ATOMS AND MOLECULES-QnA

Q1: Who proposed the atomic theory?

John Dalton proposed the atomic theory in 1808.

Q2: What does the law of conservation of mass state?

Mass can neither be created nor destroyed in a chemical reaction.

Q3: Define atom.

The smallest particle of an element that takes part in a chemical reaction.

Q4: Define molecule.

Two or more atoms chemically bonded together form a molecule.

Q5: What is the symbol of sodium?

Na (from Latin name Natrium).

Q6: Define atomic mass unit (amu).

It is one-twelfth of the mass of a carbon-12 atom.

Q7: What is the formula of water?

\(\mathrm{H_2O}\).

Q8: Who gave the law of definite proportions?

Joseph Proust.

Q9: Define valency.

The combining capacity of an atom.

Q10: What is Avogadro’s number?

\(\mathrm{6.022 × 10^{23}}\) particles per mole.

Q11: Name one polyatomic ion.

Ammonium ion \(\mathrm{(NH_4^+)}\).

Q12: What is the molecular mass of \(\mathrm{CO_2}\)?

44 u.

Q13: What is the atomicity of phosphorus \(\mathrm{(P_4)}\)?

4

Q14: Define compound.

A substance formed by chemical combination of two or more elements.

Q15: What is the chemical formula of calcium oxide?

CaO.

Q1: What is meant by a chemical formula?

A chemical formula represents the composition of a compound using symbols and numbers of atoms of each element present.

Q2: State the law of multiple proportions.

When two elements form more than one compound, the masses of one element combine with a fixed mass of the other in simple whole-number ratios.

Q3: Define relative atomic mass.

The ratio of the average mass of one atom to one-twelfth the mass of a carbon-12 atom.

Q4: Differentiate between atom and molecule.

Atom is the smallest unit of an element, while a molecule is formed when two or more atoms combine chemically.

Q5: What is the valency of aluminium and oxygen in aluminium oxide?

Aluminium = 3, Oxygen = 2.

Q6: What is meant by the term “mole”?

A mole is the amount of substance containing \(\mathrm{6.022 × 10^{23}}\) particles (atoms, molecules, or ions).

Q7: What is the molecular mass of \(\mathrm{H_2SO_4}\)?

(2×1) + 32 + (4×16) = 98 u.

Q8: What are polyatomic ions? Give examples.

Ions made of more than one atom. Example: \(\mathrm{NH_4^+, SO_4^{2-}}\).

Q9: How is molecular mass calculated?

Molecular mass = Sum of atomic masses of all atoms in a molecule.

Q10: Why is water a compound and not an element?

Because it consists of two elements—hydrogen and oxygen—combined in a fixed ratio (2:1).

Q1: Explain the postulates of Dalton’s Atomic Theory.

(1) Matter is made up of indivisible atoms. (2) Atoms of the same element are identical. (3) Atoms combine in simple ratios to form compounds. (4) Atoms cannot be created or destroyed. (5) Chemical reactions involve rearrangement of atoms.

Q2: Describe the laws of chemical combination with examples.

Law of Conservation of Mass: Total mass remains constant (e.g., NaCl formation). Law of Definite Proportions: A compound always contains elements in a fixed mass ratio (e.g., \(\mathrm{H_2O}\) has H:O = 1:8).

Q3: Explain the concept of a mole with suitable examples.

A mole represents \(\mathrm{6.022 × 10^{23}}\) particles. For example, 1 mole of \(\mathrm{H_2O}\) has \(\mathrm{6.022 × 10^{23}}\) molecules or 18 g of water.

Q4: How are chemical formulas written? Illustrate with examples.

Identify symbols and valencies, then cross-multiply to balance. Example: \(\mathrm{Mg^{2+}}\) and \(\mathrm{Cl^- \rightarrow MgCl_2}\); \(\mathrm{Al^{3+}}\) and \(\mathrm{O^{2-} \rightarrow Al_2O_3}\).

Q5: Explain how to calculate molecular mass with examples.

Add atomic masses of all atoms in a molecule. Example: \(\mathrm{CO_2}\) = 12 + (16×2) = 44 u; \(\mathrm{H_2SO_4}\) = 2(1) + 32 + 4(16) = 98 u.

Q1: Explain Dalton’s Atomic Theory in detail and discuss its limitations.

Dalton proposed that matter is made up of indivisible atoms, identical for each element, combining in fixed ratios to form compounds. Limitations: Atoms are divisible (discovery of subatomic particles), isotopes and isobars show atoms of the same element can differ in mass, and complex compounds don’t always form in simple ratios.

Q2: Describe the experimental verification of the Law of Conservation of Mass.

Lavoisier and Landolt conducted experiments showing that the total mass of reactants equals the total mass of products in a closed system. Example: When barium chloride reacts with sodium sulfate to form barium sulfate and sodium chloride, total mass remains unchanged.

Q3: How are atoms and molecules related to the concept of moles? Explain with examples.

One mole of any substance contains Avogadro’s number \(\mathrm{(6.022×10^{23})}\) of particles. Example: 1 mole of \(\mathrm{CO_2 = 6.022×10^{23}}\) molecules or 44 g. It links microscopic particles (atoms, molecules) to measurable quantities (grams).

Q4: Explain with examples how to calculate the molecular mass and formula mass of compounds.

Molecular mass = Sum of atomic masses of all atoms in a molecule (e.g., \(\mathrm{H_2O }\)= 18 u). Formula mass = Sum of atomic masses in an ionic compound’s formula unit (e.g., NaCl = 23 + 35.5 = 58.5 u).

Q5: Discuss the importance and applications of the mole concept in chemistry.

The mole concept helps convert microscopic particles into measurable quantities, balance chemical equations, calculate reactant-product quantities, determine empirical and molecular formulas, and understand gas laws. It is essential in stoichiometry and quantitative chemistry.

Q1: In a reaction, 5.3 g of sodium carbonate reacted with 6 g of acetic acid. The products were 2.2 g of carbon dioxide, 0.9 g water and 8.2 g of sodium acetate. Show that these observations are in agreement with the law of conservation of mass. sodium carbonate + acetic acid ? sodium acetate + carbon dioxide + water

Q2: Hydrogen and oxygen combine in the ratio of 1:8 by mass to form water. What mass of oxygen gas would be required to react completely with 3 g of hydrogen gas?

The question states: Hydrogen and oxygen combine in the ratio of 1:8 by mass to form water. You need to find the mass of oxygen required to react completely with 3 g of hydrogen.

Step-by-Step Solution:
If 1 g of hydrogen reacts with 8 g of oxygen,
then 3 g of hydrogen will react with:
Mass of oxygen required \[3\,g\times \frac{8}{1}=24\,g\] 24 grams of oxygen gas are required to react completely with 3 grams of hydrogen gas to form water.

Q3: Which postulate of Dalton’s atomic theory is the result of the law of conservation of mass?

The postulate of Dalton’s atomic theory that is the direct result of the law of conservation of mass is:

"Atoms can neither be created nor destroyed in a chemical reaction."

This means that during any chemical change, the total number and kind of atoms remain constant—they’re just rearranged to form new substances.
That’s why, even after a reaction, the total mass of all substances stays the same—no atoms have magically appeared or disappeared.
This beautifully connects Dalton’s ideas with Lavoisier’s law of conservation of mass, showing that matter is always conserved in chemical processes.

Q4: Which postulate of Dalton’s atomic theory can explain the law of definite proportions?

The postulate of Dalton’s atomic theory that explains the law of definite proportions is:

"Atoms of different elements combine in simple whole-number ratios to form compounds."

This means that no matter where a compound comes from or how it’s made, the elements always join together in the same exact proportions by mass because they’re combining in these fixed ratios.
So, water will always be made by combining hydrogen and oxygen in a definite mass ratio (2:16 or 1:8), no matter the source—just as Dalton described in his groundbreaking theory.
This directly supports the law of definite proportions, which says a chemical compound always contains its component elements in a fixed ratio by mass.

Q5: Define the atomic mass unit.

The atomic mass unit (abbreviated as amu or u) is a standard unit used to express the mass of atoms and subatomic particles. One atomic mass unit is defined as exactly one-twelfth the mass of a single atom of carbon-12. In other words, it’s a tiny “weighing scale” created by scientists so we can easily compare the masses of atoms, since they are far too small to measure in grams or kilograms.
Using amu gives us a simple and practical way to talk about and calculate the relative masses of different elements and molecules.

Q6: Why is it not possible to see an atom with naked eyes?

It is not possible to see an atom with the naked eyes because atoms are unimaginably tiny—so small that their size is measured in nanometers (a billionth of a meter). An atom’s diameter is typically about 0.1 nanometers, which is far smaller than the wavelength of visible light.

Since our eyes depend on light waves to detect and distinguish objects, anything much smaller than the wavelength of light (about 400–700 nanometers) remains completely invisible to us.

In other words, atoms are so minuscule that no matter how sharp your eyesight is, they simply fall below the limit of what can be seen without the help of powerful modern instruments like electron microscopes.

Q7: Write down the names of compounds represented by the following formulae: (i) \(\mathrm{Al_2 (SO_4 )_3}\) (ii) \(\mathrm{CaCl_2}\) (iii) \(\mathrm{ K_2SO_4}\) (iv) \(\mathrm{KNO_3}\) (v) \(\mathrm{CaCO_3}\)

Q8: What is meant by the term chemical formula?

A chemical formula is a concise way of expressing the composition of a substance using symbols and numbers. It tells you exactly which elements make up a compound and in what proportion they are combined. For example, the formula \(\mathrm{H_2O}\) stands for water, revealing that each molecule is made of two hydrogen atoms and one oxygen atom. In short, a chemical formula is like a secret code that reveals the building blocks of any substance in the language of chemistry.

Q9: How many atoms are present in a (i) \(\mathrm{H_2S}\) molecule and (ii) \(\mathrm{PO_4 ^{3–}}\) ion

(i) \(\mathrm{H_2S}\) molecule:
If you look at the formula\(\mathrm{H_2S}\), you'll notice it has a small “2” after the “H.” This means each molecule contains 2 atoms of hydrogen. The “S” stands for sulfur and, since it doesn’t have a subscript, there’s 1 atom of sulfur. So, altogether, one molecule of \(\mathrm{H_2S}\) has:

2 + 1 = 3 atoms (2 hydrogen, 1 sulfur)

(ii) \(\mathrm{PO_4 ^{3–}}\) ion (phosphate ion):
This formula has a “4” after the “O,” meaning there are 4 atoms of oxygen.
There’s no small number after the “P,” which means 1 atom of phosphorus. So, in one phosphate ion, there are:

1 + 4 = 5 atoms (1 phosphorus, 4 oxygen)

Summary:
In a \(\mathrm{H_2S}\) molecule, there are 3 atoms.
In a \(\mathrm{PO_4 ^{3–}}\) ion, there are 5 atoms.

Q10: Write down the formulae of (i) sodium oxide (ii) aluminium chloride (iii) sodium sulphide (iv) magnesium hydroxide

(i) Sodium oxide
Formula: \(\mathrm{Na_2O}\)

(ii) Aluminium chloride
Formula: \(\mathrm{AlCl_3}\)

(iii) Sodium sulphide
Formula: \(\mathrm{Na_2S}\)

(iv) Magnesium hydroxide
Formula: \(\mathrm{Mg(OH)_2}\)

Q11: Calculate the molecular masses of \(\mathrm{H_2 ,\; O_2 \;, Cl_2 \;, CO_2 \;, CH_4 \;, C_2 H_6\; , C_2 H_4\; , NH_3 \;, CH_3 OH}\).

Q12: Calculate the formula unit masses of \(\mathrm{ZnO,\; Na_2O,\; K_2CO_3}\) , given atomic masses of Zn = 65 u, Na = 23 u, K = 39 u, C = 12 u, and O = 16 u.