CBSE Class 9 Science (Exploration) · Complete Notes · Read only this page → score 100%
🧠 Part 1 Map · Does Mass Change?
The Big Question
Water ≠ hydrogen + oxygen in properties
But mass of water = mass of H + O
Does mass stay the same?
Activity 9.1 — Physical Change
Salt dissolved in water
Mass of solution = water + salt
No mass change
Set-up 1 — Open System
Vinegar + baking soda
Balloon NOT fixed
Final ≠ initial (gas escapes)
Set-up 2 — Closed System
Balloon fixed with thread
Gas trapped inside
Final = initial ✅
PART 1 · IS MASS CONSERVED?
🤔 Think It Over
Water can be obtained from various sources. Are all these samples of water chemically identical?
Oxygen is sometimes represented as O and sometimes as O₂. What is the difference between these symbols?
Why does dissolved salt in water conduct electricity, but sugar does not?
Where this chapter starts
In Chapter 8 you learnt: atoms with an octet of electrons in their valence shell are stable; atoms can lose, gain or share electrons to achieve an octet.
Many properties of elements are not retained when they form a compound.
Example: hydrogen and oxygen are gases — but the substance they form, water, is a liquid at ordinary temperatures.
Hydrogen gas is combustible; oxygen supports combustion; but water neither burns nor helps in burning — it extinguishes fire.
However — the mass of the water formed equals the sum of the masses of the hydrogen and oxygen that combined to form it.
🔍 The question this chapter explores: Does mass remain unchanged during physical and chemical changes?
Activity 9.1: Let us investigate a physical change
Place a clean and dry 100 mL beaker on a digital weighing balance.
Set the balance reading to zero by pressing the tare or reset button.
Pour about 50 mL of water into the beaker.
Add a spatula full of common salt to the water in the beaker.
Record the reading on the weighing balance (Fig. 9.1a).
Swirl until the added salt dissolves and record your observations (Fig. 9.1b).
EXAM · OBSERVATION
The mass of the solution is equal to the sum of the masses of water and salt taken. There is practically no change in mass during the formation of a solution, which is a physical change. This is true for all physical changes. Repeat idea: weigh a piece of paper before and after tearing it into pieces — the mass does not change.
Activity 9.2: Let us investigate a chemical change
Recall what happened when baking soda was added to vinegar — a gas, carbon dioxide, was formed. The reaction is represented as:
Vinegar + Baking soda (Sodium hydrogencarbonate) → Carbon dioxide + Other substances
Let us explore whether the mass remains the same before and after this chemical change.
Experimental set-up 1
Place a clean, dry 100 mL conical flask and a medium-sized balloon on a weighing balance.
Set the balance reading to zero by pressing the tare or reset button.
Pour about 20 mL of vinegar or lemon juice into the conical flask.
Take about 2 g of baking soda (sodium hydrogencarbonate) and put it into the balloon.
Keep the balloon filled with baking soda on the weighing balance next to the conical flask. Record the initial reading (Fig. 9.2a).
Carefully transfer the baking soda from the balloon into the conical flask containing vinegar (Fig. 9.2b).
Place the conical flask and balloon back on the weighing balance, and record the final reading (Fig. 9.2c).
What do you observe? Are the initial and the final readings same?
📌 Note (NCERT)
Keep the conical flask and the balloon on the digital weighing balance. This prevents errors caused by small traces of baking soda that may remain stuck to the balloon.
📌 Note (NCERT)
Usually in any measurement, there is an uncertainty of ±1 in the last digit. Therefore, the variation in readings of a digital weighing balance is within experimental error, and the weight can be taken as constant.
👀 LOOK HERE: In set-up 1, a brisk effervescence is observed and the final reading does NOT match the initial reading. Ask yourself why — the answer is the whole point of set-up 2.
Experimental set-up 2
Place a clean, dry 100 mL conical flask and a medium-sized balloon on a weighing balance.
Set the balance reading to zero using the tare or reset button.
Pour about 20 mL of vinegar or lemon juice into the conical flask.
Place about 2 g of baking soda (sodium hydrogencarbonate) in the balloon.
Fix the balloon to the mouth of the conical flask using a thread, without allowing the baking soda to mix with the vinegar.
Weigh the conical flask containing vinegar and the balloon containing baking soda, and record the reading (Fig. 9.3a).
Lift the other end of the balloon upwards, allowing the baking soda to fall into the vinegar (Fig. 9.3b).
What do you observe? As in set-up 1, a brisk effervescence occurs, which inflates the balloon during the reaction.
Record the final reading (Fig. 9.3c). Are the initial and final readings same in this case?
Fig. 9.3: (c) Weight of the final reaction mixture — the balloon traps the gas
EXAM · KEY CONCLUSION
In set-up 2, the final reading matches the initial reading. The total mass of vinegar and baking soda before the reaction equals the total mass of carbon dioxide and other substances formed after the reaction.
Why set-up 1 failed: the mass difference occurs because the gas produced escapes into the air (open system), resulting in a difference between the initial and final readings.
📝 Check Your Concepts — Part 1
C1 In Activity 9.1, why does the mass not change when salt dissolves in water?
Show Answer
Dissolving is a physical change — no new substance is created or destroyed. The mass of the solution equals the sum of the masses of water and salt taken. This holds for all physical changes.
C2 Why is the final reading lower than the initial reading in Experimental set-up 1?
Show Answer
Because the carbon dioxide gas produced escapes into the surroundings — it is an open system. The mass is not destroyed; it simply leaves the balance.
C3 What single change makes set-up 2 give the correct result?
Show Answer
The balloon is fixed to the mouth of the conical flask with a thread, making it a closed system. The gas produced is trapped inside the balloon, so no mass leaves — initial reading = final reading.
C4 Why must the balloon be kept on the balance along with the flask?
Show Answer
To prevent errors caused by small traces of baking soda remaining stuck to the balloon — if the balloon is removed, that leftover mass is not counted.
C5 Water is made from two gases, yet it puts out fire. What does this show?
Show Answer
That the properties of elements are not retained when they form a compound. Hydrogen is combustible and oxygen supports combustion, but water neither burns nor helps burning — it extinguishes fire. Only the mass carries over unchanged.
🧠 Part 2 Map · Law of Conservation of Mass
The Law
Lavoisier, 1789
Matter neither created nor destroyed
Applies to every chemical reaction
Activity 9.3
Sodium sulfate + Barium chloride
White precipitate forms
Mass unchanged
Open vs Closed
Gas formed → use closed system
No gas → open system is fine
Numericals
Mass of reactants = mass of products
Unitary method for ratios
PART 2 · LAW OF CONSERVATION OF MASS
9.1 Law of Conservation of Mass
Activity 9.2 demonstrates that the total mass remains the same before and after a chemical reaction.
So, matter can neither be created nor destroyed in a chemical reaction.
EXAM · DEFINITION Law of Conservation of Mass: matter can neither be created nor destroyed in a chemical reaction — proposed by Antoine Lavoisier in 1789.
👨🔬 Meet a Scientist — Antoine Lavoisier
Known as the Father of Modern Chemistry.
Proposed the Law of Conservation of Mass, which applies to every chemical reaction.
He proposed that in every operation an equal quantity of matter exists both before and after the operation.
Activity 9.3: Let us verify the law — Group activity
Place two clean and dry 100 mL conical flasks on a weighing balance, and mark them A and B.
Set the balance reading to zero using the tare or reset button.
Pour about 10 mL of 1% m/v sodium sulfate solution into Conical Flask A.
In Conical Flask B, pour about 10 mL of 1% m/v barium chloride solution.
Leave both flasks on the weighing balance undisturbed, and record the total mass of both solutions (Fig. 9.4a).
Transfer the solution from Flask B to Flask A and mix the two solutions carefully. What do you observe?
Place both flasks A and B on the balance again (Fig. 9.4b), and note the reading.
Do you observe any change in the reading after mixing the solutions?
Fig. 9.4: (a) Weight of solutions (reactants) before mixingFig. 9.4: (b) Weight of products after mixing
👀 LOOK HERE: Compare the two balance readings in Fig. 9.4a and 9.4b — they are the same. A new substance formed (white precipitate), yet not a single gram was gained or lost.
Observation: when solutions of sodium sulfate and barium chloride are mixed, a white precipitate of barium sulfate is formed along with sodium chloride.
Based on Activity 9.2, it was established that weight — and hence mass — is conserved during a chemical change.
📌 Note (NCERT)
Keep both conical flasks on the digital weighing balance. This prevents the error caused by a small amount of solution that may remain stuck to the walls of the flask during transfer.
This reaction has been carried out in the open system, since no gas is formed.
🧩 Think as a Scientist (NCERT)
You are given a chemical reaction in which zinc reacts with dilute hydrochloric acid to form zinc chloride and hydrogen gas.
Design and perform an experiment to test the hypothesis that mass is conserved during the chemical reaction. You may use a set-up different from the one shown in Activity 9.2.
Show Expected Answer
Since hydrogen gas is produced, the experiment MUST be done in a closed system, or the gas will escape and the mass will appear to decrease.
Design: take dilute HCl in a conical flask and zinc granules in a balloon. Fix the balloon tightly on the mouth of the flask with thread. Weigh the whole set-up → initial reading.
Lift the balloon so the zinc falls into the acid. Effervescence occurs and the balloon inflates with hydrogen. Weigh again → final reading.
Result: initial reading = final reading → mass is conserved, hypothesis verified.
✏️ EXAMPLE 9.1 — METHOD, STUDY THIS Q: Students place 4.0 g of calcium carbonate with 2.92 g of hydrochloric acid in a closed container. After the reaction, they measured 1.76 g of carbon dioxide, 0.72 g of water, and 4.44 g of calcium chloride. Verify whether the Law of Conservation of Mass is obeyed.
Solution:
Mass of calcium carbonate = 4.0 g · Mass of hydrochloric acid = 2.92 g
Total mass of reactants = 4.0 g + 2.92 g = 6.92 g
Mass of carbon dioxide = 1.76 g · water = 0.72 g · calcium chloride = 4.44 g
Total mass of products = 1.76 + 0.72 + 4.44 = 6.92 g
Mass of reactants = Mass of products → the Law of Conservation of Mass is obeyed.
✏️ EXAMPLE 9.2 — METHOD, STUDY THIS Q: 12 g of carbon combines with 32 g of oxygen to form 44 g of carbon dioxide. If 2.4 g of carbon reacts completely with oxygen, how much carbon dioxide will be produced?
Carbon + Oxygen → Carbon dioxide
Solution: Given 12 g of carbon gives 44 g of carbon dioxide.
1 g of carbon will give = 44/12 g of carbon dioxide
2.4 g of carbon will give = (44/12) × 2.4 = 8.8 g of carbon dioxide
Hence, 8.8 g of carbon dioxide will be produced.
🧮 Numerical Practice (NCERT Pause and Ponder) — try first!
Q1 A student burns 10 g of ethanol in an open beaker. After the reaction, no residue is left in the beaker. Does this mean the Law of Conservation of Mass is violated? Explain.
Show Answer
No, the law is not violated.
Ethanol burns to form carbon dioxide and water vapour, which are gases that escape into the air from the open beaker.
If the reaction were carried out in a closed container and the mass of oxygen used plus the gaseous products were counted, the total mass before = total mass after.
Q2 When 20 g of hydrogen reacts completely with 160 g of oxygen, how much water is formed according to the Law of Conservation of Mass?
Show Answer
Mass of water = mass of reactants = 20 g + 160 g = 180 g of water.
📝 Check Your Concepts — Part 2
C1 State the Law of Conservation of Mass and name its proposer with the year.
Show Answer
Matter can neither be created nor destroyed in a chemical reaction. Proposed by Antoine Lavoisier in 1789.
C2 Why could Activity 9.3 be done in an open system while Activity 9.2 needed a closed one?
Show Answer
In Activity 9.3 no gas is formed (only a solid precipitate), so nothing can escape. In Activity 9.2 carbon dioxide gas is produced, which would escape from an open system and make the mass appear to decrease.
C3 Name the white precipitate formed in Activity 9.3 and write the word equation.
Show Answer
The white precipitate is barium sulfate. Sodium sulfate + Barium chloride → Barium sulfate + Sodium chloride
C4 5 g of A reacts with 8 g of B to form C and 3 g of D. Find the mass of C.
Show Answer
Total reactants = 5 + 8 = 13 g. Products must also total 13 g. Mass of C = 13 − 3 = 10 g.
🧠 Part 3 Map · Constant Proportions & Dalton
Proust’s Law
Fixed ratio by mass
Water always H : O = 1 : 8
Irrespective of source
Three Names, One Law
Law of Constant Proportions
Law of Definite Proportions
Proust’s Law
Dalton’s Postulates
Matter = tiny atoms
Atoms indivisible in reactions
Combine in whole number ratios
Why It Matters
Explains both laws logically
Atoms merely rearrange
PART 3 · CONSTANT PROPORTIONS & DALTON’S THEORY
9.2 Law of Constant Proportions
Soon after Lavoisier, Joseph Proust proposed that in any compound formed by two or more elements, the elements combine in a fixed ratio by mass.
In other words, elements in a compound have definite proportions irrespective of its source.
Example — water: water collected from rivers, borewells or the ocean, when purified and analysed, is always found to contain hydrogen and oxygen in a mass ratio of 1 : 8.
Meaning: if 9 g of purified water from any source is decomposed, 1 g of hydrogen and 8 g of oxygen are always obtained.
EXAM · DEFINITION (3 NAMES)
This is known as the Law of Constant Proportions, or the Law of Definite Proportions, or sometimes as Proust’s Law.
👨🔬 Meet a Scientist — Joseph Louis Proust
A prominent French chemist known for his careful experimental work.
Contributed to the Law of Definite Proportions by showing that chemical compounds always contain elements in fixed ratios by mass.
Studied the composition of copper carbonate — showing it always contains copper, carbon and oxygen in the same proportion by mass, no matter how it was prepared or where it was found.
🧵 Threads of Curiosity — Cinnabar (hingula)
In many ancient civilisations, a red pigment derived from rocks was widely used in painting and as a colouring agent.
In India it was known as hingula; in Latin and English, as cinnabar.
Heating cinnabar yields two elements — mercury and sulfur — in mass percentages of about 86.22% and 13.78% respectively.
Most civilisations also found that grinding mercury and sulfur together in this same ratio could form cinnabar — an early demonstration of constant proportions! (The toxic nature of both prevented the process from becoming widespread.)
✏️ EXAMPLE 9.3 — METHOD, STUDY THIS Q: Sodium chloride (NaCl) contains sodium and chlorine in the mass ratio of 23 : 35.5. If 46 g of sodium reacts completely, how much chlorine is needed to form NaCl?
Solution: Mass of chlorine required = (35.5 ÷ 23) × 46 = 71 g
🧮 Numerical Practice (NCERT Pause and Ponder) — try first!
Q3 A compound consists of 40% sulfur and 60% oxygen by mass. In a sample of the same compound containing 20 g of sulfur, what mass of oxygen must be present to satisfy the Law of Constant Proportions?
Show Answer
S : O = 40 : 60 = 2 : 3. If sulfur = 20 g, oxygen = (3/2) × 20 = 30 g of oxygen.
Q4 Carbon monoxide (CO) contains carbon and oxygen in the mass ratio of 3 : 4. How much oxygen will combine with 9 g of carbon to form carbon monoxide?
Show Answer
C : O = 3 : 4. For 9 g of carbon → oxygen = (4/3) × 9 = 12 g of oxygen.
Q5 The Law of Definite Proportions holds true for compounds but not for mixtures. Give reason.
Show Answer
In a compound, elements are chemically combined in a fixed ratio by mass, which never changes regardless of source or method of preparation.
In a mixture, components are simply physically mixed in any proportion — e.g., a sugar solution can be dilute or concentrated. Hence no fixed ratio exists, and the law does not apply.
Q6 Students X and Y both prepared an oxide of copper by combining copper and oxygen in the ratios of 4:1 and 8:2, respectively. Do their results justify the Law of Constant Proportions? Explain.
Show Answer
Yes. Simplify Y’s ratio: 8 : 2 = 4 : 1.
Both students obtained the same ratio of copper to oxygen by mass, even though they used different amounts.
This is exactly what the Law of Constant Proportions states — the compound has a definite composition irrespective of the quantity taken.
❓ What if …
atoms could combine in any ratio and not in a fixed ratio? How would this affect the substances around us?
9.3 Dalton’s Atomic Theory
The two laws above formed the basis of Dalton’s Atomic Theory.
The theory logically explains why substances combine in fixed proportions and why there is no loss or gain of mass during a chemical reaction.
Dalton explained these by proposing that during a chemical reaction, atoms are indivisible and merely rearrange, rather than being created or destroyed.
EXAM · DEFINITION
A postulate is a fundamental assumption accepted as truth without formal proof, from which further ideas are formed or developed.
John Dalton postulated that:
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 simple whole numbers to form compounds.
The relative number and kinds of atoms are constant in a given compound.
🔗 Dalton’s postulates provide the basis for the modern understanding of atoms and their behaviour.
Example: hydrogen and oxygen atoms combine to form water, but the atoms themselves are not destroyed or changed into something else.
Example: when magnesium burns in air, a white powder of magnesium oxide forms — showing atoms of magnesium have combined with those of oxygen.
👨🔬 Meet a Scientist — John Dalton
Born in England. In 1793 he moved to Manchester to teach mathematics, physics and chemistry at a college.
He spent most of his life teaching and researching there.
In 1808 he presented his atomic theory, which proved to be a turning point in the study of matter.
⏸▶ Pause and Ponder (NCERT)
Q7 Assertion (A): 2 g of hydrogen combines with 16 g of oxygen to form 18 g of water. Reason (R): According to Dalton’s Atomic Theory, atoms combine in a simple whole number ratio by mass to form compounds. Choose: (i) Both A and R are true, and R is the correct explanation of A. (ii) Both A and R are true, but R is not the correct explanation of A. (iii) A is true, but R is false. (iv) A is false, but R is true.
Show Answer
Option (i). A is true (2 g + 16 g = 18 g, and H : O = 1 : 8 by mass). R is true, and it correctly explains A — atoms combining in simple whole number ratios is exactly why hydrogen and oxygen combine in this fixed proportion to give 18 g of water.
📝 Check Your Concepts — Part 3
C1 State the Law of Constant Proportions with the water example.
Show Answer
In a compound, elements are always present in a fixed ratio by mass, irrespective of the source. Water from any source always contains hydrogen and oxygen in the ratio 1 : 8 — decomposing 9 g of water always gives 1 g H and 8 g O.
C2 Give the three names of Proust’s Law.
Show Answer
Law of Constant Proportions · Law of Definite Proportions · Proust’s Law.
C3 Which two Dalton postulates explain the Law of Conservation of Mass?
Show Answer
① Atoms are indivisible and ② atoms cannot be created or destroyed in a chemical reaction — they merely rearrange. So the total mass stays the same.
C4 Which Dalton postulate explains the Law of Constant Proportions?
Show Answer
Atoms combine in the ratio of simple whole numbers to form compounds, and the relative number and kinds of atoms are constant in a given compound.
C5 Define postulate.
Show Answer
A fundamental assumption accepted as truth without formal proof, from which further ideas are formed or developed.
🧠 Part 4 Map · Covalent Bonding
Two Ways to Combine
Sharing of electrons
Transfer of electrons
Molecules of Elements
H₂ — single bond
Cl₂ — single bond
O₂ — double bond
Molecules of Compounds
HCl — H shares with Cl
H₂O — two H share with O
Naming Rules
Prefixes mono-, di-, tri-…
Second element ends in -ide
PART 4 · HOW ATOMS COMBINE — COVALENT BOND
9.4 How Atoms Combine?
Atoms of an element can combine to form a molecule of that element — e.g., a hydrogen molecule consists of two hydrogen atoms.
Atoms of different elements combine to form a molecule of a compound — e.g., one hydrogen atom + one chlorine atom form a molecule of hydrogen chloride.
EXAM · DEFINITION
A molecule is an electrically neutral entity consisting of more than one atom that is capable of independent existence and shows all the properties of that substance.
Note: some elements, such as helium, exist only as atoms because their atoms are already stable.
From Chapter 8: atoms with 8 electrons in the outermost (valence) shell — or 2 electrons if the K-shell is outermost — are stable.
If the valence shell has fewer than eight electrons, atoms share, gain or lose electrons to complete their valence shell and become stable. In this process, atoms combine to form compounds. This takes place in two ways:
Sharing of electrons — share a few or all of their valence electrons with another atom.
Transfer of electrons — transfer one or more valence electrons to another atom, or accept one or more electrons from some other atom.
EXAM · DEFINITION
When atoms combine, the total energy of the system becomes lower than the sum of the energies of the individual atoms, making the arrangement more stable. The force that holds atoms together is called a chemical bond.
9.4.1 Bonding by sharing of electrons — Covalent Bond
A. Molecules of elements
Formation of a hydrogen molecule:
Electronic configuration of hydrogen (atomic number 1): only one electron in the K-shell.
Since the K-shell can hold a total of two electrons, it needs one more electron to become stable.
Thus a hydrogen atom shares one electron with another hydrogen atom to form a hydrogen molecule (H₂).
The shared pair of electrons attracts both the nuclei and makes the molecule stable.
Fig. 9.6: Formation of a hydrogen molecule
EXAM · DEFINITION
The interaction between atoms through a shared pair of electrons is called a covalent bond. If two atoms are joined by sharing one electron each, they are joined by a single bond, depicted by a single line: H—H.
Formation of a chlorine molecule:
The chlorine atom has seven electrons in its valence shell (2, 8, 7).
Each chlorine atom requires one electron to attain a stable electronic configuration.
A chlorine molecule forms by the sharing of one electron each by two chlorine atoms. This shared pair holds the two atoms together.
Single covalent bond → chlorine (Cl₂) is represented as Cl—Cl.
Fig. 9.7: Formation of a chlorine molecule
Formation of an oxygen molecule:
Electronic configuration of oxygen (atomic number 8): 2, 6 — the valence shell has six electrons and requires two more to complete its octet.
So two oxygen atoms share two electrons each, forming an oxygen molecule (O₂).
Fig. 9.8: Formation of an oxygen molecule
EXAM · DEFINITION
In O₂, the two atoms are joined by two pairs of shared electrons and are held together by a double bond, depicted by two lines: O=O.
👀 LOOK HERE: Count the shared dots between the atoms — H₂ and Cl₂ show one pair (single bond); O₂ shows two pairs (double bond). That dot count IS the bond type.
⏸▶ Pause and Ponder (NCERT)
Q8 Nitrogen has five valence electrons. Draw the structure of the nitrogen molecule (N₂).
Show Answer
Nitrogen (Z = 7) has configuration 2, 5 → needs 3 more electrons to complete its octet.
So two nitrogen atoms share three electrons each → three shared pairs.
This is a triple bond, written as N≡N.
Drawing: two circles overlapping, each nitrogen showing 2 non-shared electrons on the outer side, and 6 dots (3 pairs) in the overlap region.
Q9 The atomic number of fluorine is 9. Explain the formation of the fluorine molecule (F₂).
Show Answer
Fluorine (Z = 9) has configuration 2, 7 → needs one electron to complete its octet.
Two fluorine atoms each share one electron, forming one shared pair.
Both atoms now effectively have 8 valence electrons → stable. It is a single covalent bond: F—F.
🤝 Together we are strong … (NCERT)
Just as atoms share electrons to form covalent bonds, we too can share and care to build strong relationships with people around us. This sharing brings unity and stability, laying the foundation for a stronger community, and ultimately, a strong nation.
B. Molecules of compounds
Formation of hydrogen chloride (HCl): follow these steps —
Write the electronic configurations of hydrogen and chlorine (H = 1; Cl = 17 → 2, 8, 7).
Calculate the electrons required by chlorine to complete its octet and by hydrogen to complete its duplet.
Both hydrogen and chlorine need one electron each → both share one electron each to form a molecule of hydrogen chloride.
Fig. 9.9: Formation of a hydrogen chloride molecule
Since H and Cl bond by sharing electrons, HCl is a covalent compound.
They share one pair of electrons → bonded by a single bond, depicted as H—Cl.
Formation of water (H₂O):
Hydrogen needs only one electron; oxygen needs two electrons for stable configurations.
Solution: two hydrogen atoms each share an electron with one oxygen atom.
Fig. 9.10: Formation of a water molecule
💧 The water molecule is represented as H₂O — indicating two hydrogen atoms and one oxygen atom.
⏸▶ Pause and Ponder (NCERT)
Q10 Show the formation of the following molecules: (i) Carbon dioxide (CO₂) (ii) Hydrogen sulfide (H₂S) (iii) Ammonia (NH₃)
Show Answer
(i) CO₂: Carbon (2, 4) needs 4 electrons; each oxygen (2, 6) needs 2. Carbon shares two electrons with each oxygen → two double bonds: O=C=O.
(ii) H₂S: Sulfur (2, 8, 6) needs 2 electrons; each hydrogen needs 1. Two hydrogen atoms share one electron each with sulfur → two single bonds: H—S—H.
(iii) NH₃: Nitrogen (2, 5) needs 3 electrons; each hydrogen needs 1. Three hydrogen atoms share one electron each with nitrogen → three single bonds. Nitrogen also keeps one non-bonded pair.
Q11 Neon (atomic number 10) neither transfers nor shares its valence electrons. Explain.
Show Answer
Neon’s configuration is 2, 8 — its outermost shell already has a complete octet. It is therefore stable and unreactive, with no need to lose, gain or share electrons. Its valency is 0.
🇮🇳 Bridging Science and Society (NCERT)
Atoms can release enormous energy when their nuclei split or combine to form new elements — called atomic or nuclear energy.
Beyond electricity generation, it is used in medicine, scientific research and space exploration.
In nuclear power plants, thermal energy from nuclear reactions produces steam that drives turbines and generates electricity — a cleaner alternative to fossil fuels.
In India, scientists like Raja Ramanna (often called the Father of the Indian Nuclear Programme) made significant contributions in developing the nation’s Nuclear Energy Programme and promoting its peaceful use for development.
C. Naming covalent compounds
Covalent compounds are named by indicating the number of atoms of each element using a prefix system.
The first element retains its regular name; the second element ends in -ide.
mono- is usually omitted for the first element but used for the second element.
If a prefix ends with ‘o’ or ‘a’ and the element starts with a vowel, drop the last vowel (e.g., monoxide, pentoxide).
If the prefix ends with ‘i’, keep it for pronunciation (e.g., dioxide, trioxide).
✏️ EXAMPLES — STUDY THIS
CO → carbon monoxide (not monooxide)
CO₂ → carbon dioxide (not monocarbon dioxide)
CS₂ → carbon disulfide, showing two sulfur atoms
PCl₃ → phosphorus trichloride, showing three chlorine atoms
SF₆ → sulfur hexafluoride, showing six fluorine atoms
N₂O₄ → dinitrogen tetroxide (not tetraoxide)
N₂O₅ → dinitrogen pentoxide
When hydrogen is the first element in the formula, no prefix is added before hydrogen, irrespective of the number of its atoms. Example: H₂S is hydrogen sulfide, not dihydrogen sulfide.
A few binary compounds are known only by their common names: H₂O (would be hydrogen monoxide) is commonly known as water; NH₃ (actually nitrogen trihydride) is known as ammonia.
📝 Check Your Concepts — Part 4
C1 Define a molecule and a chemical bond.
Show Answer
Molecule: an electrically neutral entity of more than one atom, capable of independent existence, showing all the properties of that substance. Chemical bond: the force that holds atoms together when they combine.
C2 Why does helium exist only as single atoms?
Show Answer
Helium’s K-shell already holds 2 electrons — a complete outermost shell. Its atoms are already stable, so they have no need to bond with other atoms.
C3 Distinguish between a single bond and a double bond, with an example each.
Show Answer
Single bond:one shared pair of electrons — e.g., H—H, Cl—Cl, H—Cl. Double bond:two shared pairs — e.g., O=O.
C5 Why is H₂S called hydrogen sulfide and not dihydrogen sulfide?
Show Answer
Because when hydrogen is the first element in a formula, no prefix is used before hydrogen, no matter how many hydrogen atoms are present.
🧠 Part 5 Map · Ionic Bonding
Losing Electrons
< 4 valence e⁻ → donates
Na → Na⁺ + e⁻
Positive = cation
Gaining Electrons
> 4 valence e⁻ → accepts
Cl + e⁻ → Cl⁻
Negative = anion
Ionic Bond
Electrostatic force of attraction
Between opposite charges
Na⁺ + Cl⁻ → NaCl
Naming
Cation first, then anion
Simple anions end in -ide
Polyatomic ions: exceptions
PART 5 · IONIC BOND
9.4.2 Bonding by electron transfer — Ionic bond
If the valence shell of an atom has less than four electrons, it generally donates its valence electrons to achieve a stable electronic configuration.
Atoms with more than 4 valence electrons usually gain or share electrons to complete an octet.
Example: sodium chloride (common salt), chemical formula NaCl.
Formation of the sodium cation (Na⁺)
Atomic number of sodium is 11 → configuration 2, 8, 1. Its valence shell contains only one electron.
It attains a stable configuration after losing this valence electron.
After losing one electron it has 11 protons and 10 electrons → it becomes a positively charged species, called a sodium cation, represented as Na⁺.
Fig. 9.11: Formation of a sodium cation
Formation of the chloride anion (Cl⁻)
Chlorine has seven valence electrons (2, 8, 7).
It attains a stable configuration by gaining one electron from another atom.
After gaining one extra electron it acquires a negative charge and is called a chloride anion, represented as Cl⁻.
Fig. 9.12: Formation of a chloride anion
👀 LOOK HERE: Note the square brackets and the charge sign in Figs. 9.11 and 9.12. Brackets mean “this is an ION, not an atom” — always draw them in the exam.
EXAM · DEFINITION
Cations and anions are collectively called ions. Once sodium and chloride ions form, they are held together by the electrostatic force of attraction due to their opposite charges.
The electrostatic force of attraction between oppositely charged ions that holds them together is called an ionic bond.
Na· + :Cl: → Na⁺ + :Cl:⁻ → NaCl
Fig. 9.13: Formation of sodium chloride by transfer of electron
🧵 Threads of Curiosity — Crystals of ionic compounds
Ionic compounds usually do not remain as single units. They form three-dimensional (3-D) crystals in which ions are arranged in a repeating pattern.
In sodium chloride, each Na⁺ is surrounded by six Cl⁻ ions, and each Cl⁻ is surrounded by six Na⁺ ions.
This regular, repeating 3-D pattern is known as a crystal structure.
The crystal structure is represented as a crystal lattice, with ions depicted as points or dots — it helps visualise the arrangement of ions in the crystal.
Some elements like sulfur have six electrons in their outer shell and need to gain two electrons to complete their octet. On gaining two electrons, sulfur acquires two units of negative charge and is represented as S²⁻.
❓ What if …
we could see atoms directly? How would it help scientists and what challenges would it cause?
⏸▶ Pause and Ponder (NCERT)
Q12 What kind of ion will oxygen (O) form?
Show Answer
Oxygen (2, 6) has 6 valence electrons and needs 2 more → it gains 2 electrons → forms an anion with 2 units of negative charge: the oxide ion, O²⁻.
Q13 Fill in the blanks: Among magnesium and chlorine, magnesium atom can give two electrons to become Mg²⁺. However, chlorine can take only one electron to become ______. Now, ______ ion of magnesium and ______ ions of chlorine combine to give magnesium chloride.
Show Answer
Chlorine becomes Cl⁻ (chloride ion). Now, one ion of magnesium and two ions of chlorine combine to give magnesium chloride (MgCl₂).
Q14 Show the formation of cations of potassium (K) and calcium (Ca) atoms, and the formation of their corresponding chlorides using diagrams.
Show Answer
Potassium (Z = 19): 2, 8, 8, 1 → loses 1 electron → K⁺ (2, 8, 8). It gives that electron to one chlorine atom → Cl⁻. Formula: KCl.
Calcium (Z = 20): 2, 8, 8, 2 → loses 2 electrons → Ca²⁺ (2, 8, 8). Those 2 electrons go to two chlorine atoms → 2 Cl⁻. Formula: CaCl₂.
Drawing: show the metal atom’s shells before and after (in square brackets with + or 2+), and each chlorine as 2, 8, 8 in brackets with a − sign.
Q15 Illustrate how sodium sulfide (Na₂S) is formed.
So two sodium atoms each donate one electron to one sulfur atom.
Two Na⁺ ions and one S²⁻ ion are held by an ionic bond → formula Na₂S.
A. Naming ionic compounds
The name of the cation is written first, followed by the name of the anion.
Names of simple anions end with -ide.
Generally, metals form cations and non-metals form anions.
Ionic compounds are typically formed when metals combine with non-metals — e.g., sodium chloride, calcium oxide, magnesium sulfide.
Some ions are formed by the combination of atoms of two or more elements — these are called polyatomic ions. Names of polyatomic ions generally do not end with -ide.
Table 9.1: (a) Some common monoatomic ions
Name of ion
Formula
Valency
Sodium
Na⁺
1
Lithium
Li⁺
1
Potassium
K⁺
1
Silver
Ag⁺
1
Calcium
Ca²⁺
2
Barium
Ba²⁺
2
Iron (Ferrous)
Fe²⁺
2
Iron (Ferric)
Fe³⁺
3
Copper (Cuprous)
Cu⁺
1
Copper (Cupric)
Cu²⁺
2
Magnesium
Mg²⁺
2
Zinc
Zn²⁺
2
Aluminium
Al³⁺
3
Fluoride
F⁻
1
Chloride
Cl⁻
1
Bromide
Br⁻
1
Iodide
I⁻
1
Oxide
O²⁻
2
Sulfide
S²⁻
2
Table 9.1: (b) Some common polyatomic ions
Name of ion
Formula
Valency
Hydroxide
OH⁻
1
Nitrate
NO₃⁻
1
Hydrogencarbonate
HCO₃⁻
1
Carbonate
CO₃²⁻
2
Sulfate
SO₄²⁻
2
Ammonium
NH₄⁺
1
⭐ Learn Table 9.1 by heart — every formula-writing question in the exercises depends on knowing these valencies.
📝 Check Your Concepts — Part 5
C1 Define cation, anion and ionic bond.
Show Answer
Cation: a positively charged ion formed when an atom loses electrons. Anion: a negatively charged ion formed when an atom gains electrons. Ionic bond: the electrostatic force of attraction between oppositely charged ions that holds them together.
C2 Why does Na⁺ carry a positive charge although sodium atom is neutral?
Show Answer
The sodium atom has 11 protons and 11 electrons (neutral). After losing one electron it has 11 protons but only 10 electrons — one extra positive charge remains unbalanced → Na⁺.
C3 In a NaCl crystal, how many Cl⁻ ions surround each Na⁺ ion?
Show Answer
Six chloride ions surround each sodium ion, and six sodium ions surround each chloride ion, in a repeating 3-D crystal structure.
C4 Write the formula and valency of: ferric ion, cuprous ion, sulfate ion, ammonium ion.
An ion formed by the combination of atoms of two or more elements, carrying an overall charge. Examples: carbonate (CO₃²⁻) and ammonium (NH₄⁺). Their names generally do not end in -ide.
🧠 Part 6 Map · Writing Chemical Formulae
Covalent — 3 Steps
Write symbols
Write valencies
Criss-cross as subscripts
Ionic — 4 Steps
Cation first, then anion
Charges under symbols
Criss-cross numbers only
Divide by common factor
Golden Rules
Valency 1 is not written
Mg₂O₂ → simplified to MgO
Charges never shown in formula
Brackets Rule
2 or more same polyatomic ions
Al(OH)₃ not AlOH₃
One ion → no bracket
PART 6 · WRITING CHEMICAL FORMULAE
9.5 Writing Chemical Formulae
You have learnt to write formulae by finding the number of electrons shared or transferred. There is a quicker way — the criss-cross method.
9.5.1 Writing chemical formulae of covalent compounds
(i) Write the symbols of the constituent elements of the compound.
(ii) Write the valencies of these elements (refer to Table 9.1).
(iii) Crossover the valencies of the combining atoms and write them as subscripts after the symbols of elements.
✏️ EXAMPLE — HYDROGEN CHLORIDE
Symbol of element: HCl
Valency: 1 ✕ 1
Criss-cross → H₁Cl₁ → the formula of the compound would be HCl. Rule: if the valency is one after criss-crossing, it is not written.
✏️ EXAMPLE — HYDROGEN SULFIDE
Symbol of element: HS
Valency: 1 ✕ 2
Criss-cross → the formula of the compound would be H₂S.
✏️ EXAMPLE — CARBON TETRACHLORIDE
Symbol of element: CCl
Valency: 4 ✕ 1
Criss-cross → the formula of the compound would be CCl₄.
9.5.2 Writing chemical formulae of ionic compounds
(i) Write the symbol of the cation first, followed by the symbol of the anion.
(ii) Write the charges under the symbols rather than as superscripts.
(iii) Crossover the charges (only the numbers) to obtain the formula.
(iv) The chemical formula gives the simplest ratio of the elements. After criss-crossing, the subscripts are divided by a common factor, if any. Example: subscripts 2 and 4 are divided by 2 → 1 and 2.
📌 Note (NCERT)
The charges on the ions are NOT indicated in the formula of the compound.
Thus, in calcium chloride there are two chloride ions (Cl⁻) for each calcium ion (Ca²⁺). The positive and negative charges must balance each other, and the overall structure must be neutral.
✏️ EXAMPLE — MAGNESIUM OXIDE (simplifying)
Symbol: MgO
Charge: 2+ ✕ 2−
Criss-cross gives Mg₂O₂ — but here the valencies of the two elements are the same, so it is simply written as MgO.
This method can also be used to write formulae of compounds of metals with polyatomic ions, such as calcium carbonate.
✏️ EXAMPLE — CALCIUM CARBONATE
Symbol: CaCO₃
Charge: 2+ ✕ 2−
Formula: CaCO₃
Here the valencies of the two ions are the same, so Ca₂(CO₃)₂ is simply written as CaCO₃.
Thus in magnesium hydroxide there are two hydroxide ions (OH⁻) for each magnesium ion (Mg²⁺).
EXAM · BRACKETS RULE
Use brackets ( ) when there are two or more polyatomic ions of the same type in a formula.
Example: in aluminium hydroxide, the bracket around OH with subscript 3 indicates three hydroxide ions bound to one aluminium ion. Brackets are NOT required when only one polyatomic anion is present.
✏️ EXAMPLE — ALUMINIUM HYDROXIDE
Symbol: AlOH
Charge: 3+ ✕ 1−
Formula: Al(OH)₃
⚠️ Note: the formula is Al(OH)₃, NOT AlOH₃.
👀 LOOK HERE — the three traps students fall into: ① valency 1 is never written as a subscript ② same valencies must be simplified (Mg₂O₂ → MgO) ③ two or more identical polyatomic ions ALWAYS take brackets.
🧮 Numerical Practice (NCERT Pause and Ponder) — try first!
Q16 Name the following: (i) CO₂ (ii) NO₂ (iii) SF₆ (iv) PCl₃
Q17 Write the formula for: (i) Sodium hydrogencarbonate (ii) Sulfur dioxide (iii) Ferric chloride (iv) Cuprous oxide
Show Answer
(i) Na⁺ + HCO₃⁻ → NaHCO₃
(ii) Sulfur dioxide → SO₂
(iii) Fe³⁺ + Cl⁻ → FeCl₃
(iv) Cu⁺ + O²⁻ → Cu₂O
Q18 Write the formulae for the compounds formed from the following pairs of ions: (i) Fe³⁺ and OH⁻ (ii) K⁺ and CO₃²⁻
Show Answer
(i) Fe³⁺ and OH⁻ → criss-cross → Fe(OH)₃ (brackets needed — three hydroxide ions)
(ii) K⁺ and CO₃²⁻ → criss-cross → K₂CO₃ (no brackets — only one carbonate ion)
📝 Check Your Concepts — Part 6
C1 List the four steps for writing the formula of an ionic compound.
Show Answer
① Cation symbol first, then anion ② write charges under the symbols ③ criss-cross the numbers only ④ divide subscripts by a common factor if any, to give the simplest ratio.
C2 Why is magnesium oxide written MgO and not Mg₂O₂?
Show Answer
Because the chemical formula must show the simplest whole number ratio. Both ions have valency 2, so the subscripts 2 and 2 are divided by the common factor 2 → MgO.
C3 When are brackets used in a formula? Give one example with and one without.
Show Answer
Brackets are used when two or more polyatomic ions of the same type are present. With brackets: Al(OH)₃, Mg(OH)₂, Al₂(SO₄)₃. Without: CaCO₃, NaHCO₃, K₂CO₃ (only one polyatomic ion).
PART 7 · PROPERTIES · MOLECULAR MASS · FORMULA UNIT MASS
9.6 Properties of the Ionic and the Covalent Compounds
Activity 9.4: Let us experiment
Collect samples of some compounds, such as camphor, sodium chloride, copper sulfate, sugar and naphthalene.
(A) Solubility in (i) water, (ii) kerosene, and (iii) petrol
Try dissolving each sample separately in water, kerosene and petrol.
Record your observations in Table 9.2.
(B) Electrical conductivity in the water
⚠️ Safety first (NCERT)
Do not touch the electrodes when they are connected to the battery — use a low-voltage battery to avoid the risk of shock. Petrol and kerosene are flammable liquids, so be careful while working with them.
Take two carbon or metal electrodes and insert them into a piece of cardboard by drilling two holes.
Connect one end of each electrode to the terminals of the 9 V battery and to the light bulb.
Test the electrical conductivity of each solid sample and observe whether the bulb glows. Record your observations.
Test the electrical conductivity of each sample dissolved in water by transferring the solution to the beaker one at a time and observing whether the bulb glows.
Record your observations with other samples given in Table 9.2.
Group the compounds showing similar properties listed in Table 9.2.
Fig. 9.15: Experimental set-up for electrical conductivity of a solution
Table 9.2: Solubility and electrical conductivity observations (fill in your own observations)
Compound
Solubility in
Electrical conductivity of compounds in
water
kerosene
petrol
solid state
water
Camphor
Insoluble
Soluble
Soluble
No
No
Sodium chloride
Soluble
Insoluble
Insoluble
No
Yes
Copper sulfate
Soluble
Insoluble
Insoluble
No
Yes
Sugar
Soluble
Insoluble
Insoluble
No
No
Naphthalene
Insoluble
Soluble
Soluble
No
No
Conclusions from Activity 9.4
Ionic compounds like sodium chloride and copper sulfate are generally soluble in water but insoluble in solvents such as kerosene and petrol.
Most covalent compounds, such as camphor and naphthalene, are insoluble in water but dissolve in kerosene and petrol.
Ionic compounds do not conduct electricity in the solid state because their ions are held in fixed positions by strong forces.
To conduct electricity, ions must be free to move — which occurs only when ionic compounds are dissolved in water.
Some covalent compounds, such as sugar, are soluble in water but do not provide ions in solution — therefore they do not conduct electricity.
Other covalent compounds, such as camphor and naphthalene, also do not conduct electricity.
Ionic compounds generally have high melting and boiling points due to strong inter-ionic attractions, whereas covalent compounds usually have low melting and boiling points.
Q (NCERT in-text) Predict whether ionic and covalent compounds would conduct electricity in the molten state (the melted state of a substance).
Show Answer
Ionic compounds: YES. On melting, the strong forces holding the ions in fixed lattice positions are overcome — the ions become free to move and carry current.
Covalent compounds: NO. They contain no ions at all, so even in the molten state there are no charged particles free to carry current.
⏸▶ Pause and Ponder (NCERT)
Q19 What type of chemical bond is present in a solid compound that does not conduct electricity in the solid state but conducts electricity when dissolved in water?
Show Answer
An ionic bond. In the solid state its ions are locked in fixed positions; on dissolving in water the ions become free to move and conduct electricity.
Q20 Metal M, with two electrons in its valence shell (M shell), reacts with oxygen to form a compound that is slightly soluble in water. Predict its: (i) formula (ii) type of bond (iii) electrical conductivity of its aqueous solution.
Show Answer
The metal has 2 electrons in the M shell → configuration 2, 8, 2 → it is magnesium (Mg), valency 2. Oxygen valency 2.
(i) Formula: MO (i.e., MgO — criss-cross gives M₂O₂, simplified to MO).
(ii) Type of bond: ionic (metal + non-metal, formed by transfer of two electrons).
(iii) Conductivity: its aqueous solution conducts electricity, because the dissolved portion provides free-moving ions (though it is only slightly soluble, so conductivity is weak).
9.7 Molecular Mass of Covalent Compounds
Since you know the formulae of covalent compounds, you can find the masses of their molecules by simply adding up the masses of the atoms present in them.
✏️ EXAMPLE 9.4 — METHOD, STUDY THIS Molecular mass of water (H₂O)
Atomic mass — H = 1 u; O = 16 u
Molecular mass of H₂O = (1 u × 2) + (16 u × 1) = 18 u
✏️ EXAMPLE 9.5 — METHOD, STUDY THIS Molecular mass of carbon dioxide (CO₂)
Atomic mass — C = 12 u; O = 16 u
Molecular mass of CO₂ = (12 u × 1) + (16 u × 2) = 44 u
⚠️ Remember: in ionic compounds the ions form 3-D crystals — i.e., ionic compounds do NOT form molecules. That is why they need a different term ⬇
9.8 Formula Unit Mass of Ionic Compounds
EXAM · DEFINITION
In ionic compounds, the collection of the simplest whole number ratio of ions is termed a formula unit. The mass of a formula unit is called the formula unit mass.
✏️ EXAMPLE 9.6 — METHOD, STUDY THIS Formula unit mass of sodium oxide (Na₂O)
Atomic mass — Na = 23 u; O = 16 u
Formula unit mass of Na₂O = (23 u × 2) + (16 u × 1) = 62 u
✏️ EXAMPLE 9.7 — METHOD, STUDY THIS (with brackets!) Formula unit mass of calcium nitrate, Ca(NO₃)₂
Atomic mass — Ca = 40 u; N = 14 u; O = 16 u
Formula unit mass = (40 u × 1) + {(14 u × 1) + (16 u × 3)} × 2
= 40 + {14 + 48} × 2 = 40 + 124 = 164 u
👉 Note how the bracket contents are calculated first, then multiplied by the subscript outside.
🧮 Numerical Practice (NCERT Pause and Ponder) — try first!
Q21 Find the molecular mass of nitric acid (HNO₃). Atomic mass — H = 1 u; N = 14 u; O = 16 u.
Q22 Find the molecular mass of methane (CH₄). Atomic mass — C = 12 u; H = 1 u.
Show Answer
CH₄ = (12 × 1) + (1 × 4) = 12 + 4 = 16 u
Q23 Find the formula unit mass of potassium chloride (KCl). Atomic mass — K = 39 u; Cl = 35.5 u.
Show Answer
KCl = 39 + 35.5 = 74.5 u
Q24 Find the formula unit mass of magnesium hydroxide, Mg(OH)₂. Atomic mass — Mg = 24 u; O = 16 u; H = 1 u.
Show Answer
Mg(OH)₂ = 24 + {16 + 1} × 2 = 24 + 34 = 58 u
🌍 Understanding atoms, molecules and chemical bonding reveals the hidden science behind everything we use and consume each day — it shows how tiny atoms combine to build the universe we live in!
At a Glance (NCERT Summary)
Mass can neither be created nor destroyed in a chemical reaction. This is known as the Law of Conservation of Mass.
A compound always contains the same elements combined in a fixed ratio by mass, no matter how it is formed or from where it is obtained. This is called the Law of Definite Proportions.
A molecule is defined as an electrically neutral entity consisting of more than one atom that can exist independently and shows all its chemical properties.
Atoms combine to form molecules of elements or compounds to become stable. Atoms are held together by a force called a chemical bond.
A covalent bond is formed by the sharing of electrons between atoms.
An ionic bond is formed by the transfer of electrons between atoms, where one atom loses electrons and the other gains electrons to form cations and anions, respectively.
The chemical formula of a covalent compound represents the elements and number of atoms of each element present in it.
The chemical formula of an ionic compound represents the simplest whole number ratio of atoms of different elements present in it.
Molecular mass is the total mass of a molecule, calculated by adding the atomic masses of all the atoms constituting it.
Formula unit mass of an ionic compound is the sum of the atomic masses of all the atoms present in a formula unit (simplest whole number ratio of ions in an ionic compound).
📝 Check Your Concepts — Part 7
C1 Why does solid NaCl not conduct electricity while its solution does?
Show Answer
In the solid, ions are held in fixed positions by strong electrostatic forces and cannot move. In solution, the ions become free to move and carry the current.
C2 Sugar dissolves in water but its solution does not conduct electricity. Why?
Show Answer
Sugar is a covalent compound — it dissolves as neutral molecules and does not provide ions in solution. With no free-moving charged particles, no current flows.
C3 Why do ionic compounds have high melting points?
Show Answer
Because of strong inter-ionic attractions between oppositely charged ions in the crystal lattice — a large amount of heat energy is needed to break them. Covalent compounds have weaker forces, so low melting and boiling points.
C4 Why do we say “formula unit mass” for NaCl instead of “molecular mass”?
Show Answer
Because ionic compounds like NaCl exist as 3-D crystals of ions, not as separate molecules. NaCl only shows the simplest whole number ratio of ions (a formula unit), so we call its mass the formula unit mass.
C5 Calculate the molecular mass of H₂SO₄ (H = 1, S = 32, O = 16).
Q1 A particular element (A) has one electron in its third shell. There is another element (B) with six electrons in its second shell. (i) How many electrons does A tend to give or take to become stable? (ii) What kind of ion would it form? (iii) How many electrons does B tend to give or take to become stable? (iv) What kind of ion would it form? (v) If A and B were to combine, what kind of bond would be formed? (vi) What would be the formula for the compound thus formed?
Show Answer
A has one electron in the third shell → configuration 2, 8, 1 → sodium (Na). B has six electrons in the second shell → 2, 6 → oxygen (O).
(i) A tends to give 1 electron.
(ii) It forms a cation, Na⁺.
(iii) B tends to take 2 electrons.
(iv) It forms an anion, O²⁻.
(v) An ionic bond (electrons transferred from metal to non-metal).
(vi) Criss-cross Na⁺ and O²⁻ → Na₂O.
Q2 An element X has six electrons in its outer shell and forms a diatomic molecule. (i) Why would that be so? (ii) What kind of bond would it form? (iii) Draw the structure of the molecule it would form. (iv) A certain other element Y has two electrons in its second shell. Draw the structure of the molecule that X would form with Y.
Show Answer
X has 6 valence electrons → oxygen (2, 6). Y has 2 electrons in the second shell → beryllium (2, 2).
(i) X needs 2 more electrons to complete its octet. Two X atoms each share two electrons, so both attain 8 → they exist as a stable diatomic molecule (X₂).
(ii) A covalent bond — specifically a double bond (two shared pairs).
(iii) Structure: two overlapping circles, each atom showing 4 non-shared electrons, with 4 dots (2 pairs) in the overlap → O=O.
(iv) Y (Be, 2 valence electrons) shares its 2 electrons with X → each Be–O link is a shared pair. Structure: Be in the centre sharing electrons with O, drawn as Be=O (or in the ionic view Be²⁺ and O²⁻ giving BeO).
Q3 You want to design a new ionic compound, where the total positive charge is 6+ and the total negative charge is 6−. Which of the following combinations gives the correct number of ions? (i) 2 Al³⁺ and 3 Cl⁻ (ii) 3 Mg²⁺ and 1 PO₄³⁻ (iii) 2 Fe³⁺ and 3 O²⁻ (iv) 3 Ca²⁺ and 2 SO₄²⁻
Show Answer
Check each: positive total vs negative total —
(i) 2 × 3+ = 6+ but 3 × 1− = 3− ❌ (not balanced)
(ii) 3 × 2+ = 6+ but 1 × 3− = 3− ❌
(iii) 2 × 3+ = 6+ and 3 × 2− = 6− ✅
(iv) 3 × 2+ = 6+ but 2 × 2− = 4− ❌
Correct answer: (iii) 2 Fe³⁺ and 3 O²⁻ → formula Fe₂O₃.
Q4 Choose the correct statement(s) and correct the false statement(s). (i) Elements are made up of molecules and compounds are made up of atoms. (ii) The molecule of a compound is always made up of two or more atoms of the same kind. (iii) One molecule of nitrogen gas contains three nitrogen atoms. (iv) Water is made of two hydrogen atoms, covalently bonded with one oxygen atom.
Show Answer
(i) ❌ False. Correction: elements are made up of atoms (which may exist as molecules of the same kind), and compounds are made up of molecules containing atoms of different elements.
(ii) ❌ False. Correction: the molecule of a compound is made up of two or more atoms of different kinds.
(iii) ❌ False. Correction: one molecule of nitrogen gas contains two nitrogen atoms (N₂).
(iv) ✅ Correct. Water (H₂O) has two hydrogen atoms covalently bonded to one oxygen atom.
Q5 Write the chemical formulae for: (i) Aluminium nitrate (ii) Calcium oxide (iii) Ferric oxide
The other options show the wrong total number of electrons (they represent the neutral atom or incorrect arrangements), so they cannot be the chloride ion.
Tip: always count the outermost shell first — for any negative ion of a Group 17 element, the outer shell must be complete with 8.
Q8 Determine the formula unit mass of: (i) Ammonium nitrate (NH₄NO₃) (ii) Phosphoric acid (H₃PO₄) (iii) Sodium hydrogencarbonate (NaHCO₃)
(Atomic masses used: H = 1, C = 12, N = 14, O = 16, Na = 23, P = 31.)
Q9 Write the formulae for the compounds formed by the reaction of: (i) Magnesium and nitrogen (ii) Lithium and nitrogen (iii) Sodium and sulfur (iv) Aluminium and oxygen
Show Answer
(i) Mg²⁺ and N³⁻ → Mg₃N₂ (magnesium nitride)
(ii) Li⁺ and N³⁻ → Li₃N (lithium nitride)
(iii) Na⁺ and S²⁻ → Na₂S (sodium sulfide)
(iv) Al³⁺ and O²⁻ → Al₂O₃ (aluminium oxide)
Q10 Complete Table 9.3 by writing the formulae of the compounds formed by the cations on the left and the anions at the top. LiNO₃ is given as an example.
Show Answer
Table 9.3 — completed
NO₃⁻
SO₄²⁻
PO₄³⁻
NH₄⁺
NH₄NO₃
(NH₄)₂SO₄
(NH₄)₃PO₄
Li⁺
LiNO₃
Li₂SO₄
Li₃PO₄
Al³⁺
Al(NO₃)₃
Al₂(SO₄)₃
AlPO₄
Cu²⁺
Cu(NO₃)₂
CuSO₄
Cu₃(PO₄)₂
Note where brackets appear — only when two or more of the same polyatomic ion are present.
Q11 5.3 g of sodium carbonate and 6.0 g of acetic acid react to produce 2.2 g of carbon dioxide, 0.9 g of water, and 8.2 g of sodium acetate. Verify whether the law of conservation of mass is valid.
Show Answer
Total mass of reactants = 5.3 g + 6.0 g = 11.3 g
Total mass of products = 2.2 g + 0.9 g + 8.2 g = 11.3 g
Mass of reactants = Mass of products → the Law of Conservation of Mass is valid.
Q12 If a species has 11 protons, 12 neutrons and 10 electrons then (i) what is its atomic number and mass number? (ii) is it neutral, a cation or an anion? Explain. (iii) write its electronic configuration. (iv) name the species.
Show Answer
(i) Atomic number = protons = 11; Mass number = 11 + 12 = 23.
(ii) It is a cation — it has 11 protons but only 10 electrons, so there is one extra positive charge (charge = +1).
(iii) With 10 electrons: 2, 8.
(iv) The species is the sodium ion, Na⁺.
Q13 Two elements, A and B, have the configurations — A: 2, 8, 5 B: 2, 8, 7. (i) Which element is more reactive? (ii) Will A and B form ionic or covalent bonds when they combine? Explain using electron transfer or sharing. (iii) Predict the formula of the compound they would form.
Show Answer
A (2, 8, 5) is phosphorus; B (2, 8, 7) is chlorine.
(i) B is more reactive — it needs only one electron to complete its octet, whereas A needs three. The closer an atom is to a complete octet, the more readily it reacts.
(ii) Both are non-metals and both need to gain electrons — neither can donate easily. So they achieve stability by sharing electrons → a covalent bond. A shares one electron with each of three B atoms.
(iii) Valency of A = 3, valency of B = 1 → formula AB₃ (i.e., PCl₃).
Q14 Assertion (A): Copper sulfate conducts electricity in the molten state but not in the solid state. Reason (R): Copper and sulfate ions are fixed in the lattice in molten state, while in solid state they can move freely. Choose: (i) Both A and R are true, and R is the correct explanation of A. (ii) Both A and R are true, but R is not the correct explanation of A. (iii) A is true, but R is false. (iv) A is false, but R is true.
Show Answer
Option (iii) — A is true, but R is false.
A is true: molten copper sulfate conducts; solid copper sulfate does not.
R is false because it states the situation backwards. The correct statement: ions are fixed in the lattice in the SOLID state, and move freely in the MOLTEN state.
Q15 The species ²⁷Al, ⁸⁰Br⁻ and ²⁰¹Hg²⁺ have 13, 35 and 80 protons, respectively. How many electrons and neutrons do they have?
Show Answer
Species
Protons
Electrons
Neutrons (A − p)
²⁷Al (neutral)
13
13
27 − 13 = 14
⁸⁰Br⁻ (gained 1 e⁻)
35
36
80 − 35 = 45
²⁰¹Hg²⁺ (lost 2 e⁻)
80
78
201 − 80 = 121
Rule: negative charge → add that many electrons; positive charge → subtract that many electrons. Neutrons are unaffected by charge.
The Journey Beyond (NCERT Activities)
Design and perform an experiment to show and compare that water always contains hydrogen and oxygen in the same ratio, regardless of its source.
Compare atoms and ions of any three elements. Show the number of electrons before and after ion formation using bar graphs.
Make a card game with cations and anions. Possible ideas may include — picking cards from the pile or from open cards and matching them with cards in your hand to form compounds (you may discard any unwanted cards); and asking other players for cards to form compounds using the cards they already have.
To learn more about molecules you can explore: phet.colorado.edu/sims/html/build-a-molecule/latest/build-a-molecule_all.html
🗝 The Quest Continues …
Are there any chemical changes that do not obey the Law of Conservation of Mass?
