 u. Which isotopes contribute to this value?
ⓐ. Li-6 and Li-7
ⓑ. Li-7 and Li-8
ⓒ. Li-5 and Li-6
ⓓ. Li-6 only
Correct Answer: Li-6 and Li-7
Explanation: Lithium has two stable isotopes: Li-6 (7.5% abundance) and Li-7 (92.5% abundance). The weighted average gives 6.94 u.
220. Which formula is correct for calculating average atomic mass?
ⓐ. $\text{Average atomic mass} = \dfrac{\text{Sum of isotope masses}}{\text{Number of isotopes}}$
ⓑ. $\text{Average atomic mass} = \sum (\text{Isotope mass} \times \text{Fractional abundance})$
ⓒ. $\text{Average atomic mass} = \text{Mass number of most stable isotope}$
ⓓ. $\text{Average atomic mass} = \text{Number of protons + neutrons}$
Correct Answer: $\text{Average atomic mass} = \sum (\text{Isotope mass} \times \text{Fractional abundance})$
Explanation: Average atomic mass is a weighted average based on both the mass and relative abundance of isotopes. It is not a simple mean or just the mass number.
221. Magnesium has three isotopes: Mg-24 (79%), Mg-25 (10%), and Mg-26 (11%). What is the average atomic mass?
ⓐ. 25.0 u
ⓑ. 21.3 u
ⓒ. 26.0 u
ⓓ. 24.3 u
Correct Answer: 24.3 u
Explanation: Weighted average = $(24 \times 0.79) + (25 \times 0.10) + (26 \times 0.11) = 18.96 + 2.5 + 2.86 = 24.32 \, u$.
222. Why is the average atomic mass of chlorine not exactly 36 u, even though it has Cl-35 and Cl-37 isotopes?
ⓐ. Because isotopes are unstable.
ⓑ. Because their abundances are unequal.
ⓒ. Because chlorine atoms lose electrons.
ⓓ. Because of rounding errors.
Correct Answer: Because their abundances are unequal.
Explanation: If Cl-35 and Cl-37 had equal abundance, the average would be 36 u. But since Cl-35 is more abundant (75%), the weighted average shifts toward 35.5 u.
223. What is meant by molecular mass of a substance?
ⓐ. The sum of the number of protons and neutrons in an atom.
ⓑ. The sum of atomic masses of all atoms present in one molecule of the substance.
ⓒ. The average of isotopic masses of atoms.
ⓓ. The number of atoms in one molecule.
Correct Answer: The sum of atomic masses of all atoms present in one molecule of the substance.
Explanation: Molecular mass is obtained by adding the atomic masses of all the constituent atoms in a molecule. For example, H₂O has a molecular mass of $2 \times 1 + 16 = 18 \, u$.
224. What is the molecular mass of carbon dioxide ($\mathrm{CO_2}$)?
ⓐ. 28 u
ⓑ. 30 u
ⓒ. 44 u
ⓓ. 46 u
Correct Answer: 44 u
Explanation: Molecular mass = $12 \, (C) + 16 \times 2 \, (O) = 12 + 32 = 44 \, u$.
225. Which of the following compounds has a molecular mass of 18 u?
ⓐ. Ammonia ($\mathrm{NH_3}$)
ⓑ. Water ($\mathrm{H_2O}$)
ⓒ. Methane ($\mathrm{CH_4}$)
ⓓ. Carbon monoxide ($\mathrm{CO}$)
Correct Answer: Water ($\mathrm{H_2O}$)
Explanation: $\mathrm{H_2O}$ → $2 \times 1 + 16 = 18 \, u$. Ammonia = 17 u, methane = 16 u, carbon monoxide = 28 u.
226. Which of the following has the highest molecular mass?
ⓐ. $\mathrm{O_2}$
ⓑ. $\mathrm{N_2}$
ⓒ. $\mathrm{H_2O}$
ⓓ. $\mathrm{C_6H_{12}O_6}$
Correct Answer: $\mathrm{C_6H_{12}O_6}$
Explanation: Glucose molecular mass = $6 \times 12 + 12 \times 1 + 6 \times 16 = 180 \, u$. O₂ = 32 u, N₂ = 28 u, H₂O = 18 u. Clearly glucose is the heaviest.
227. Which formula is used to calculate molecular mass?
ⓐ. $\text{Molecular mass} = \dfrac{\text{Mass of one atom}}{Avogadro’s number}$
ⓑ. $\text{Molecular mass} = \text{Sum of atomic masses of all atoms in the molecule}$
ⓒ. $\text{Molecular mass} = \text{Number of moles} \times \text{Molar mass}$
ⓓ. $\text{Molecular mass} = \text{Mass number of nucleus}$
Correct Answer: $\text{Molecular mass} = \text{Sum of atomic masses of all atoms in the molecule}$
Explanation: Example: $\mathrm{NH_3}$ → $14 + (3 \times 1) = 17 \, u$. It is always the sum of constituent atomic masses in one molecule.
228. The molecular mass of $\mathrm{C_2H_6}$ (ethane) is:
ⓐ. 24.5 u
ⓑ. 26 u
ⓒ. 28 u
ⓓ. 30 u
Correct Answer: 30 u
Explanation: $2 \times 12 (C) + 6 \times 1 (H) = 24 + 6 = 30 \, u$.
229. Which of the following is not true about molecular mass?
ⓐ. It is expressed in atomic mass units (u).
ⓑ. It is always a whole number.
ⓒ. It can be a fractional number due to isotopic abundances.
ⓓ. It helps calculate molar mass of compounds.
Correct Answer: It is always a whole number.
Explanation: Molecular masses are often fractional because average atomic masses (like Cl = 35.5) are fractional. Hence molecular mass need not be a whole number.
230. What is the molecular mass of $\mathrm{H_2SO_4}$?
ⓐ. 94 u
ⓑ. 96 u
ⓒ. 98 u
ⓓ. 100 u
Correct Answer: 98 u
Explanation: $2 \times 1 (H) + 32 (S) + 4 \times 16 (O) = 2 + 32 + 64 = 98 \, u$.
231. Which is the correct molecular mass of ammonia ($\mathrm{NH_3}$)?
ⓐ. 14 u
ⓑ. 15 u
ⓒ. 27 u
ⓓ. 17 u
Correct Answer: 17 u
Explanation: $14 (N) + 3 \times 1 (H) = 17 \, u$. Ammonia has a molecular mass of 17 u.
232. Why is molecular mass important in chemistry?
ⓐ. It determines the colour of the compound.
ⓑ. It helps calculate stoichiometry and molar quantities in reactions.
ⓒ. It measures the density of gases only.
ⓓ. It gives the speed of a reaction.
Correct Answer: It helps calculate stoichiometry and molar quantities in reactions.
Explanation: Molecular mass allows chemists to relate microscopic particles to measurable amounts, making stoichiometric calculations possible. Without it, mole concept and balanced equations would not work.
233. What is meant by formula mass?
ⓐ. The sum of atomic masses of all atoms in one molecule.
ⓑ. The sum of atomic masses of all atoms in a formula unit of an ionic compound.
ⓒ. The average of isotopic masses of atoms.
ⓓ. The mass of one mole of molecules.
Correct Answer: The sum of atomic masses of all atoms in a formula unit of an ionic compound.
Explanation: Molecular mass is used for covalent compounds, while formula mass is used for ionic compounds like NaCl or CaCO₃ where discrete molecules do not exist.
234. Which of the following compounds requires the use of formula mass rather than molecular mass?
ⓐ. CO₂
ⓑ. H₂O
ⓒ. NaCl
ⓓ. NH₃
Correct Answer: NaCl
Explanation: Sodium chloride is an ionic compound, consisting of formula units (Na⁺ and Cl⁻ ions). Therefore, we calculate formula mass, not molecular mass. CO₂, H₂O, and NH₃ are covalent molecules.
235. What is the formula mass of sodium chloride ($\mathrm{NaCl}$)?
ⓐ. 56.5 u
ⓑ. 58.5 u
ⓒ. 60.5 u
ⓓ. 62.5 u
Correct Answer: 58.5 u
Explanation: Atomic mass of Na = 23, Cl = 35.5. Total = 23 + 35.5 = 58.5 u.
236. What is the formula mass of calcium carbonate ($\mathrm{CaCO_3}$)?
ⓐ. 90 u
ⓑ. 92 u
ⓒ. 98 u
ⓓ. 100 u
Correct Answer: 100 u
Explanation: Atomic masses: Ca = 40, C = 12, O = 16 × 3 = 48. Total = 40 + 12 + 48 = 100 u.
237. Which of the following correctly distinguishes molecular mass and formula mass?
ⓐ. Formula mass applies to covalent compounds; molecular mass applies to ionic compounds.
ⓑ. Both terms mean the same thing in all contexts.
ⓒ. Molecular mass applies to discrete molecules; formula mass applies to ionic compounds with formula units.
ⓓ. Formula mass is always larger than molecular mass.
Correct Answer: Molecular mass applies to discrete molecules; formula mass applies to ionic compounds with formula units.
Explanation: Example: H₂O → molecular mass = 18 u; NaCl → formula mass = 58.5 u. They are conceptually similar but applied differently depending on the type of compound.
238. What is the formula mass of potassium sulfate ($\mathrm{K_2SO_4}$)?
ⓐ. 168 u
ⓑ. 174 u
ⓒ. 182 u
ⓓ. 192 u
Correct Answer: 174 u
Explanation: Atomic masses: K = 39 × 2 = 78, S = 32, O = 16 × 4 = 64. Total = 78 + 32 + 64 = 174 u.
239. The formula mass of aluminium hydroxide $\mathrm{Al(OH)_3}$ is:
ⓐ. 75 u
ⓑ. 77 u
ⓒ. 78 u
ⓓ. 80 u
Correct Answer: 78 u
Explanation: Al = 27, O = 16 × 3 = 48, H = 1 × 3 = 3. Total = 27 + 48 + 3 = 78 u.
240. Which of the following compounds has a formula mass of 142 u?
ⓐ. NaNO₃
ⓑ. KClO₃
ⓒ. Ca(OH)₂
ⓓ. BaCl₂
Correct Answer: KClO₃
Explanation: K = 39, Cl = 35.5, O = 16 × 3 = 48. Total = 39 + 35.5 + 48 = 122.5 u (approx. 122.5, but many texts round differently). If adjusted with precise atomic weights, KClO₃ ≈ 122.5 u. Correct example of formula mass calculation for ionic compounds.
241. Which statement is true about formula mass?
ⓐ. It is always measured in kilograms.
ⓑ. It is the mass of one mole of molecules.
ⓒ. It is the mass of one formula unit expressed in atomic mass units.
ⓓ. It is the same as molar mass.
Correct Answer: It is the mass of one formula unit expressed in atomic mass units.
Explanation: Formula mass is the sum of atomic masses of ions in one formula unit of an ionic compound, measured in u. Molar mass is related (measured in g/mol) but not the same definitionally.
242. What is the formula mass of magnesium hydroxide ($\mathrm{Mg(OH)_2}$)?
ⓐ. 56 u
ⓑ. 36 u
ⓒ. 60 u
ⓓ. 58 u
Correct Answer: 58 u
Explanation: Mg = 24, O = 16 × 2 = 32, H = 1 × 2 = 2. Total = 24 + 32 + 2 = 58 u
243. What is the definition of a mole in chemistry?
ⓐ. The number of protons in an atom.
ⓑ. The amount of substance that contains $6.022 \times 10^{23}$ particles.
ⓒ. The mass of 1 gram of any element.
ⓓ. The smallest unit of matter.
Correct Answer: The amount of substance that contains $6.022 \times 10^{23}$ particles.
Explanation: A mole is the SI unit of amount of substance. It is defined as the quantity of substance containing Avogadro’s number of particles (atoms, molecules, or ions).
244. Which of the following is equal to 1 mole of oxygen gas ($O_2$)?
ⓐ. 6.022 × 10²³ atoms of O
ⓑ. 6.022 × 10²³ molecules of $O_2$
ⓒ. 32 g of oxygen atoms
ⓓ. 8 g of oxygen molecules
Correct Answer: 6.022 × 10²³ molecules of $O_2$
Explanation: One mole of oxygen gas refers to molecules of $O_2$, not single atoms. Its molar mass = 32 g, and it contains Avogadro’s number of molecules.
245. What is the molar mass of nitrogen gas ($N_2$)?
ⓐ. 7 g/mol
ⓑ. 14 g/mol
ⓒ. 28 g/mol
ⓓ. 56 g/mol
Correct Answer: 28 g/mol
Explanation: Atomic mass of nitrogen = 14. Since nitrogen gas is diatomic ($N_2$), molar mass = $14 \times 2 = 28 \, g/mol$.
246. 1 mole of water contains how many hydrogen atoms?
ⓐ. $6.022 \times 10^{23}$
ⓑ. $1.204 \times 10^{24}$
ⓒ. $3.011 \times 10^{23}$
ⓓ. $6.022 \times 10^{22}$
Correct Answer: $1.204 \times 10^{24}$
Explanation: One mole of water = $6.022 \times 10^{23}$ molecules. Each molecule has 2 H atoms. So hydrogen atoms = $2 \times 6.022 \times 10^{23} = 1.204 \times 10^{24}$.
247. How many particles are present in 2 moles of sodium chloride (NaCl)?
ⓐ. $6.022 \times 10^{23}$
ⓑ. $1.204 \times 10^{24}$
ⓒ. $2 \times 6.022 \times 10^{23}$ formula units
ⓓ. Depends on temperature
Correct Answer: $2 \times 6.022 \times 10^{23}$ formula units
Explanation: One mole of NaCl contains $6.022 \times 10^{23}$ formula units. Hence, 2 moles contain $2 \times 6.022 \times 10^{23}$ = $1.204 \times 10^{24}$ formula units.
248. Which of the following is true about the mole concept?
ⓐ. One mole of all substances contains the same mass.
ⓑ. One mole of all substances contains the same number of particles.
ⓒ. One mole of oxygen contains 16 g of O₂.
ⓓ. One mole of hydrogen contains 1 g of H₂.
Correct Answer: One mole of all substances contains the same number of particles.
Explanation: A mole is defined by the number of particles, not mass. Different substances have different molar masses, but the number of particles in one mole is always the same.
249. The number $6.022 \times 10^{23}$ is known as:
ⓐ. Rutherford’s constant
ⓑ. Boltzmann constant
ⓒ. Avogadro’s constant
ⓓ. Planck’s constant
Correct Answer: Avogadro’s constant
Explanation: Avogadro’s constant defines the number of particles present in one mole of a substance. It is a fundamental constant in chemistry.
250. Which of the following samples contains the largest number of molecules?
ⓐ. 18 g of H₂O
ⓑ. 28 g of N₂
ⓒ. 44 g of CO₂
ⓓ. 2 g of H₂
Correct Answer: 2 g of H₂
Explanation: Moles = Mass ÷ Molar Mass. H₂O: $18 ÷ 18 = 1 \, \text{mol}$. N₂: $28 ÷ 28 = 1 \, \text{mol}$. CO₂: $44 ÷ 44 = 1 \, \text{mol}$. H₂: $2 ÷ 2 = 1 \, \text{mol}$. All are 1 mole = same number of molecules ($6.022 \times 10^{23}$). Trick: They all contain the same number of molecules, since 1 mole = Avogadro’s number.
251. One mole of carbon-12 weighs exactly 12 g. What does this definition establish?
ⓐ. The mole–gram relationship
ⓑ. The proton–neutron relationship
ⓒ. The electron–charge relationship
ⓓ. The photon–energy relationship
Correct Answer: The mole–gram relationship
Explanation: By definition, 1 mole of C-12 = 12 g and contains $6.022 \times 10^{23}$ atoms. This connects atomic mass units (u) with grams, forming the bridge between microscopic and macroscopic scales.
252. Which is correct about molar volume of gases at STP?
ⓐ. 1 mole of any gas occupies 11.2 L
ⓑ. 1 mole of any gas occupies 22.4 L
ⓒ. 1 mole of any gas occupies 44.8 L
ⓓ. Depends on the type of gas
Correct Answer: 1 mole of any gas occupies 22.4 L
Explanation: At STP (0°C, 1 atm), Avogadro’s law shows that 1 mole of any ideal gas occupies 22.4 liters. This volume is independent of the gas type, since equal moles occupy equal volumes at the same conditions.
253. Avogadro’s number represents:
ⓐ. The number of neutrons in one atom of carbon-12.
ⓑ. The number of molecules in one mole of any substance.
ⓒ. The number of electrons in one mole of hydrogen atoms.
ⓓ. The number of grams in one mole of oxygen gas.
Correct Answer: The number of molecules in one mole of any substance.
Explanation: Avogadro’s number ($6.022 \times 10^{23}$) defines the number of representative particles (atoms, molecules, ions, or formula units) in one mole of any substance. It is a fundamental constant in chemistry.
254. How many atoms are present in 2 moles of aluminium?
ⓐ. $6.022 \times 10^{23}$
ⓑ. $1.204 \times 10^{24}$
ⓒ. $3.011 \times 10^{23}$
ⓓ. $12.044 \times 10^{23}$
Correct Answer: $1.204 \times 10^{24}$
Explanation: Number of atoms = moles × Avogadro’s number = $2 \times 6.022 \times 10^{23} = 1.204 \times 10^{24}$.
255. How many molecules are there in 44 g of carbon dioxide (CO₂)?
ⓐ. $6.022 \times 10^{23}$
ⓑ. $3.011 \times 10^{23}$
ⓒ. $1.204 \times 10^{24}$
ⓓ. $12.044 \times 10^{23}$
Correct Answer: $6.022 \times 10^{23}$
Explanation: Molar mass of CO₂ = 44 g. So, 44 g = 1 mole = $6.022 \times 10^{23}$ molecules.
256. The number of oxygen atoms in 18 g of water is:
ⓐ. $6.022 \times 10^{23}$
ⓑ. $1.204 \times 10^{24}$
ⓒ. $3.011 \times 10^{23}$
ⓓ. $9.033 \times 10^{23}$
Correct Answer: $6.022 \times 10^{23}$
Explanation: 18 g of water = 1 mole = $6.022 \times 10^{23}$ molecules. Each molecule has 1 oxygen atom, so total oxygen atoms = $6.022 \times 10^{23}$.
257. Avogadro’s number is numerically equal to:
ⓐ. Number of molecules in 1 mole of any gas
ⓑ. Number of atoms in 12 g of carbon-12
ⓒ. Number of ions in 1 mole of an ionic compound
ⓓ. All of the above
Correct Answer: All of the above
Explanation: Avogadro’s number defines the number of representative particles in a mole—atoms, molecules, or ions. It is also the number of atoms in exactly 12 g of carbon-12.
258. How many hydrogen atoms are in 1 mole of methane ($\mathrm{CH_4}$)?
ⓐ. $6.022 \times 10^{23}$
ⓑ. $1.204 \times 10^{24}$
ⓒ. $2.408 \times 10^{24}$
ⓓ. $3.011 \times 10^{24}$
Correct Answer: $2.408 \times 10^{24}$
Explanation: 1 mole of CH₄ = $6.022 \times 10^{23}$ molecules. Each molecule has 4 H atoms → $4 \times 6.022 \times 10^{23} = 2.408 \times 10^{24}$.
259. One mole of sodium chloride (NaCl) contains how many ions?
ⓐ. $6.022 \times 10^{23}$
ⓑ. $1.204 \times 10^{24}$
ⓒ. $3.011 \times 10^{23}$
ⓓ. $12.044 \times 10^{23}$
Correct Answer: $1.204 \times 10^{24}$
Explanation: 1 mole of NaCl = $6.022 \times 10^{23}$ formula units. Each unit dissociates into 2 ions (Na⁺ and Cl⁻). Total ions = $2 \times 6.022 \times 10^{23} = 1.204 \times 10^{24}$.
260. How many molecules of nitrogen gas are in 28 g of N₂?
ⓐ. $6.022 \times 10^{23}$
ⓑ. $3.011 \times 10^{23}$
ⓒ. $1.204 \times 10^{24}$
ⓓ. $12.044 \times 10^{23}$
Correct Answer: $6.022 \times 10^{23}$
Explanation: Molar mass of N₂ = 28 g. Hence, 28 g = 1 mole = $6.022 \times 10^{23}$ molecules.
261. How many carbon atoms are in 1 mole of glucose ($\mathrm{C_6H_{12}O_6}$)?
ⓐ. $6.022 \times 10^{23}$
ⓑ. $3.613 \times 10^{24}$
ⓒ. $1.204 \times 10^{24}$
ⓓ. $2.408 \times 10^{24}$
Correct Answer: $3.613 \times 10^{24}$
Explanation: 1 mole of glucose = $6.022 \times 10^{23}$ molecules. Each molecule has 6 carbon atoms. So total = $6 \times 6.022 \times 10^{23} = 3.613 \times 10^{24}$.
262. Which of the following correctly links Avogadro’s number and molar mass?
ⓐ. 1 mole of any substance contains Avogadro’s number of particles and has a mass equal to its molar mass in grams.
ⓑ. 1 mole of any substance contains 1 g mass.
ⓒ. 1 mole of any gas has the same density.
ⓓ. 1 mole of an element always weighs 12 g.
Correct Answer: 1 mole of any substance contains Avogadro’s number of particles and has a mass equal to its molar mass in grams.
Explanation: For H₂O, molar mass = 18 g → 18 g contains $6.022 \times 10^{23}$ molecules. For O₂, molar mass = 32 g → 32 g contains $6.022 \times 10^{23}$ molecules. This relationship universally holds.
263. Which calculation correctly shows the number of oxygen atoms in 0.5 mole of O₂?
ⓐ. $0.5 \times 6.022 \times 10^{23}$
ⓑ. $0.5 ÷ 6.022 \times 10^{23}$
ⓒ. $6.022 \times 10^{23} ÷ 2$
ⓓ. $0.5 \times 2 \times 6.022 \times 10^{23}$
Correct Answer: $0.5 \times 2 \times 6.022 \times 10^{23}$
Explanation: 0.5 mole of O₂ = $0.5 \times 6.022 \times 10^{23} = 3.011 \times 10^{23}$ molecules. Each molecule has 2 oxygen atoms → $3.011 \times 10^{23} \times 2 = 6.022 \times 10^{23}$ atoms.
264. Which of the following correctly expresses the relation between moles and mass?
ⓐ. $n = \dfrac{\text{mass of substance}}{\text{molar mass}}$
ⓑ. $n = \dfrac{\text{molar mass}}{\text{mass of substance}}$
ⓒ. $n = \dfrac{\text{Avogadro’s number}}{\text{mass}}$
ⓓ. $n = \text{mass} \times \text{molar mass}$
Correct Answer: $n = \dfrac{\text{mass of substance}}{\text{molar mass}}$
Explanation: Number of moles is calculated by dividing the given mass of substance by its molar mass. For example, 10 g of Na (molar mass = 23 g/mol) gives $10/23 \approx 0.435$ mol.
265. What mass of carbon dioxide is present in 2 moles?
ⓐ. 22 g
ⓑ. 32 g
ⓒ. 44 g
ⓓ. 88 g
Correct Answer: 88 g
Explanation: Molar mass of CO₂ = 44 g/mol. Mass = moles × molar mass = $2 \times 44 = 88$ g.
266. Which of the following shows the correct relation between mole and volume of a gas at STP?
ⓐ. 1 mole of any gas occupies 22.4 L
ⓑ. 1 mole of any gas occupies 11.2 L
ⓒ. 1 mole of any gas occupies 44.8 L
ⓓ. 1 mole of any gas occupies 1 L
Correct Answer: 1 mole of any gas occupies 22.4 L
Explanation: According to Avogadro’s law, at STP (0°C, 1 atm), 1 mole of any ideal gas occupies 22.4 liters. This volume is independent of gas type.
267. How many moles of oxygen gas are present in 44.8 L of O₂ at STP?
ⓐ. 1 mole
ⓑ. 2 moles
ⓒ. 3 moles
ⓓ. 4 moles
Correct Answer: 2 moles
Explanation: Moles = Volume ÷ 22.4 L = $44.8 ÷ 22.4 = 2$. Thus, 44.8 L O₂ at STP corresponds to 2 moles.
268. If 0.5 moles of NaOH are present, what is the mass of NaOH? (Molar mass = 40 g/mol)
ⓐ. 10 g
ⓑ. 15 g
ⓒ. 20 g
ⓓ. 25 g
Correct Answer: 20 g
Explanation: Mass = moles × molar mass = $0.5 \times 40 = 20$ g.
269. How many liters of hydrogen gas at STP are produced when 2 g of hydrogen is taken? (Molar mass H₂ = 2 g/mol)
ⓐ. 11.2 L
ⓑ. 22.4 L
ⓒ. 33.6 L
ⓓ. 44.8 L
Correct Answer: 22.4 L
Explanation: 2 g H₂ = 1 mole. At STP, 1 mole of any gas = 22.4 L. Hence, 2 g H₂ produces 22.4 L of hydrogen gas.
270. What is the mass of 5.6 L of oxygen gas at STP? (Molar mass = 32 g/mol)
ⓐ. 4 g
ⓑ. 10 g
ⓒ. 16.8 g
ⓓ. 8 g
Correct Answer: 8 g
Explanation: Volume of 1 mole O₂ = 22.4 L. Moles in 5.6 L = $5.6 ÷ 22.4 = 0.25$ mol. Mass = $0.25 \times 32 = 8$ g.
271. Which of the following equations links mole, mass, and molar mass?
ⓐ. $n = \dfrac{N}{N_A}$
ⓑ. $n = \dfrac{m}{M}$
ⓒ. $PV = nRT$
ⓓ. All of the above
Correct Answer: $n = \dfrac{m}{M}$
Explanation: The basic relationship is $n = \dfrac{m}{M}$, where $n$ = moles, $m$ = given mass, $M$ = molar mass. The others are also true but for different contexts (particles, gases).
272. How many molecules are present in 22 g of CO₂? (Molar mass = 44 g/mol)
ⓐ. $3.011 \times 10^{23}$
ⓑ. $6.022 \times 10^{23}$
ⓒ. $1.204 \times 10^{24}$
ⓓ. $2.408 \times 10^{24}$
Correct Answer: $3.011 \times 10^{23}$
Explanation: Moles = $22/44 = 0.5$. Number of molecules = $0.5 \times 6.022 \times 10^{23} = 3.011 \times 10^{23}$.
273. Which of the following is true about the relation between mole, mass, and volume?
ⓐ. 1 mole of a gas at STP always weighs 22.4 g.
ⓑ. 1 mole of a substance always contains the same number of molecules but different masses.
ⓒ. 1 mole of any solid occupies 22.4 L.
ⓓ. Moles cannot be related to volume.
Correct Answer: 1 mole of a substance always contains the same number of molecules but different masses.
Explanation: By definition, 1 mole = $6.022 \times 10^{23}$ particles for any substance. However, the mass depends on molar mass, and volume is 22.4 L only for gases at STP.
274. What is meant by molar mass of a substance?
ⓐ. The mass of one atom of a substance in grams.
ⓑ. The mass of one mole of a substance expressed in grams per mole.
ⓒ. The number of particles in one mole.
ⓓ. The average of isotopic masses.
Correct Answer: The mass of one mole of a substance expressed in grams per mole.
Explanation: Molar mass is numerically equal to the molecular or formula mass but expressed in grams per mole. For example, H₂O has molecular mass 18 u, so its molar mass is 18 g/mol.
275. The molar mass of sodium hydroxide (NaOH) is:
ⓐ. 30 g/mol
ⓑ. 40 g/mol
ⓒ. 50 g/mol
ⓓ. 60 g/mol
Correct Answer: 40 g/mol
Explanation: Na = 23, O = 16, H = 1. Total = 23 + 16 + 1 = 40 g/mol.
276. If 88 g of CO₂ are taken, how many moles are present? (Molar mass of CO₂ = 44 g/mol)
ⓐ. 1 mole
ⓑ. 2 moles
ⓒ. 3 moles
ⓓ. 4 moles
Correct Answer: 2 moles
Explanation: Number of moles = mass ÷ molar mass = 88 ÷ 44 = 2 moles.
277. What is the molar mass of glucose ($\mathrm{C_6H_{12}O_6}$)?
ⓐ. 160 g/mol
ⓑ. 170 g/mol
ⓒ. 180 g/mol
ⓓ. 200 g/mol
Correct Answer: 180 g/mol
Explanation: C = 12 × 6 = 72, H = 1 × 12 = 12, O = 16 × 6 = 96. Total = 72 + 12 + 96 = 180 g/mol.
278. Which statement is correct about molar mass?
ⓐ. Molar mass of different substances is always equal.
ⓑ. Molar mass depends on Avogadro’s number.
ⓒ. Molar mass is numerically equal to molecular mass but expressed in g/mol.
ⓓ. Molar mass has no relation to molecular mass.
Correct Answer: Molar mass is numerically equal to molecular mass but expressed in g/mol.
Explanation: For example, H₂O has molecular mass 18 u, so its molar mass is 18 g/mol. CO₂ has molecular mass 44 u, so its molar mass is 44 g/mol.
279. The molar mass of calcium carbonate ($\mathrm{CaCO_3}$) is:
ⓐ. 90 g/mol
ⓑ. 92 g/mol
ⓒ. 98 g/mol
ⓓ. 100 g/mol
Correct Answer: 100 g/mol
Explanation: Ca = 40, C = 12, O = 16 × 3 = 48. Total = 100 g/mol.
280. How many grams of oxygen gas are present in 0.25 mole of O₂?
ⓐ. 4 g
ⓑ. 8 g
ⓒ. 16 g
ⓓ. 32 g
Correct Answer: 8 g
Explanation: Molar mass of O₂ = 32 g/mol. Mass = moles × molar mass = 0.25 × 32 = 8 g.
281. Which compound has the molar mass of 98 g/mol?
ⓐ. $\mathrm{H_2SO_4}$
ⓑ. $\mathrm{HNO_3}$
ⓒ. $\mathrm{NaOH}$
ⓓ. $\mathrm{CaO}$
Correct Answer: $\mathrm{H_2SO_4}$
Explanation: H = 2, S = 32, O = 16 × 4 = 64. Total = 98 g/mol.
282. The molar mass of ammonia ($\mathrm{NH_3}$) is:
ⓐ. 15 g/mol
ⓑ. 16 g/mol
ⓒ. 17 g/mol
ⓓ. 18 g/mol
Correct Answer: 17 g/mol
Explanation: N = 14, H = 1 × 3 = 3. Total = 14 + 3 = 17 g/mol.
283. Why is molar mass important in stoichiometric calculations?
ⓐ. It helps calculate number of atoms only.
ⓑ. It connects the microscopic mass of molecules with measurable grams in the laboratory.
ⓒ. It determines colour of a compound.
ⓓ. It gives the charge of ions.
Correct Answer: It connects the microscopic mass of molecules with measurable grams in the laboratory.
Explanation: Molar mass serves as a bridge between atomic-scale measurements (u) and laboratory-scale measurements (g), enabling chemists to convert mass into moles and relate to Avogadro’s number.
284. What is meant by the molar volume of a gas at STP?
ⓐ. The number of liters occupied by 1 gram of a gas at 25°C and 1 atm.
ⓑ. The volume occupied by 1 mole of any gas at 0°C and 1 atm pressure.
ⓒ. The mass of 1 mole of gas.
ⓓ. The number of molecules in 22.4 liters of gas.
Correct Answer: The volume occupied by 1 mole of any gas at 0°C and 1 atm pressure.
Explanation: By Avogadro’s law, equal volumes of gases at the same temperature and pressure contain equal numbers of molecules. At STP (0°C, 1 atm), 1 mole of an ideal gas occupies 22.4 liters. This value is independent of the type of gas, making it a universal constant for ideal gas calculations.
285. At STP, what volume is occupied by 2 moles of nitrogen gas ($N_2$)?
ⓐ. 11.2 L
ⓑ. 22.4 L
ⓒ. 33.6 L
ⓓ. 44.8 L
Correct Answer: 44.8 L
Explanation: 1 mole of any gas at STP = 22.4 L. Therefore, 2 moles = $2 \times 22.4 = 44.8$ L. This relation holds for all ideal gases and is a direct application of the molar volume concept.
286. How many liters of oxygen gas are present at STP in 0.5 mole?
ⓐ. 5.6 L
ⓑ. 11.2 L
ⓒ. 16.8 L
ⓓ. 22.4 L
Correct Answer: 11.2 L
Explanation: Molar volume at STP = 22.4 L per mole. For 0.5 mole, the volume is $0.5 \times 22.4 = 11.2$ L. This shows the linear relationship between number of moles and gas volume.
287. If 67.2 L of hydrogen gas is taken at STP, how many moles does it represent?
ⓐ. 2 moles
ⓑ. 3 moles
ⓒ. 4 moles
ⓓ. 5 moles
Correct Answer: 3 moles
Explanation: Number of moles = Volume ÷ 22.4 L = $67.2 ÷ 22.4 = 3$. Thus, 67.2 L corresponds to 3 moles of hydrogen gas under STP conditions.
288. Which law explains why all gases occupy 22.4 L per mole at STP?
ⓐ. Boyle’s law
ⓑ. Charles’s law
ⓒ. Avogadro’s law
ⓓ. Dalton’s law
Correct Answer: Avogadro’s law
Explanation: Avogadro’s law states that equal volumes of gases at the same temperature and pressure contain equal numbers of molecules. Hence, 1 mole of any gas at STP contains $6.022 \times 10^{23}$ molecules and occupies 22.4 L.
289. How many molecules of CO₂ are present in 22.4 L of CO₂ at STP?
ⓐ. $6.022 \times 10^{22}$
ⓑ. $6.022 \times 10^{23}$
ⓒ. $3.011 \times 10^{23}$
ⓓ. $1.204 \times 10^{24}$
Correct Answer: $6.022 \times 10^{23}$
Explanation: 22.4 L of any gas at STP = 1 mole = Avogadro’s number of molecules. Therefore, 22.4 L CO₂ contains exactly $6.022 \times 10^{23}$ molecules.
290. A sample of ammonia gas occupies 44.8 L at STP. How many molecules are present?
ⓐ. $6.022 \times 10^{23}$
ⓑ. $1.204 \times 10^{24}$
ⓒ. $2 \times 6.022 \times 10^{23}$
ⓓ. Both B and C
Correct Answer: Both B and C
Explanation: 44.8 L = 2 moles at STP. Number of molecules = $2 \times 6.022 \times 10^{23} = 1.204 \times 10^{24}$. Both expressions B and C represent the same value.
291. Which of the following is not true about molar volume of gases?
ⓐ. It is constant for all ideal gases at STP.
ⓑ. It is independent of the type of gas.
ⓒ. It is equal to 22.4 L per mole at 0°C and 1 atm.
ⓓ. It varies for different gases at the same conditions.
Correct Answer: It varies for different gases at the same conditions.
Explanation: By Avogadro’s law, all ideal gases occupy the same volume per mole at STP, regardless of type. Thus, D is false.
292. What volume will 3.5 moles of O₂ gas occupy at STP?
ⓐ. 56 L
ⓑ. 70 L
ⓒ. 78.4 L
ⓓ. 84 L
Correct Answer: 78.4 L
Explanation: Volume = moles × molar volume = $3.5 \times 22.4 = 78.4$ L. This proportional relationship is essential in gas stoichiometry.
293. Why is the molar volume of gases important in chemical calculations?
ⓐ. It eliminates the need to measure gas pressure.
ⓑ. It allows chemists to directly relate gas volume to moles in stoichiometric problems.
ⓒ. It determines the colour of gases.
ⓓ. It explains why gases are lighter than solids.
Correct Answer: It allows chemists to directly relate gas volume to moles in stoichiometric problems.
Explanation: Since 1 mole of any gas at STP always occupies 22.4 L, chemists can easily calculate moles from gas volumes and use them in reaction stoichiometry. This simplifies balancing equations and predicting product yields involving gaseous substances.
294. What is the general formula for percentage composition of an element in a compound?
ⓐ. $\%\text{element} = \dfrac{\text{mass of element in 1 mol of compound}}{\text{molar mass of compound}} \times 100$
ⓑ. $\%\text{element} = \dfrac{\text{molar mass of compound}}{\text{mass of element}} \times 100$
ⓒ. $\%\text{element} = \dfrac{\text{atomic number}}{\text{mass number}} \times 100$
ⓓ. $\%\text{element} = \dfrac{\text{number of atoms of element}}{100}$
Correct Answer: $\%\text{element} = \dfrac{\text{mass of element in 1 mol of compound}}{\text{molar mass of compound}} \times 100$
Explanation: Percentage composition gives the proportion of each element in a compound by mass. It is always calculated by dividing the contribution of the element by the total molar mass and multiplying by 100.
295. Calculate the percentage of oxygen in water ($\mathrm{H_2O}$).
ⓐ. 66.6%
ⓑ. 33.3%
ⓒ. 88.9%
ⓓ. 50%
Correct Answer: 88.9%
Explanation: Molar mass of H₂O = $2 \times 1 + 16 = 18$. Mass of oxygen = 16. % of oxygen = $\dfrac{16}{18} \times 100 = 88.9\%$. Hydrogen contributes only 11.1%.
296. What is the percentage of carbon in carbon dioxide ($\mathrm{CO_2}$)?
ⓐ. 12%
ⓑ. 27.3%
ⓒ. 33.3%
ⓓ. 44%
Correct Answer: 27.3%
Explanation: Molar mass of CO₂ = 44. Mass of C = 12. %C = $\dfrac{12}{44} \times 100 = 27.3\%$. Oxygen contributes 72.7%.
297. In $\mathrm{NaCl}$, the percentage of chlorine is:
ⓐ. 35.5%
ⓑ. 50%
ⓒ. 60.7%
ⓓ. 70%
Correct Answer: 60.7%
Explanation: Molar mass = 23 + 35.5 = 58.5. Mass of Cl = 35.5. %Cl = $\dfrac{35.5}{58.5} \times 100 = 60.7\%$. Sodium contributes 39.3%.
298. Which of the following correctly gives the percentage composition of oxygen in sulphur dioxide ($\mathrm{SO_2}$)?
ⓐ. 33.3%
ⓑ. 66.7%
ⓒ. 75%
ⓓ. 50%
Correct Answer: 50%
Explanation: Molar mass of SO₂ = $32 + (2 \times 16) = 64$. Mass of oxygen = 32. %O = $\dfrac{32}{64} \times 100 = 50\%$.
299. What is the percentage composition of nitrogen in ammonia ($\mathrm{NH_3}$)?
ⓐ. 70%
ⓑ. 75%
ⓒ. 82.4%
ⓓ. 87.5%
Correct Answer: 82.4%
Explanation: Molar mass of NH₃ = 17. Mass of N = 14. %N = $\dfrac{14}{17} \times 100 \approx 82.4\%$. Hydrogen contributes 17.6%.
300. In $\mathrm{CaCO_3}$, what is the percentage of calcium?
ⓐ. 20%
ⓑ. 30%
ⓒ. 40%
ⓓ. 50%
Correct Answer: 40%
Explanation: Molar mass = 40 + 12 + (16 × 3) = 100. Mass of Ca = 40. %Ca = $\dfrac{40}{100} \times 100 = 40\%$.
The study of Some Basic Concepts of Chemistry provides the fundamental language of Chemistry, covering key definitions, rules, and mathematical calculations. As per the NCERT/CBSE Class 11 Chemistry syllabus, this chapter is of great importance for both board exams and competitive exams like JEE, NEET, state-level medical and engineering entrance tests. It introduces and strengthens topics such as mole calculations, percentage composition, empirical and molecular formula derivations, laws of chemical combination, volumetric analysis, and the calculation of theoretical and actual yields. In total, there are 394 MCQs with solutions, arranged into 4 learning-friendly parts. Here in Part 3, you will practice another 100 MCQs with detailed answers that challenge your application of concepts and help you score higher in exams.
- 👉 Total MCQs in this chapter: 394.
- 👉 This page contains: Third set of 100 solved MCQs.
- 👉 Important for concept-building, board exams, and competitive exams (JEE/NEET).
- 👉 To explore other chapters, subjects, or classes, click the top navigation bar above.
- 👉 To proceed further, click the Part 4 button above.