Atom and Molecule notes with NCERT solution – Class 9 Science Chapter Explained
Complete notes, NCERT solutions and practice questions
Table of Contents
Introduction to Atoms and Molecules
The concept of atoms dates back to ancient times. In India, around 500 BC, an Indian philosopher Maharishi Kanad postulated that if we keep dividing matter (padarth), we would eventually reach a stage where we cannot divide it any further. He named these particles Parmanu. Around the same era, ancient Greek philosophers like Democritus and Leucippus suggested similar ideas, with Democritus calling these indivisible particles atoms (meaning indivisible).
In the modern scientific context, atoms are the basic units of matter and the defining structure of elements. They are made up of three subatomic particles: protons, neutrons, and electrons. Atoms combine to form molecules, which are the smallest identifiable units of a substance that retain all the properties of that substance.
Key Points:
- Atoms are the fundamental building blocks of matter
- Molecules are groups of two or more atoms bonded together
- Compounds are substances made up of two or more different elements in fixed ratio
Laws of Chemical Combination
By the end of the eighteenth century, scientists had established the difference between elements and compounds and became interested in understanding how elements combine. Antoine Lavoisier established two important laws of chemical combination through extensive experimentation.
Law of Conservation of Mass
According to the law of conservation of mass, mass can neither be created nor destroyed in a chemical reaction. This means that the total mass of the products in a chemical reaction is equal to the total mass of the reactants.
Example:
When 5.3 g of sodium carbonate reacts with 6 g of ethanoic acid, the products formed are 2.2 g of carbon dioxide, 0.9 g of water, and 8.2 g of sodium ethanoate.
Total mass of reactants: 5.3 g + 6 g = 11.3 g
Total mass of products: 8.2 g + 2.2 g + 0.9 g = 11.3 g
Since the total mass of reactants equals the total mass of products, the law of conservation of mass is verified.
Law of Constant Proportions
Also known as the law of definite proportions, this law states that in a chemical substance, the elements are always present in definite proportions by mass. This means that a pure chemical compound, regardless of its source or method of preparation, always contains the same elements combined in the same fixed ratio by mass.
Example:
In water (H₂O), hydrogen and oxygen are always present in the ratio of 1:8 by mass. This means that 9 g of water will always contain 1 g of hydrogen and 8 g of oxygen, regardless of the source of water.
Similarly, in ammonia (NH₃), nitrogen and hydrogen are always present in the ratio of 14:3 by mass.
Dalton’s Atomic Theory
John Dalton, a British scientist, proposed his atomic theory in 1808, which provided explanations for the laws of chemical combination. The postulates of Dalton’s atomic theory are:
- All matter is made up of very tiny particles called atoms, which participate in chemical reactions.
- Atoms are indivisible particles, which cannot be created or destroyed in a chemical reaction.
- Atoms of a given element are identical in mass and chemical properties.
- Atoms of different elements have different masses and chemical properties.
- Atoms combine in the ratio of small whole numbers to form compounds.
- The relative number and kinds of atoms are constant in a given compound.
While Dalton’s atomic theory successfully explained the laws of chemical combination, we now know that atoms can be further divided into subatomic particles (protons, neutrons, and electrons). Also, atoms of the same element can have different masses (isotopes).
What is an Atom?
An atom is the smallest particle of an element that can participate in a chemical reaction. Atoms are incredibly small – their radius is measured in nanometers (1 nm = 10⁻⁹ m). To give you an idea, millions of atoms stacked would barely make a layer as thick as a sheet of paper.
Relative Sizes:
- Atom of hydrogen: 10⁻¹⁰ m
- Molecule of water: 10⁻⁹ m
- Molecule of hemoglobin: 10⁻⁸ m
- Grain of sand: 10⁻⁴ m
Atoms of most elements are not capable of existing independently. They combine with other atoms to form molecules or ions.
Atomic Mass
One of the remarkable concepts that Dalton’s atomic theory proposed was that of atomic mass. According to the theory, each element has a characteristic atomic mass. Since determining the mass of an individual atom is difficult, scientists determined relative atomic masses using the laws of chemical combinations.
Atomic Mass Unit
The atomic mass unit (u) is defined as exactly one-twelfth (1/12th) the mass of one atom of carbon-12. The relative atomic masses of all elements have been found with respect to an atom of carbon-12.
| Element | Atomic Mass (u) |
|---|---|
| Hydrogen | 1 |
| Carbon | 12 |
| Nitrogen | 14 |
| Oxygen | 16 |
| Sodium | 23 |
| Magnesium | 24 |
| Sulfur | 32 |
| Chlorine | 35.5 |
| Calcium | 40 |
What is a Molecule?
A molecule is a group of two or more atoms that are chemically bonded together, meaning they are tightly held together by attractive forces. A molecule can be defined as the smallest particle of an element or a compound that is capable of an independent existence and shows all the properties of that substance. Atoms of the same element or different elements can join together to form molecules.
Molecules of Elements
The molecules of an element are constituted by the same type of atoms. The number of atoms constituting a molecule is known as its atomicity.
| Element | Atomicity | Example |
|---|---|---|
| Argon, Helium | Monoatomic | Ar, He |
| Hydrogen, Oxygen, Nitrogen, Chlorine | Diatomic | H₂, O₂, N₂, Cl₂ |
| Ozone | Triatomic | O₃ |
| Phosphorus | Tetra-atomic | P₄ |
| Sulfur | Polyatomic | S₈ |
Metals and some other elements, such as carbon, do not have a simple structure but consist of a very large and indefinite number of atoms bonded together.
Molecules of Compounds
Atoms of different elements join together in definite proportions to form molecules of compounds. For example:
| Compound | Combining Elements | Ratio by Mass |
|---|---|---|
| Water (H₂O) | Hydrogen, Oxygen | 1:8 |
| Ammonia (NH₃) | Nitrogen, Hydrogen | 14:3 |
| Carbon dioxide (CO₂) | Carbon, Oxygen | 3:8 |
Ions and Polyatomic Ions
Compounds composed of metals and non-metals contain charged species called ions. Ions may consist of a single charged atom or a group of atoms that have a net charge. An ion can be negatively charged (anion) or positively charged (cation).
For example, in sodium chloride (NaCl), the constituent particles are positively charged sodium ions (Na⁺) and negatively charged chloride ions (Cl⁻).
A group of atoms carrying a charge is known as a polyatomic ion. Examples include:
| Polyatomic Ion | Symbol | Charge |
|---|---|---|
| Ammonium | NH₄⁺ | +1 |
| Hydroxide | OH⁻ | -1 |
| Nitrate | NO₃⁻ | -1 |
| Carbonate | CO₃²⁻ | -2 |
| Sulfate | SO₄²⁻ | -2 |
| Phosphate | PO₄³⁻ | -3 |
Writing Chemical Formulas
The chemical formula of a compound is a symbolic representation of its composition. It shows the constituent elements and the number of atoms of each combining element. To write chemical formulas, we need to know the valencies of elements.
Valency is the combining power (or capacity) of an element. It can be thought of as the number of hands or arms that an atom has to make bonds with other atoms.
Rules for Writing Chemical Formulas:
- The valencies or charges on the ions must balance.
- When a compound consists of a metal and a non-metal, the name or symbol of the metal is written first.
- In compounds formed with polyatomic ions, the formula of the ion is enclosed in brackets if more than one such ion is present.
To write a chemical formula, we need to:
- Write the symbols of the constituent elements or ions.
- Write the valencies or charges below the respective symbols.
- Cross-multiply the valencies to find the subscripts for each element or ion.
- Simplify the ratios if possible.
Example 1: Magnesium Chloride
Magnesium (Mg) has a valency of +2, while chlorine (Cl) has a valency of -1.
To form a neutral compound, we need 1 Mg²⁺ and 2 Cl⁻.
Therefore, the formula of magnesium chloride is MgCl₂.
Example 2: Aluminium Oxide
Aluminium (Al) has a valency of +3, while oxygen (O) has a valency of -2.
To form a neutral compound, we need 2 Al³⁺ and 3 O²⁻.
Therefore, the formula of aluminium oxide is Al₂O₃.
Example 3: Calcium Hydroxide
Calcium (Ca) has a valency of +2, while the hydroxide ion (OH⁻) has a valency of -1.
To form a neutral compound, we need 1 Ca²⁺ and 2 OH⁻.
Therefore, the formula of calcium hydroxide is Ca(OH)₂.
Note: We use brackets to enclose the polyatomic ion since there are two hydroxide ions.
Mole Concept and Molecular Mass
Molecular Mass
The molecular mass of a substance is the sum of the atomic masses of all the atoms in a molecule of the substance. It is expressed in atomic mass units (u).
Example: Calculate the molecular mass of water (H₂O)
Atomic mass of hydrogen (H) = 1 u
Atomic mass of oxygen (O) = 16 u
Molecular mass of water (H₂O) = (2 × 1) + (1 × 16) = 2 + 16 = 18 u
Formula Unit Mass
The formula unit mass of a substance is the sum of the atomic masses of all atoms in a formula unit of a compound. It is calculated in the same manner as molecular mass. The term is generally used for ionic compounds, where the constituent particles are ions rather than molecules.
Example: Calculate the formula unit mass of sodium chloride (NaCl)
Atomic mass of sodium (Na) = 23 u
Atomic mass of chlorine (Cl) = 35.5 u
Formula unit mass of NaCl = 23 + 35.5 = 58.5 u
Mole Concept
The mole is a fundamental unit in chemistry that links the atomic and macroscopic scales. One mole of a substance is defined as the amount that contains exactly 6.022 × 10²³ elementary entities (atoms, molecules, ions, etc.). This number is known as Avogadro’s number.
The mass of one mole of a substance, expressed in grams, is called its molar mass. The molar mass of an element in grams is numerically equal to its atomic mass in atomic mass units (u).
Example 1: Find the mass of one atom of carbon
Given that one mole of carbon atoms weighs 12 g.
One mole of carbon contains 6.022 × 10²³ atoms.
Mass of one atom of carbon = 12 g ÷ 6.022 × 10²³ = 1.993 × 10⁻²³ g
Example 2: Compare the number of atoms in 100 g of sodium and 100 g of iron
Atomic mass of sodium (Na) = 23 u
Atomic mass of iron (Fe) = 56 u
Number of moles in 100 g of Na = 100 ÷ 23 = 4.35 moles
Number of atoms in 100 g of Na = 4.35 × 6.022 × 10²³ = 2.62 × 10²⁴ atoms
Number of moles in 100 g of Fe = 100 ÷ 56 = 1.79 moles
Number of atoms in 100 g of Fe = 1.79 × 6.022 × 10²³ = 1.08 × 10²⁴ atoms
Therefore, 100 g of sodium contains more atoms than 100 g of iron.
NCERT Questions and Solutions
In-Text Questions
Question 1: In a reaction, 5.3 g of sodium carbonate reacted with 6 g of ethanoic acid. The products were 2.2 g of carbon dioxide, 0.9 g water and 8.2 g of sodium ethanoate. Show that these observations are in agreement with the law of conservation of mass.
Answer: According to the law of conservation of mass, the total mass of the reactants should equal the total mass of the products.
Mass of reactants = Mass of sodium carbonate + Mass of ethanoic acid = 5.3 g + 6 g = 11.3 g
Mass of products = Mass of carbon dioxide + Mass of water + Mass of sodium ethanoate = 2.2 g + 0.9 g + 8.2 g = 11.3 g
Since the mass of reactants equals the mass of products, the observations are in agreement with the law of conservation of mass.
Question 2: 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?
Answer: The ratio of hydrogen to oxygen by mass in water is 1:8.
If 1 g of hydrogen combines with 8 g of oxygen, then 3 g of hydrogen will combine with:
Mass of oxygen = 8 g × 3 = 24 g
Therefore, 24 g of oxygen gas would be required to react completely with 3 g of hydrogen gas.
Question 3: Which postulate of Dalton’s atomic theory is the result of the law of conservation of mass?
Answer: The postulate of Dalton’s atomic theory that is the result of the law of conservation of mass is: “Atoms cannot be created or destroyed in a chemical reaction.” This means that the total number of atoms before and after a chemical reaction remains the same, which aligns with the law of conservation of mass.
Question 4: Which postulate of Dalton’s atomic theory can explain the law of definite proportions?
Answer: The postulate of Dalton’s atomic theory that explains the law of definite proportions is: “The relative number and kinds of atoms are constant in a given compound.” This means that a compound always contains the same elements in the same proportion by mass, regardless of its source or method of preparation.
Question 5: Define the atomic mass unit.
Answer: The atomic mass unit (u) is defined as exactly one-twelfth (1/12th) the mass of one atom of carbon-12. It is a standard unit used to express atomic and molecular masses.
Question 6: Why is it not possible to see an atom with naked eyes?
Answer: Atoms are extremely small, with sizes on the order of 10⁻¹⁰ meters (0.1 nanometers). This is far smaller than the wavelength of visible light (400-700 nanometers), making it impossible for the human eye to resolve individual atoms. Even with powerful optical microscopes, atoms cannot be seen because their size is smaller than the wavelength of light used for imaging.
Exercise Questions
Question 1: A 0.24 g sample of compound of oxygen and boron was found by analysis to contain 0.096 g of boron and 0.144 g of oxygen. Calculate the percentage composition of the compound by weight.
Answer:
Total mass of the compound = 0.24 g
Mass of boron = 0.096 g
Mass of oxygen = 0.144 g
Percentage of boron = (0.096 ÷ 0.24) × 100 = 40%
Percentage of oxygen = (0.144 ÷ 0.24) × 100 = 60%
Therefore, the percentage composition of the compound is 40% boron and 60% oxygen by weight.
Question 2: When 3.0 g of carbon is burnt in 8.00 g oxygen, 11.00 g of carbon dioxide is produced. What mass of carbon dioxide will be formed when 3.00 g of carbon is burnt in 50.00 g of oxygen? Which law of chemical combination will govern your answer?
Answer:
When 3.0 g of carbon is burnt in 8.00 g of oxygen, 11.00 g of carbon dioxide is produced. This means all 3.0 g of carbon combined with 8.00 g of oxygen to form carbon dioxide.
When 3.00 g of carbon is burnt in 50.00 g of oxygen, only 3.00 g of carbon is available to react. According to the first reaction, 3.00 g of carbon will need only 8.00 g of oxygen to react completely. The remaining oxygen (50.00 – 8.00 = 42.00 g) will remain unreacted.
Therefore, the mass of carbon dioxide formed will still be 11.00 g.
This is governed by the Law of Constant Proportions, which states that elements combine in a fixed ratio by mass to form a compound. In this case, carbon and oxygen combine in a fixed ratio (regardless of how much oxygen is available) to form carbon dioxide.
Question 3: What are polyatomic ions? Give examples.
Answer:
Polyatomic ions are charged entities composed of two or more atoms that behave as a single unit in chemical reactions.
Examples:
- Hydroxide ion (OH⁻)
- Ammonium ion (NH₄⁺)
- Carbonate ion (CO₃²⁻)
- Sulfate ion (SO₄²⁻)
- Nitrate ion (NO₃⁻)
- Phosphate ion (PO₄³⁻)
Question 4: Write the chemical formulae of the following:
(a) Magnesium chloride
(b) Calcium oxide
(c) Copper nitrate
(d) Aluminium chloride
(e) Calcium carbonate
Answer:
(a) Magnesium chloride: MgCl₂
(b) Calcium oxide: CaO
(c) Copper nitrate: Cu(NO₃)₂
(d) Aluminium chloride: AlCl₃
(e) Calcium carbonate: CaCO₃
Question 5: Give the names of the elements present in the following compounds:
(a) Quick lime
(b) Hydrogen bromide
(c) Baking powder
(d) Potassium sulfate
Answer:
(a) Quick lime (CaO): Calcium and Oxygen
(b) Hydrogen bromide (HBr): Hydrogen and Bromine
(c) Baking powder (NaHCO₃): Sodium, Hydrogen, Carbon, and Oxygen
(d) Potassium sulfate (K₂SO₄): Potassium, Sulfur, and Oxygen
Question 6: Calculate the molar mass of the following substances:
(a) Ethyne, C₂H₂
(b) Sulfur molecule, S₈
(c) Phosphorus molecule, P₄ (Atomic mass of phosphorus = 31)
(d) Hydrochloric acid, HCl
(e) Nitric acid, HNO₃
Answer:
(a) Ethyne, C₂H₂:
Molar mass = (2 × 12) + (2 × 1) = 24 + 2 = 26 g/mol
(b) Sulfur molecule, S₈:
Molar mass = 8 × 32 = 256 g/mol
(c) Phosphorus molecule, P₄:
Molar mass = 4 × 31 = 124 g/mol
(d) Hydrochloric acid, HCl:
Molar mass = 1 + 35.5 = 36.5 g/mol
(e) Nitric acid, HNO₃:
Molar mass = 1 + 14 + (3 × 16) = 1 + 14 + 48 = 63 g/mol
Question 7: What is the mass of:
(a) 1 mole of nitrogen atoms?
(b) 4 moles of aluminium atoms (Atomic mass of aluminium = 27)?
(c) 10 moles of sodium sulfite (Na₂SO₃)?
Answer:
(a) Mass of 1 mole of nitrogen atoms:
Atomic mass of nitrogen = 14 u
Mass of 1 mole of nitrogen atoms = 14 g
(b) Mass of 4 moles of aluminium atoms:
Atomic mass of aluminium = 27 u
Mass of 1 mole of aluminium = 27 g
Mass of 4 moles of aluminium = 4 × 27 = 108 g
(c) Mass of 10 moles of sodium sulfite (Na₂SO₃):
Molar mass of Na₂SO₃ = (2 × 23) + 32 + (3 × 16) = 46 + 32 + 48 = 126 g/mol
Mass of 10 moles of Na₂SO₃ = 10 × 126 = 1260 g
Question 8: Convert into moles:
(a) 12 g of oxygen gas
(b) 20 g of water
(c) 22 g of carbon dioxide
Answer:
(a) Conversion of 12 g of oxygen gas (O₂) into moles:
Molar mass of O₂ = 2 × 16 = 32 g/mol
Number of moles = Given mass ÷ Molar mass = 12 ÷ 32 = 0.375 moles
(b) Conversion of 20 g of water (H₂O) into moles:
Molar mass of H₂O = (2 × 1) + 16 = 18 g/mol
Number of moles = Given mass ÷ Molar mass = 20 ÷ 18 = 1.11 moles
(c) Conversion of 22 g of carbon dioxide (CO₂) into moles:
Molar mass of CO₂ = 12 + (2 × 16) = 12 + 32 = 44 g/mol
Number of moles = Given mass ÷ Molar mass = 22 ÷ 44 = 0.5 moles
Question 9: What is the mass of:
(a) 0.2 mole of oxygen atoms?
(b) 0.5 mole of water molecules?
Answer:
(a) Mass of 0.2 mole of oxygen atoms:
Atomic mass of oxygen = 16 u
Mass of 1 mole of oxygen atoms = 16 g
Mass of 0.2 mole of oxygen atoms = 0.2 × 16 = 3.2 g
(b) Mass of 0.5 mole of water molecules:
Molar mass of H₂O = (2 × 1) + 16 = 18 g/mol
Mass of 0.5 mole of water = 0.5 × 18 = 9 g
Question 10: Calculate the number of molecules of sulfur (S₈) present in 16 g of solid sulfur.
Answer:
Molar mass of S₈ = 8 × 32 = 256 g/mol
Number of moles of S₈ = Given mass ÷ Molar mass = 16 ÷ 256 = 0.0625 moles
Number of molecules = Number of moles × Avogadro’s number
Number of molecules = 0.0625 × 6.022 × 10²³ = 3.76 × 10²² molecules
Question 11: Calculate the number of aluminium ions present in 0.051 g of aluminium oxide. (Hint: The mass of an ion is the same as that of an atom of the same element. Atomic mass of Al = 27 u)
Answer:
The formula of aluminium oxide is Al₂O₃.
Molar mass of Al₂O₃ = (2 × 27) + (3 × 16) = 54 + 48 = 102 g/mol
Number of moles of Al₂O₃ = Given mass ÷ Molar mass = 0.051 ÷ 102 = 0.0005 moles
Each mole of Al₂O₃ contains 2 moles of Al ions.
Number of moles of Al ions = 2 × Number of moles of Al₂O₃ = 2 × 0.0005 = 0.001 moles
Number of Al ions = Number of moles of Al ions × Avogadro’s number
Number of Al ions = 0.001 × 6.022 × 10²³ = 6.022 × 10²⁰ ions
Additional Practice Questions
Question 1: Define the following terms:
(a) Atomicity
(b) Valency
(c) Formula unit mass
(d) Mole
Answer:
(a) Atomicity: The number of atoms constituting a molecule is known as its atomicity. For example, the atomicity of hydrogen (H₂) is 2, and the atomicity of ozone (O₃) is 3.
(b) Valency: The combining capacity or power of an element is known as its valency. It represents the number of chemical bonds an atom can form with other atoms.
(c) Formula unit mass: The sum of the atomic masses of all atoms in a formula unit of a compound. It is expressed in atomic mass units (u).
(d) Mole: One mole is the amount of a substance that contains exactly 6.022 × 10²³ elementary entities (atoms, molecules, ions, etc.). It is the SI unit of the amount of substance.
Question 2: Calculate the molecular mass of the following:
(a) H₂SO₄
(b) C₆H₁₂O₆
(c) Ca₃(PO₄)₂
Answer:
(a) H₂SO₄:
Molecular mass = (2 × 1) + 32 + (4 × 16) = 2 + 32 + 64 = 98 u
(b) C₆H₁₂O₆ (glucose):
Molecular mass = (6 × 12) + (12 × 1) + (6 × 16) = 72 + 12 + 96 = 180 u
(c) Ca₃(PO₄)₂:
Formula unit mass = (3 × 40) + (2 × 31) + (8 × 16) = 120 + 62 + 128 = 310 u
Question 3: What is the difference between a molecule and an atom?
Answer:
An atom is the smallest particle of an element that can participate in a chemical reaction. It cannot usually exist independently.
A molecule is a group of two or more atoms chemically bonded together. It is the smallest particle of an element or compound that can exist independently and retain all the properties of that substance.
For example, oxygen exists as diatomic molecules (O₂) in nature, not as individual oxygen atoms. Similarly, a water molecule (H₂O) consists of two hydrogen atoms and one oxygen atom bonded together.
Question 4: If 5 moles of oxygen atoms weigh 80 g, then calculate the weight of one atom of oxygen in grams.
Answer:
Weight of 5 moles of oxygen atoms = 80 g
Weight of 1 mole of oxygen atoms = 80 ÷ 5 = 16 g
1 mole of oxygen atoms contains 6.022 × 10²³ atoms (Avogadro’s number)
Weight of one atom of oxygen = 16 ÷ (6.022 × 10²³) = 2.66 × 10⁻²³ g
Question 5: A sample of vitamin C is found to contain 2.58 × 10²⁴ oxygen atoms. How many moles of oxygen atoms are present in the sample?
Answer:
Number of oxygen atoms in the sample = 2.58 × 10²⁴ atoms
Number of moles of oxygen atoms = Number of atoms ÷ Avogadro’s number
Number of moles of oxygen atoms = 2.58 × 10²⁴ ÷ 6.022 × 10²³ = 4.28 moles
Question 6: Which has more number of atoms – 1 g of hydrogen or 1 g of helium?
Answer:
Atomic mass of hydrogen (H) = 1 u (Hydrogen exists as H₂ in nature)
Atomic mass of helium (He) = 4 u
Number of moles in 1 g of hydrogen atoms = 1 ÷ 1 = 1 mole
Number of atoms in 1 g of hydrogen = 1 × 6.022 × 10²³ = 6.022 × 10²³ atoms
Number of moles in 1 g of helium atoms = 1 ÷ 4 = 0.25 mole
Number of atoms in 1 g of helium = 0.25 × 6.022 × 10²³ = 1.506 × 10²³ atoms
Therefore, 1 g of hydrogen has more atoms than 1 g of helium.
Question 7: Write the chemical formulas for the following compounds and calculate their molecular masses:
(a) Sodium carbonate
(b) Ammonium sulfate
(c) Magnesium hydroxide
Answer:
(a) Sodium carbonate (Na₂CO₃):
Molecular mass = (2 × 23) + 12 + (3 × 16) = 46 + 12 + 48 = 106 u
(b) Ammonium sulfate ((NH₄)₂SO₄):
Molecular mass = (2 × 14) + (8 × 1) + 32 + (4 × 16) = 28 + 8 + 32 + 64 = 132 u
(c) Magnesium hydroxide (Mg(OH)₂):
Molecular mass = 24 + (2 × 16) + (2 × 1) = 24 + 32 + 2 = 58 u
Question 8: Explain how the law of constant proportions is a logical consequence of Dalton’s atomic theory.
Answer:
Dalton’s atomic theory states that atoms of different elements combine in simple whole-number ratios to form compounds. Additionally, it states that the relative number and kinds of atoms are constant in a given compound.
As a logical consequence, when elements combine to form a compound, they will always combine in the same fixed proportion by mass, regardless of the source of the compound or how it is prepared. This is precisely what the law of constant proportions states.
For example, according to Dalton’s theory, water is always formed when two hydrogen atoms combine with one oxygen atom. Since atoms of the same element have the same mass, the mass ratio of hydrogen to oxygen in water will always be the same (1:8 by mass). This fixed ratio of elements by mass in a compound is what the law of constant proportions describes.
Question 9: How many atoms are present in:
(a) 52 moles of He
(b) 52 u of He
Answer:
(a) Number of atoms in 52 moles of He:
Number of atoms = Number of moles × Avogadro’s number
Number of atoms = 52 × 6.022 × 10²³ = 3.13 × 10²⁵ atoms
(b) Number of atoms in 52 u of He:
The atomic mass of He is 4 u, which means one atom of helium has a mass of 4 u.
Number of atoms in 52 u of He = 52 ÷ 4 = 13 atoms
Question 10: Why do atoms of most elements not exist independently?
Answer:
Atoms of most elements do not exist independently because they are not stable in their isolated form. Atoms tend to either gain, lose, or share electrons to achieve a stable electronic configuration, typically resembling that of the nearest noble gas.
This tendency leads atoms to combine with other atoms through chemical bonds (ionic, covalent, or metallic) to form more stable structures like molecules or crystal lattices. For example, hydrogen exists as H₂ molecules, not as individual H atoms, because the shared electron pair in the H-H bond provides greater stability than isolated hydrogen atoms.
Only noble gases (like helium, neon, argon) generally exist as individual atoms because they already have stable electronic configurations with filled outer shells.
