Class 9 Science Chapter 9:
Atomic Foundations of Matter
Learn chemistry in simple English. Understand the laws of chemical combination, Dalton's atomic theory, ionic and covalent bonds, chemical formulae, molecular mass and formula unit mass. This chapter includes easy examples, solved numericals and important exam questions with answers.
1. Introduction: What Is Matter?
Everything around us is made of matter. Air, water, soil, books, food and our bodies are examples of matter. Matter has mass and occupies space.
Matter is made up of extremely small particles called atoms. Atoms can join together to form molecules and compounds. Different arrangements of atoms produce different substances.
Important words to understand
Atom
The smallest unit of an element that retains its chemical identity.
Example: He represents one helium atom.Molecule
A group of two or more atoms chemically bonded together.
Examples: H₂, O₂ and H₂O.Element
A pure substance made of only one type of atom.
Examples: Iron, oxygen and copper.Compound
A substance formed when different elements combine chemically in a fixed proportion.
Examples: Water (H₂O) and carbon dioxide (CO₂).2. Laws of Chemical Combination
When elements combine to form compounds, they follow certain scientific rules. These rules are called the laws of chemical combination.
Law 1: Law of Conservation of Mass
This law is associated with the French scientist Antoine Lavoisier. It states that mass is neither created nor destroyed during a chemical reaction.
Reactants are the substances that take part in a reaction. Products are the new substances formed.
Example: If 12 g of carbon reacts completely with 32 g of oxygen, 44 g of carbon dioxide is formed.
12 g + 32 g = 44 gThe mass has not disappeared. The atoms have rearranged to form a new substance.
Law 2: Law of Constant Proportions
This law is associated with Joseph Proust. It states that a pure chemical compound always contains the same elements combined in the same fixed proportion by mass, regardless of its source or method of preparation.
Example: Water
Every sample of pure water contains hydrogen and oxygen in a mass ratio of 1:8.
| Sample of pure water | Hydrogen mass | Oxygen mass | Ratio H:O |
|---|---|---|---|
| 9 g | 1 g | 8 g | 1:8 |
| 18 g | 2 g | 16 g | 1:8 |
| 27 g | 3 g | 24 g | 1:8 |
Notice that the quantities change, but the ratio remains the same. That is why water is always represented by the chemical formula H₂O.
Conservation of mass: The total mass before and after a chemical reaction remains equal.
Constant proportions: A pure compound always contains its elements in a fixed ratio by mass.
3. Dalton's Atomic Theory
John Dalton proposed an atomic theory in the early nineteenth century. His theory helped explain why elements combine in definite proportions.
Main ideas of Dalton's theory
- All matter is made up of tiny particles called atoms.
- Atoms of different elements have different properties.
- Atoms combine in simple whole-number ratios to form compounds.
- Chemical reactions involve the rearrangement of atoms.
- In ordinary chemical reactions, atoms are not created or destroyed.
Why is Dalton's theory important?
It explains why chemical compounds have definite compositions and why chemical reactions rearrange atoms instead of creating new matter from nothing.
4. How Do Atoms Combine?
Atoms combine through chemical bonds. Many atoms become more stable when their outermost electron shell is filled. Atoms can share electrons or transfer electrons to form bonds.
Valence electrons
Electrons in the outermost shell of an atom are called valence electrons. These electrons play an important role in chemical bonding.
| Atom | Electronic configuration | Valence electrons |
|---|---|---|
| Hydrogen | 1 | 1 |
| Carbon | 2, 4 | 4 |
| Oxygen | 2, 6 | 6 |
| Sodium | 2, 8, 1 | 1 |
| Magnesium | 2, 8, 2 | 2 |
| Chlorine | 2, 8, 7 | 7 |
The first electron shell can hold up to 2 electrons. In the simple shell model used at this level, the second shell can hold up to 8. Many atoms become more stable when their outer shell has a complete arrangement of electrons.
What is an ion?
An ion is an atom or group of atoms carrying an electric charge because electrons have been lost or gained.
Positive ion (cation)
An atom loses one or more electrons and becomes positively charged.
Na → Na⁺ + e⁻Sodium forms a positive ion.
Negative ion (anion)
An atom gains one or more electrons and becomes negatively charged.
Cl + e⁻ → Cl⁻Chlorine forms a negative ion.
5. Covalent and Ionic Bonds
A. Covalent bond: Sharing of electrons
A covalent bond forms when atoms share one or more pairs of electrons. This type of bond is commonly found in molecules formed by non-metals.
Each hydrogen atom has one electron. By sharing a pair of electrons, the two hydrogen atoms achieve a stable arrangement for their first shell.
H + H → H₂Other examples: O₂, N₂, H₂O, CO₂ and CH₄.
B. Ionic bond: Transfer of electrons
An ionic bond forms because oppositely charged ions attract one another. It commonly forms when electrons are transferred from a metal atom to a non-metal atom.
- Sodium loses one electron and becomes Na⁺.
- Chlorine gains that electron and becomes Cl⁻.
- The oppositely charged ions attract each other.
- The resulting compound has the formula NaCl.
| Feature | Covalent compounds | Ionic compounds |
|---|---|---|
| How the bond forms | Sharing of electrons | Attraction between oppositely charged ions formed through electron transfer |
| Common elements | Usually non-metals | Often metals and non-metals |
| Structure | Often separate molecules | Usually a large ionic lattice |
| Melting point | Often relatively low for simple molecular substances | Usually high |
| Electrical conductivity | Usually poor | Conduct when molten or dissolved in water if mobile ions are present |
| Examples | H₂O, CO₂, CH₄ | NaCl, MgO, CaCl₂ |
6. Writing Chemical Formulae
A chemical formula tells us which elements are present and how many atoms of each element are represented in a molecule or formula unit.
Understanding symbols and subscripts
- H₂O contains 2 hydrogen atoms and 1 oxygen atom.
- CO₂ contains 1 carbon atom and 2 oxygen atoms.
- CaCl₂ contains 1 calcium ion for every 2 chloride ions in the formula unit.
- In 3H₂O, the coefficient 3 means three water molecules, containing 6 hydrogen atoms and 3 oxygen atoms in total.
Method 1: Writing ionic formulae using charges
Follow these steps:
- Write the positive ion first and the negative ion second.
- Write their charges.
- Choose the smallest whole-number ratio that makes the total charge zero.
- Write the numbers as subscripts and omit the number 1.
Magnesium ion = Mg²⁺
Chloride ion = Cl⁻
One Mg²⁺ needs two Cl⁻ ions to balance its charge.
Mg²⁺ + 2Cl⁻ → MgCl₂The formula is MgCl₂.
Aluminium ion = Al³⁺
Oxide ion = O²⁻
Two aluminium ions give a total charge of +6. Three oxide ions give a total charge of −6.
2Al³⁺ + 3O²⁻ → Al₂O₃The formula is Al₂O₃.
Common ions and their charges
| Positive ions | Negative ions |
|---|---|
| Sodium: Na⁺ | Chloride: Cl⁻ |
| Potassium: K⁺ | Oxide: O²⁻ |
| Magnesium: Mg²⁺ | Sulphide: S²⁻ |
| Calcium: Ca²⁺ | Hydroxide: OH⁻ |
| Aluminium: Al³⁺ | Nitrate: NO₃⁻ |
| Ammonium: NH₄⁺ | Sulphate: SO₄²⁻ |
| Carbonate: CO₃²⁻ |
Formulae of compounds to remember
| Compound | Formula |
|---|---|
| Water | H₂O |
| Carbon dioxide | CO₂ |
| Sodium chloride | NaCl |
| Magnesium oxide | MgO |
| Magnesium chloride | MgCl₂ |
| Calcium chloride | CaCl₂ |
| Aluminium oxide | Al₂O₃ |
| Calcium carbonate | CaCO₃ |
| Sodium sulphate | Na₂SO₄ |
| Calcium hydroxide | Ca(OH)₂ |
| Ammonium chloride | NH₄Cl |
7. Properties of Ionic and Covalent Compounds
Ionic compounds
- Usually consist of positive and negative ions arranged in a lattice.
- Often have high melting and boiling points.
- Many are crystalline solids at room temperature.
- Conduct electricity when molten or in a solution if their ions are mobile.
- Many dissolve in water, although solubility varies from compound to compound.
Covalent compounds
- Form when atoms share electrons.
- Many simple molecular compounds have relatively low melting and boiling points.
- Many do not conduct electricity because they do not contain mobile ions.
- They can be solids, liquids or gases at room temperature.
8. Molecular Mass and Formula Unit Mass
What is atomic mass?
Atomic mass is commonly expressed in unified atomic mass units (u). One unified atomic mass unit is defined as one-twelfth of the mass of a carbon-12 atom.
Useful approximate atomic masses for calculations:
| Element | Symbol | Approximate atomic mass |
|---|---|---|
| Hydrogen | H | 1 u |
| Carbon | C | 12 u |
| Nitrogen | N | 14 u |
| Oxygen | O | 16 u |
| Sodium | Na | 23 u |
| Magnesium | Mg | 24 u |
| Sulphur | S | 32 u |
| Chlorine | Cl | 35.5 u |
| Calcium | Ca | 40 u |
Molecular mass
Molecular mass is the sum of the atomic masses of all atoms present in one molecule.
Molecular mass = Sum of atomic masses of all atomsFormula unit mass
Ionic compounds form extended lattices rather than separate molecules. Their formula unit mass is calculated by adding the atomic masses represented in the chemical formula.
For example, NaCl has a formula unit mass of 23 + 35.5 = 58.5 u.
9. Solved Numericals: Step-by-Step
Read each question carefully. Write the formula, substitute the values and show your calculation. This makes your answer easier to understand and earns credit for working.
Numerical 1: Find the mass of carbon dioxide formed.
Question: 12 g of carbon reacts completely with 32 g of oxygen. What mass of carbon dioxide is formed?
Mass of reactants = Mass of products.
Total mass = 12 g + 32 g = 44 g.
Numerical 2: Find the missing mass.
Question: In a chemical reaction, 25 g of reactants produce 9 g of one product and another product of unknown mass. Find the missing mass, assuming all products are included.
Numerical 3: Apply the law of constant proportions.
Question: Water contains hydrogen and oxygen in a mass ratio of 1:8. Calculate the masses of hydrogen and oxygen in 45 g of pure water.
Oxygen = 8 × 5 = 40 g.
Numerical 4: Calculate the molecular mass of water (H₂O).
Given: H = 1 u, O = 16 u.
Numerical 5: Calculate the molecular mass of carbon dioxide (CO₂).
Given: C = 12 u, O = 16 u.
CO₂ = 12 + (2 × 16) = 12 + 32 = 44 u.
Numerical 6: Calculate the molecular mass of sulphuric acid (H₂SO₄).
Given: H = 1 u, S = 32 u, O = 16 u.
H₂SO₄ = (2 × 1) + 32 + (4 × 16)
= 2 + 32 + 64 = 98 u.
Numerical 7: Calculate the formula unit mass of sodium chloride (NaCl).
Given: Na = 23 u, Cl = 35.5 u.
Formula unit mass = 23 + 35.5 = 58.5 u.
Numerical 8: Calculate the formula unit mass of calcium carbonate (CaCO₃).
Given: Ca = 40 u, C = 12 u, O = 16 u.
CaCO₃ = 40 + 12 + (3 × 16)
= 40 + 12 + 48 = 100 u.
Numerical 9: Calculate the formula unit mass of magnesium chloride (MgCl₂).
Given: Mg = 24 u, Cl = 35.5 u.
MgCl₂ = 24 + (2 × 35.5)
= 24 + 71 = 95 u.
Numerical 10: Find the mass ratio of hydrogen and oxygen in water.
Question: One molecule of water is H₂O. Calculate the ratio of the total mass of hydrogen to oxygen.
10. Important Short Questions and Answers
Q1. What is the law of conservation of mass?
Q2. What is the law of constant proportions?
Q3. Who proposed the atomic theory?
Q4. What are valence electrons?
Q5. What is a covalent bond?
Q6. What is an ionic bond?
Q7. What is a cation?
Q8. What is an anion?
Q9. What is the difference between a molecule and a compound?
Q10. What does the formula H₂O tell us?
Q11. Why is the formula of magnesium chloride MgCl₂?
Q12. Why does molten sodium chloride conduct electricity?
Q13. What is molecular mass?
Q14. What is formula unit mass?
11. Important Long Questions and Answers
Q1. Explain the law of conservation of mass with an example.
The law of conservation of mass states that mass is neither created nor destroyed during a chemical reaction.
For example, 12 g of carbon reacts with 32 g of oxygen to form 44 g of carbon dioxide.
12 g + 32 g = 44 gThe total mass before the reaction is 44 g, and the total mass after the reaction is also 44 g. Therefore, the law is obeyed.
Q2. Explain Dalton's atomic theory.
- Matter is made up of atoms.
- Atoms of different elements have different properties.
- Atoms combine in simple whole-number ratios to form compounds.
- Chemical reactions involve the rearrangement of atoms.
- Atoms are conserved in ordinary chemical reactions.
Dalton's theory was important because it helped explain the fixed composition of compounds and the conservation of mass during ordinary chemical reactions. Some parts of the original theory were later revised as scientists discovered subatomic particles and isotopes.
Q3. Differentiate between ionic and covalent bonding.
In covalent bonding, atoms share electron pairs. Water (H₂O) and carbon dioxide (CO₂) are examples of covalent substances.
In ionic bonding, electrons are transferred from one atom to another, forming oppositely charged ions that attract each other. Sodium chloride (NaCl) and magnesium oxide (MgO) are examples.
Ionic compounds generally have high melting points and conduct electricity when molten or in suitable aqueous solutions. Many simple covalent substances have lower melting points and do not conduct electricity.
Q4. Explain how to write the chemical formula of aluminium oxide.
Step 1: Write the ions: Al³⁺ and O²⁻.
Step 2: Find the smallest ratio that balances the charges.
Two aluminium ions give +6, and three oxide ions give −6. Their total charge is zero.
2(+3) + 3(−2) = 0Step 3: Write the formula as Al₂O₃.
Q5. Calculate the molecular mass of glucose (C₆H₁₂O₆).
Molecular mass = (6 × 12) + (12 × 1) + (6 × 16)
= 72 + 12 + 96
Molecular mass = 180 u12. Multiple-Choice Questions (MCQs)
Q1. Who is associated with the law of conservation of mass?
A. John Dalton
B. Antoine Lavoisier
C. Joseph Proust
D. Niels Bohr
Correct answer: B. Antoine Lavoisier
Q2. Which law explains the fixed mass ratio of elements in pure water?
A. Law of conservation of energy
B. Law of constant proportions
C. Law of gravity
D. Law of motion
Correct answer: B. Law of constant proportions
Q3. What is the molecular mass of H₂O?
A. 16 u
B. 17 u
C. 18 u
D. 20 u
Correct answer: C. 18 u
Q4. Which compound is ionic?
A. H₂O
B. CO₂
C. NaCl
D. CH₄
Correct answer: C. NaCl
Q5. What is the correct formula of aluminium oxide?
A. AlO
B. Al₂O₃
C. Al₃O₂
D. AlO₂
Correct answer: B. Al₂O₃
Q6. A covalent bond is formed by:
A. Sharing electrons
B. Sharing protons
C. Losing neutrons
D. Splitting the nucleus
Correct answer: A. Sharing electrons
Q7. What is the formula unit mass of NaCl?
A. 35.5 u
B. 46 u
C. 58.5 u
D. 60 u
Correct answer: C. 58.5 u
Q8. Which ion is positively charged?
A. Cl⁻
B. O²⁻
C. Na⁺
D. S²⁻
Correct answer: C. Na⁺
Practice numericals: Try these yourself
- Calculate the molecular mass of NH₃. (N = 14 u, H = 1 u)
- Calculate the molecular mass of CH₄. (C = 12 u, H = 1 u)
- Calculate the formula unit mass of MgO. (Mg = 24 u, O = 16 u)
- In a reaction, 18 g of reactants produce 10 g of one product. What is the mass of the other product if there are only two products?
- Find the masses of hydrogen and oxygen in 72 g of pure water, using the mass ratio 1:8.
- 17 u
- 16 u
- 40 u
- 8 g
- Hydrogen = 8 g; oxygen = 64 g
13. Quick Revision: Remember These Points
- Conservation of mass: Total reactant mass equals total product mass.
- Constant proportions: A pure compound has a fixed mass ratio of its elements.
- Dalton: Proposed an influential atomic theory.
- Covalent bond: Formed by sharing electrons.
- Ionic bond: Attraction between oppositely charged ions.
- Cation: Positive ion.
- Anion: Negative ion.
- Molecular mass: Sum of atomic masses in a molecule.
- Formula unit mass: Sum of atomic masses represented by an ionic formula.
- Water: H₂O; hydrogen-to-oxygen mass ratio = 1:8.
- Carbon dioxide: CO₂; molecular mass = 44 u.
- Sodium chloride: NaCl; formula unit mass = 58.5 u.
