Atomic Foundations of Matter — Complete Animated Notes

⚖️ Laws of Chemical Combination

Before understanding atoms and molecules, scientists studied how elements combine during chemical reactions. Two fundamental laws govern these combinations. These laws were established through careful experiments and laid the groundwork for Dalton's atomic theory.

🔬 Key Scientists

Antoine Lavoisier (1774) -- Father of Modern Chemistry -- Law of Conservation of Mass
Joseph Proust (1799) -- French chemist -- Law of Constant Proportions

🎓 Why These Laws Matter

These laws provided the experimental evidence that Dalton used to propose his atomic theory. They are the foundation of all quantitative chemistry!

⚖️ Law of Conservation of Mass (Lavoisier, 1774)

Antoine Lavoisier established that mass is neither created nor destroyed in a chemical reaction. The total mass of the reactants equals the total mass of the products.

Reactants

Products

⚖️ Mass Before = Mass After

Reactants

Products

→

⚖️ Same atoms, rearranged — MASS CONSERVED!

NCERT Activity 9.1

🔬 BaCl₂ + Na₂SO₄ Reaction

BaCl₂ + Na₂SO₄ → BaSO₄↓ + 2NaCl

If we take 20.8 g of BaCl₂ and 14.2 g of Na₂SO₄:

Total mass of reactants = 20.8 + 14.2 = 35.0 g

We get 23.3 g of BaSO₄ and 11.7 g of NaCl:

Total mass of products = 23.3 + 11.7 = 35.0 g

✔ Mass is conserved!

🔮 Law of Constant Proportions (Proust, 1799)

Joseph Proust proved that in a pure chemical compound, elements are always present in a definite proportion by mass, regardless of the source of the compound.

💧 Water (H₂O)

Hydrogen : Oxygen = 1 : 8 by mass

Whether from a river, well, or lab -- the ratio is always the same!

🔥 Carbon Dioxide (CO₂)

Carbon : Oxygen = 3 : 8 by mass

Whether from burning coal, breathing, or decomposition -- always fixed.

⚡ Ammonia (NH₃)

Nitrogen : Hydrogen = 14 : 3 by mass

Same ratio in lab-made or natural ammonia.

💡 Memory Trick: Lavoisier = Loss (no loss of mass). Proust = Proportion (fixed proportion). "LP" = Laws Pair!

Pause & Ponder

If 12 g of carbon burns completely in 32 g of oxygen, what is the mass of CO₂ formed?

By the Law of Conservation of Mass: mass of CO₂ = 12 + 32 = 44 g. No mass is lost!

Tap to reveal answer

🔬 Dalton's Atomic Theory

In 1808, John Dalton proposed the first scientific atomic theory to explain the laws of chemical combination. His postulates formed the foundation of modern chemistry.

Postulate 1: Matter is Made of Atoms

All matter is made of tiny, indivisible particles called atoms.

Postulate 2: Atoms are Indivisible

Atoms can neither be created nor destroyed. They cannot be divided into smaller particles.

Postulate 3: Same Element = Identical Atoms

All atoms of a given element are identical in mass, size, and chemical properties.

Postulate 4: Different Elements = Different Atoms

Atoms of different elements differ in mass, size, and chemical properties.

Postulate 5: Simple Whole-Number Ratios

Atoms combine in small, whole-number ratios to form compounds (e.g., 1:1, 1:2, 2:3).

Postulate 6: Conservation

In a chemical reaction, atoms are rearranged but not created or destroyed.

✔ Merits

Explained laws of chemical combination
Laid the foundation of modern atomic science
Distinguished elements from compounds

❌ Limitations (Later Discoveries)

Atoms ARE divisible (protons, neutrons, electrons)
Isotopes: atoms of same element can have different masses
Atoms of different elements can have same mass (isobars)
Atoms can be created/destroyed in nuclear reactions

📜 Dalton's Symbols (Hover to see modern!)

Dalton used simple circles and marks as symbols for elements. Hover on each to see the modern symbol!

○HHydrogen

●CCarbon

⊙OOxygen

⊕NNitrogen

⊛SSulfur

∅PPhosphorus

📚 Key Idea: Dalton's theory successfully explained both the Law of Conservation of Mass and the Law of Constant Proportions. It stated that atoms are rearranged during a reaction (conservation) and they combine in fixed ratios (constant proportions).

Pause & Ponder

If Dalton said atoms are indivisible, how do we explain the discovery of electrons, protons, and neutrons?

Dalton's postulate about indivisibility was a limitation. Later discoveries by Thomson (electron, 1897), Rutherford (nucleus/proton, 1911), and Chadwick (neutron, 1932) proved that atoms are divisible into subatomic particles. However, Dalton's basic idea that atoms are the fundamental units of chemical reactions remains valid -- we don't break apart atoms in ordinary chemical reactions.

Tap to reveal answer

📚 Dalton Was Colour-Blind!

John Dalton suffered from colour blindness. In fact, colour blindness was once called "Daltonism" in his honour. Despite this, his contributions to chemistry were revolutionary!

🎓 Ancient Wisdom

The Indian philosopher Maharishi Kanad (600 BCE) described the concept of "parmanu" (atom) centuries before Western scientists. His ideas in the Vaisheshika Sutras were remarkably similar to Dalton's later work!

⚛️ What is an Atom?

An atom is the smallest particle of an element that maintains the chemical identity of that element. Atoms are incredibly tiny!

🔎 Size of an Atom

Atomic radii are measured in nanometres (nm) or angstroms.

1 nm = 10⁻⁹ m (one-billionth of a metre)
1 Å (angstrom) = 10⁻¹⁰ m
Typical atomic radius: 1 to 3 Å

💰

1 coin (~2 cm)

1 atom (~10⁻¹⁰ m)

🔎 10⁸ atoms side by side = 1 cm!

🔮 Hydrogen (H)

Radius: 0.37 Å

Smallest atom!

🔮 Carbon (C)

Radius: 0.77 Å

Basis of organic chemistry

🔮 Oxygen (O)

Radius: 0.73 Å

Essential for life!

🧠 Mind-Blowing Scale!

If an atom were the size of a cricket ball, then the cricket ball would be the size of the Earth! Atoms are unimaginably small.

🌎 How Many Atoms?

A single drop of water contains about 5 × 10²¹ atoms. That is 5,000,000,000,000,000,000,000 atoms!

🔮 Atomic Mass Unit (u): The mass of atoms is measured in atomic mass units (u). 1 u = 1/12 of the mass of a carbon-12 atom = 1.66 × 10⁻²⁴ g. Hydrogen atom mass ≈ 1 u, Carbon atom mass ≈ 12 u, Oxygen atom mass ≈ 16 u.

🔮 Can Atoms Exist Independently?

Most atoms cannot exist independently. They combine with other atoms to form molecules or ions. However, atoms of noble gases (He, Ne, Ar, etc.) can exist independently because their outermost electron shells are completely filled.

✔ Exist Independently

Noble gases: He, Ne, Ar, Kr, Xe
Some metals in vapour form: Na, Fe

❌ Cannot Exist Alone

H, O, N, Cl, F -- form diatomic molecules
P -- forms P₄ molecule
S -- forms S₈ molecule

Pause & Ponder

An atom is so tiny that even millions placed side by side cannot be seen. How did scientists figure out they exist?

Scientists used indirect evidence: (1) The laws of chemical combination showed fixed ratios, suggesting basic building blocks. (2) Brownian motion (random movement of pollen grains in water) suggested invisible particles bumping into them. (3) Modern instruments like electron microscopes and scanning tunnelling microscopes can now actually "see" individual atoms!

Tap to reveal answer

🔢 Symbols of Atoms

Dalton was the first to use symbols for elements (circles with markings). Today, the IUPAC (International Union of Pure and Applied Chemistry) approves names and symbols for elements.

📝 Symbol Rules

Symbol = one or two letters
First letter is always CAPITAL
Second letter (if any) is always small
Example: Co = Cobalt, CO = Carbon monoxide (a compound!)

🌐 Latin Names

Some symbols come from Latin/other language names. These are often confusing because the symbol does not match the English name!

📑 First 20 Elements

Z	Name	Symbol	Z	Name	Symbol
1	Hydrogen	H	11	Sodium	Na
2	Helium	He	12	Magnesium	Mg
3	Lithium	Li	13	Aluminium	Al
4	Beryllium	Be	14	Silicon	Si
5	Boron	B	15	Phosphorus	P
6	Carbon	C	16	Sulfur	S
7	Nitrogen	N	17	Chlorine	Cl
8	Oxygen	O	18	Argon	Ar
9	Fluorine	F	19	Potassium	K
10	Neon	Ne	20	Calcium	Ca

🎓 Latin-Origin Symbols

Element	Symbol	Latin Name
Sodium	Na	Natrium
Potassium	K	Kalium
Iron	Fe	Ferrum
Gold	Au	Aurum
Silver	Ag	Argentum
Copper	Cu	Cuprum
Mercury	Hg	Hydrargyrum
Tin	Sn	Stannum
Lead	Pb	Plumbum
Tungsten	W	Wolfram

💡 Memory Trick for Latin Symbols: "Na K Fe Au Ag Cu Hg Sn Pb W" -- Remember: "Naughty Kids Fear Aunt Agnes' Cup of Hot Snacks with Pb Weight!"

🔭 Mini Periodic Table (First 20 Elements)

Elements with Latin-origin symbols are highlighted in gold. Hover for details!

1

H

1

2

He

4

3

Li

7

4

Be

9

5

B

11

6

C

12

7

N

14

8

O

16

9

F

19

10

Ne

20

11

Na

23

12

Mg

24

13

Al

27

14

Si

28

15

P

31

16

S

32

17

Cl

35.5

18

Ar

40

19

K

39

20

Ca

40

INTERACTIVE GAME

🎮 Match the Element to its Symbol!

Click an element name, then click its matching symbol. Match all pairs!

Element Names:

Symbols:

🧬 Molecules

A molecule is the smallest particle of an element or compound that can exist independently and retains all the properties of that substance.

⚛️ Molecules of Elements

The atoms of many elements exist as molecules. The number of atoms in a molecule is called its atomicity.

⚪

Monoatomic

Single atom
He, Ne, Ar
Noble gases

⚪⚪

Diatomic

2 atoms
H₂, O₂, N₂, Cl₂
Most common

⚪⚪⚪

Triatomic

3 atoms
O₃ (ozone)

🅏

Polyatomic

4+ atoms
P₄, S₈

🔮 Molecules of Compounds

When atoms of different elements combine, they form molecules of compounds.

💧 Water (H₂O)

2 hydrogen + 1 oxygen atoms. Ratio by mass: H:O = 1:8

💨 Carbon Dioxide (CO₂)

1 carbon + 2 oxygen atoms. Ratio by mass: C:O = 3:8

⚡ Ammonia (NH₃)

1 nitrogen + 3 hydrogen atoms. Ratio by mass: N:H = 14:3

💧 Hydrochloric Acid (HCl)

1 hydrogen + 1 chlorine atom. Ratio by mass: H:Cl = 1:35.5

🎬 Animated Water Molecule (H₂O)

O

104.5°

H

H₂O — Bent Shape · 2 Covalent Bonds · 2 Lone Pairs on O

🎬 Watch: H₂ Molecule Formation

H

H + H → H₂ (covalent bond forms!)

📚 Remember: Elements like hydrogen, oxygen, nitrogen, fluorine, chlorine, bromine, and iodine exist as diatomic molecules (H₂, O₂, N₂, F₂, Cl₂, Br₂, I₂). Use the mnemonic: HONClBrIF (say "Honk-le Brif")!

Pause & Ponder

Noble gases (He, Ne, Ar) exist as single atoms (monoatomic). Why don't they form molecules like H₂ or O₂?

Noble gases have completely filled outermost electron shells (He has 2, others have 8). Since their electron configuration is already stable, they have no need to share or transfer electrons. Therefore, they exist as single, independent atoms and do not form molecules.

Tap to reveal answer

Pause & Ponder

Can you tell the difference between 2H and H₂?

2H means 2 separate hydrogen atoms. H₂ means 1 hydrogen molecule made of 2 hydrogen atoms bonded together. The number before the symbol indicates the number of atoms or molecules, while the subscript indicates the number of atoms in one molecule.

Tap to reveal answer

NCERT EXAMPLES

📝 Reading Chemical Symbols

3H₂O = 3 molecules of water = 3 × (2H + 1O) = 6 hydrogen atoms + 3 oxygen atoms

2CO₂ = 2 molecules of carbon dioxide = 2 × (1C + 2O) = 2 carbon atoms + 4 oxygen atoms

5NaCl = 5 formula units of sodium chloride = 5 Na⁺ ions + 5 Cl⁻ ions

P₄ = 1 molecule of phosphorus containing 4 phosphorus atoms (atomicity = 4)

S₈ = 1 molecule of sulfur containing 8 sulfur atoms (atomicity = 8)

⚡ Ions

Not all compounds are made of molecules. Many are made of ions -- charged atoms or groups of atoms.

An ion is an atom or group of atoms that carries an electric charge. Ions form when atoms gain or lose electrons.

➕ Cation (Positive Ion)

Formed when an atom loses electrons.

Examples: Na⁺, K⁺, Ca²⁺, Al³⁺, Fe²⁺, Fe³⁺

Cation = Cat goes positive (paws up!)

➖ Anion (Negative Ion)

Formed when an atom gains electrons.

Examples: Cl⁻, O²⁻, S²⁻, N³⁻

Anion = A Negative ION

🎬 Watch: Na → Na⁺ + Cl → Cl⁻ (Ion Formation)

Stage 1: Na (2,8,1) and Cl (2,8,7) — Both unstable

Na

2, 8, 1

Cl

2, 8, 7

Na⁺

Cl⁻

✅ Stable NaCl — Electrostatic Attraction Holds Ions Together!

🔮 Polyatomic Ions

A polyatomic ion is a group of atoms that carries a charge and acts as a single unit.

Polyatomic Ion	Formula	Charge
Ammonium	NH₄⁺	+1
Hydroxide	OH⁻	-1
Nitrate	NO₃⁻	-1
Bicarbonate	HCO₃⁻	-1
Sulphate	SO₄²⁻	-2
Carbonate	CO₃²⁻	-2
Sulphite	SO₃²⁻	-2
Phosphate	PO₄³⁻	-3

📑 Common Ions Table

Valency 1 (+)	Valency 2 (+)	Valency 3 (+)	Valency 1 (-)	Valency 2 (-)	Valency 3 (-)
Na⁺	Mg²⁺	Al³⁺	Cl⁻	O²⁻	N³⁻
K⁺	Ca²⁺	Fe³⁺	Br⁻	S²⁻	PO₄³⁻
H⁺	Zn²⁺		NO₃⁻	SO₄²⁻
NH₄⁺	Fe²⁺		OH⁻	CO₃²⁻
Ag⁺	Cu²⁺		HCO₃⁻

🔮 How Ions Form: Step by Step

1

Na atom has 11 electrons (2,8,1)

2

Na loses 1 electron from outermost shell

3

Na⁺ formed (2,8) -- now has 10 e⁻ and 11 p⁺

4

Cl atom has 17 electrons (2,8,7)

5

Cl gains 1 electron in outermost shell

6

Cl⁻ formed (2,8,8) -- now has 18 e⁻ and 17 p⁺

Pause & Ponder

Why does sodium lose an electron while chlorine gains one?

Sodium has 1 electron in its outermost shell (2,8,1). It is easier to lose 1 electron than gain 7. Chlorine has 7 electrons in its outermost shell (2,8,7). It is easier to gain 1 electron than lose 7. Both achieve a stable octet (8 electrons) in their outermost shell.

Tap to reveal answer

💡 Tip: Ionic compounds like NaCl, MgO, CaCl₂ do NOT form molecules. They exist as a large collection of ions arranged in a crystal lattice. We write their formula unit, not molecular formula!

🔗 How Atoms Combine

Atoms combine to achieve a stable electron configuration (like noble gases). They do this by sharing or transferring electrons.

⚡ Covalent Bond (Electron Sharing)

A covalent bond forms when two atoms share one or more pairs of electrons. Both atoms contribute electrons to the shared pair.

💡 Sharing ≠ Exchanging!
In a covalent bond, atoms share electrons — the shared pair sits between both atoms and both count it as their own. Nobody gives, nobody takes.

Why sharing? Both H and O are non-metals. Neither is willing to fully give up electrons. So they compromise by sharing. Ionic bonds happen when a metal (eager to give) meets a non-metal (eager to take) — like Na + Cl.

💧 Why H₂O is Covalent — Step by Step

Atom	Valence e⁻	Needs	What happens
Oxygen (O)	6	2 more for octet (8)	Shares 1e⁻ with each H
Hydrogen (H)	1	1 more for duplet (2)	Shares its 1e⁻ with O

H : O : H
··
·· ← lone pairs (O keeps these)

Each : between H and O = 1 shared pair (covalent bond).
O sees: 2 shared pairs + 2 lone pairs = 8 electrons (octet ✓).
Each H sees: 1 shared pair = 2 electrons (duplet ✓).
The 2 lone pairs push the bonds downward → bent shape (104.5°).

🎯 Remember: Sharing = Covalent (both hold on). Transfer = Ionic (one gives, one takes). Non-metal + Non-metal = Covalent. Metal + Non-metal = Ionic.

Single Bond

One shared pair of electrons

Examples: H₂, HCl, H₂O

Represented by: —

Double Bond

Two shared pairs of electrons

Examples: O₂, CO₂

Represented by: =

Triple Bond

Three shared pairs of electrons

Examples: N₂

Represented by: ≡

🔮 Lewis Dot Structures (Electron-Dot Diagrams)

A Lewis dot structure shows the valence electrons of atoms as dots arranged around the element symbol. Shared pairs are shown between atoms.

Lewis Dot Structures of Common Molecules

H—H

Hydrogen (H₂)

H : H

Single bond
1 shared pair

Cl—Cl

Chlorine (Cl₂)

:Cl : Cl:

Single bond
3 lone pairs each

O=O

Oxygen (O₂)

::O :: O::

Double bond
2 shared pairs

N≡N

Nitrogen (N₂)

:N ::: N:

Triple bond
3 shared pairs

H—O—H

Water (H₂O)

H : O : H

Bent shape (104.5°)
2 lone pairs on O

O=C=O

Carbon Dioxide (CO₂)

::O :: C :: O::

Linear shape
2 double bonds

📝 IUPAC Naming of Covalent Compounds

Covalent compounds are named using prefixes to indicate the number of atoms of each element:

Prefix	Number	Prefix	Number
Mono	1	Hexa	6
Di	2	Hepta	7
Tri	3	Octa	8
Tetra	4	Nona	9
Penta	5	Deca	10

💡 Rules: (1) First element retains its name. (2) Second element ends in -ide. (3) Mono is dropped for the first element. (4) Drop the final vowel of a prefix before a vowel-starting element name (e.g., mono + oxide = monoxide, not monooxide).

📝 IUPAC Naming Examples

Formula	IUPAC Name
CO	Carbon monoxide
CO₂	Carbon dioxide
PCl₃	Phosphorus trichloride
SF₆	Sulfur hexafluoride
N₂O₄	Dinitrogen tetroxide
N₂O₅	Dinitrogen pentoxide
CCl₄	Carbon tetrachloride

⚡ Ionic Bond (Electron Transfer)

An ionic bond forms when one atom transfers electrons to another. This creates oppositely charged ions held together by electrostatic attraction.

The resulting compound has a 3D regular arrangement called a crystal lattice.

📚 NaCl Formation: Sodium (2,8,1) loses 1 electron → Na⁺ (2,8). Chlorine (2,8,7) gains 1 electron → Cl⁻ (2,8,8). The Na⁺ and Cl⁻ ions attract each other and arrange in a crystal lattice.

📊 Ionic vs Covalent Compounds

Property	Ionic Compounds	Covalent Compounds
Formation	Electron transfer	Electron sharing
Water solubility	Generally soluble	Generally insoluble
Organic solvent	Insoluble	Generally soluble
Conductivity (solid)	Non-conducting (ions fixed)	Non-conducting
Conductivity (solution)	Conducting (ions move freely)	Non-conducting
Melting/Boiling point	High	Low
Structure	Crystal lattice	Discrete molecules

💡 Did you know? Sugar dissolves in water but does NOT conduct electricity — because it remains as molecules, not ions. This is a key test to distinguish ionic from covalent compounds!

PAUSE & THINK

Diamond (carbon) has a very high melting point, yet it is a covalent compound. How?

Diamond is a covalent network solid — each carbon atom is bonded to 4 others in a continuous 3D network. There are no individual molecules, so you have to break strong covalent bonds to melt it. Most covalent compounds have weak intermolecular forces, but network solids are the exception!

👉 Tap to think

📝 Writing Chemical Formulae

A chemical formula represents the composition of a molecule or compound using element symbols and numbers.

🔗 Valency

The valency of an element is its combining capacity -- the number of bonds it can form. It is determined by the number of electrons an atom gains, loses, or shares to achieve a stable configuration.

📋 Common Valencies

H = 1, O = 2, N = 3, C = 4
Na = 1, Mg = 2, Al = 3
Cl = 1, S = 2, P = 3 (or 5)
Fe = 2 or 3, Cu = 1 or 2

📝 Rules for Formulae

Write symbols of combining elements
Write valency below each
Criss-cross the valencies
Simplify if needed

🔮 Valency at a Glance

H

Valency: 1

O

Valency: 2

N

Valency: 3

C

Valency: 4

Na

Valency: 1

Mg

Valency: 2

Al

Valency: 3

Cl

Valency: 1

S

Valency: 2

Ca

Valency: 2

Fe

Val: 2 or 3

K

Valency: 1

🔄 The Criss-Cross Method

This is the easiest way to write chemical formulae. The valency of one element becomes the subscript of the other.

Example: Aluminium Oxide

Al

O

Valency: Al = 3, O = 2

⤵

Criss-cross: Al gets 2, O gets 3

⤵

Al₂O₃

NCERT EXAMPLE

📝 More Examples of Criss-Cross

NaCl: Na(1) + Cl(1) → Na₁Cl₁ = NaCl

MgO: Mg(2) + O(2) → Mg₂O₂ = MgO (simplify 2:2 to 1:1)

CaCl₂: Ca(2) + Cl(1) → Ca₁Cl₂ = CaCl₂

Na₂O: Na(1) + O(2) → Na₂O₁ = Na₂O

MgCl₂: Mg(2) + Cl(1) → Mg₁Cl₂ = MgCl₂

Ca(OH)₂: Ca(2) + OH(1) → Ca₁(OH)₂ = Ca(OH)₂

Na₂SO₄: Na(1) + SO₄(2) → Na₂(SO₄)₁ = Na₂SO₄

INTERACTIVE TOOL

🛠 Formula Builder

Select a cation and an anion to build a chemical formula using the criss-cross method!

Cation (+):

Anion (-):

📚 Rule for Polyatomic Ions: When a polyatomic ion needs a subscript greater than 1, enclose it in brackets. Example: Ca(OH)₂, Al₂(SO₄)₃, (NH₄)₃PO₄.

📝 Worked Examples: Practice Writing Formulae

WORKED EXAMPLE 1

📝 Aluminium Sulphate

Step 1: Aluminium ion: Al³⁺ (valency 3), Sulphate ion: SO₄²⁻ (valency 2)

Step 2: Write symbols: Al and SO₄

Step 3: Criss-cross valencies: Al gets subscript 2, SO₄ gets subscript 3

Step 4: Since SO₄ is polyatomic and has subscript 3, enclose in brackets

Formula: Al₂(SO₄)₃

WORKED EXAMPLE 2

📝 Ammonium Phosphate

Step 1: Ammonium ion: NH₄⁺ (valency 1), Phosphate ion: PO₄³⁻ (valency 3)

Step 2: Write symbols: NH₄ and PO₄

Step 3: Criss-cross: NH₄ gets subscript 3, PO₄ gets subscript 1

Step 4: NH₄ is polyatomic with subscript 3, so enclose in brackets

Formula: (NH₄)₃PO₄

WORKED EXAMPLE 3

📝 Iron(III) Oxide (Ferric Oxide)

Step 1: Iron(III) ion: Fe³⁺ (valency 3), Oxide ion: O²⁻ (valency 2)

Step 2: Write symbols: Fe and O

Step 3: Criss-cross: Fe gets subscript 2, O gets subscript 3

Formula: Fe₂O₃ (this is rust!)

Pause & Ponder

Iron can form two types of chlorides: FeCl₂ and FeCl₃. How is this possible if the law of constant proportions says elements combine in fixed ratios?

The law of constant proportions applies to a specific compound. FeCl₂ (ferrous chloride) always has Fe:Cl = 56:71 by mass. FeCl₃ (ferric chloride) always has Fe:Cl = 56:106.5 by mass. Iron has variable valency (2 and 3), so it forms two different compounds, each with its own fixed ratio. The law is not violated -- each compound individually has constant proportions!

Tap to reveal answer

📊 Molecular Mass

The molecular mass of a substance is the sum of the atomic masses of all atoms in a molecule of that substance.

NCERT EXAMPLE

📊 Calculating Molecular Mass

Water (H₂O): 2(1) + 1(16) = 2 + 16 = 18 u

CO₂: 1(12) + 2(16) = 12 + 32 = 44 u

HNO₃: 1(1) + 1(14) + 3(16) = 1 + 14 + 48 = 63 u

H₂SO₄: 2(1) + 1(32) + 4(16) = 2 + 32 + 64 = 98 u

C₆H₁₂O₆ (Glucose): 6(12) + 12(1) + 6(16) = 72 + 12 + 96 = 180 u

📝 Formula Unit Mass

For ionic compounds (which don't form molecules), we use formula unit mass instead of molecular mass. It is calculated the same way.

📝 Formula Unit Mass Examples

NaCl: 23 + 35.5 = 58.5 u

CaCO₃: 40 + 12 + 3(16) = 40 + 12 + 48 = 100 u

MgO: 24 + 16 = 40 u

Pause & Ponder

What is the difference between molecular mass and formula unit mass?

Molecular mass is used for covalent/molecular compounds (like H₂O, CO₂, NH₃) that exist as discrete molecules. Formula unit mass is used for ionic compounds (like NaCl, CaCO₃, MgO) that exist as ions in a crystal lattice, not as molecules. Calculation method is the same -- just the terminology differs!

Tap to reveal answer

🔢 Common Molecular Masses

H₂ = 2 u
O₂ = 32 u
H₂O = 18 u
CO₂ = 44 u
NH₃ = 17 u
CH₄ = 16 u

🔢 Common Formula Unit Masses

NaCl = 58.5 u
KCl = 74.5 u
CaCO₃ = 100 u
NaOH = 40 u
MgO = 40 u
CaO = 56 u

INTERACTIVE CALCULATOR

📊 Molecular Mass Calculator

Click a preset compound or type a formula to calculate its molecular mass!

📸 Mole Concept

The mole is a counting unit used in chemistry, just like "dozen" means 12 items. One mole of any substance contains 6.022 × 10²³ particles (atoms, molecules, or ions).

🔮 Avogadro's Number

Nᵃ = 6.022 × 10²³

Named after Amedeo Avogadro. This is the number of particles in 1 mole of any substance.

🔢 Key Relationship

1 mole of atoms = atomic mass in grams

1 mole of molecules = molecular mass in grams

This mass is called the molar mass.

📸 How Big is a Mole?

6.022 × 10²³

= 602,200,000,000,000,000,000,000 particles!

🌊 If you had a mole of grains of sand, it would cover the entire Earth several meters deep!

⭐ A mole of stars would be more than all the stars in the observable universe!

⏰ If you counted 1 million atoms per second, it would take 19 billion years to count 1 mole!

🔮 Mole-Mass-Particle Triangle

NCERT EXAMPLES

📊 Mole Calculations

Q: Calculate the number of moles of 46 g of sodium (Na).

Atomic mass of Na = 23 u, so molar mass = 23 g/mol

Moles = mass / molar mass = 46 / 23 = 2 moles

Q: Calculate the mass of 0.5 mole of water (H₂O).

Molecular mass of H₂O = 18 u, so molar mass = 18 g/mol

Mass = moles × molar mass = 0.5 × 18 = 9 g

Q: How many molecules are in 36 g of water?

Moles = 36 / 18 = 2 moles

Molecules = 2 × 6.022 × 10²³ = 1.2044 × 10²⁴ molecules

Q: Calculate the number of atoms in 0.1 mole of carbon.

Atoms = 0.1 × 6.022 × 10²³ = 6.022 × 10²² atoms

INTERACTIVE CALCULATOR

📊 Mole Calculator

Enter the molar mass and either mass or moles to calculate the other values!

Molar Mass (g/mol):

Mass (grams):

OR enter Moles:

💡 Quick Formulas:
Moles = Mass / Molar mass
Mass = Moles × Molar mass
Number of particles = Moles × 6.022 × 10²³
Moles = Number of particles / 6.022 × 10²³

Pause & Ponder

1 mole of carbon atoms and 1 mole of sodium atoms both contain the same number of atoms. But do they have the same mass?

No! Both contain 6.022 × 10²³ atoms, but they have different masses. 1 mole of C = 12 g and 1 mole of Na = 23 g. The number of particles is the same, but the mass depends on the atomic mass of the element.

Tap to reveal answer

Pause & Ponder

Which has more atoms: 16 g of oxygen or 12 g of carbon?

16 g of O = 1 mole of O atoms = 6.022 × 10²³ atoms. 12 g of C = 1 mole of C atoms = 6.022 × 10²³ atoms. Both have the same number of atoms! This is the beauty of the mole concept -- 1 mole of any element always contains the same number of atoms.

Tap to reveal answer

📅 Mole Day!

Mole Day is celebrated on October 23 (10/23) from 6:02 AM to 6:02 PM, in honour of Avogadro's number (6.02 × 10²³). Chemistry enthusiasts worldwide celebrate with activities and experiments!

🔬 Why Carbon-12?

The atomic mass unit is defined as 1/12 of the mass of carbon-12 because carbon-12 is the most abundant, stable isotope. It provides a convenient standard for comparing masses of all other atoms.

ADDITIONAL EXAMPLES

📊 More Mole Problems

Q: What is the mass of 1 mole of H₂ molecules?

Molecular mass of H₂ = 2 × 1 = 2 u

Mass of 1 mole of H₂ = 2 g

Q: Calculate the number of molecules in 11 g of CO₂.

Molar mass of CO₂ = 44 g/mol

Moles = 11/44 = 0.25 mol

Molecules = 0.25 × 6.022 × 10²³ = 1.5055 × 10²³ molecules

Q: How many moles are in 3.011 × 10²³ atoms of iron?

Moles = Number of atoms / Avogadro's number

= 3.011 × 10²³ / 6.022 × 10²³ = 0.5 mol

Q: What is the mass of 0.2 mole of NaOH?

Molecular mass of NaOH = 23 + 16 + 1 = 40 u

Mass = 0.2 × 40 = 8 g

🧪 NCERT Activities — Animated

ACTIVITY 9.1 — ANIMATED

⚖️ Dissolving Salt — Conservation of Mass (Physical Change)

Aim: Verify that mass is conserved during dissolution (a physical change).

Procedure: Weigh salt, weigh water, dissolve salt in water, weigh the solution.

Salt (5g)

Water (50g)

55g

Total Mass: 55 g

Step 1: Salt + Water placed on left pan

⚖️ Observation: Mass of salt (5g) + Mass of water (50g) = Mass of solution (55g). Mass is conserved!

Conclusion: Mass is conserved during physical changes like dissolution.

ACTIVITY 9.2 — ANIMATED

💧 Vinegar + Baking Soda: Open vs Closed System

Aim: Show why a closed system is needed to verify the law of conservation of mass in a gas-producing reaction.

Reaction: NaHCO₃ + CH₃COOH → CH₃COONa + H₂O + CO₂↑

🚫 Open System

CO₂CO₂CO₂

Before: 80g
After: 80g

Click "Start" to see

✅ Closed System

Before: 80g
After: 80g

Click "Start" to see

💡 Key Insight: In the open flask, CO₂ gas escapes into the air. The mass appears to decrease. In the closed flask (with balloon), ALL products are retained — mass stays exactly the same! The law is always valid; we just need to account for ALL products.

ACTIVITY 9.3 — ANIMATED

🔬 Precipitation Reaction — Conservation of Mass

Aim: Verify law of conservation of mass using a precipitation reaction in a closed flask.

Reaction: BaCl₂(aq) + Na₂SO₄(aq) → BaSO₄(s)↓ + 2NaCl(aq)

Na₂SO₄

BaCl₂

⚖️

Mass: 120.0 g

Observation: White precipitate of BaSO₄ forms. Mass before = Mass after = 120.0 g

Aim: Show that water always contains H and O in a fixed ratio (1:8 by mass), regardless of source.

🚰

Tap Water

H

O

H:O = ?

🌊

River Water

H

O

H:O = ?

🌲

Well Water

H

O

H:O = ?

🌧️

Rain Water

H

O

H:O = ?

Pause & Ponder

When you burn a piece of paper, the ash left behind weighs less than the original paper. Does this violate the law of conservation of mass?

No! When paper burns, it produces CO₂ and H₂O gases which escape into the air. If we could collect ALL the products (ash + CO₂ + H₂O + other gases), their total mass would equal the mass of paper + oxygen used. The reaction takes place in an open system, so we lose track of gaseous products.

👉 Tap to reveal answer

📋 Quick Revision

Law of Conservation of Mass

Mass can neither be created nor destroyed in a chemical reaction. Total mass of reactants = total mass of products.

Law of Constant Proportions

A pure compound always contains the same elements combined in the same fixed ratio by mass.

Dalton's Atomic Theory

All matter is made of indivisible atoms. Atoms of same element are identical. Atoms combine in simple ratios.

Atom

Smallest particle of an element. Size: ~10⁻¹⁰ m. Mass measured in atomic mass units (u).

IUPAC Symbols

1 or 2 letters. First always capital. Some from Latin: Na, K, Fe, Au, Ag, Cu, Hg, Sn, Pb, W.

Molecules

Group of bonded atoms. Atomicity: mono, di, tri, poly. H₂, O₂, H₂O, CO₂, NH₃.

Ions

Charged atoms/groups. Cation (+): Na⁺, Ca²⁺. Anion (-): Cl⁻, SO₄²⁻. Polyatomic: OH⁻, NH₄⁺.

Chemical Formulae

Use criss-cross method. Valency of one element becomes subscript of the other. Simplify if needed.

Molecular Mass

Sum of atomic masses of all atoms in a molecule. Measured in u (atomic mass units).

Mole Concept

1 mole = 6.022 × 10²³ particles = molar mass in grams. Moles = mass / molar mass.

💡 Key Formulas at a Glance:
• Molecular mass = Sum of atomic masses of all atoms
• Moles = Mass (g) / Molar mass (g/mol)
• Number of particles = Moles × 6.022 × 10²³
• 1 u = 1.66 × 10⁻²⁴ g
• Molar mass (numerically) = Atomic/Molecular mass

📑 Important Distinctions

Concept	Definition	Example
Atom	Smallest particle of an element that takes part in a reaction	H, O, Na, C
Molecule	Group of atoms bonded together; smallest unit that can exist independently	H₂, O₂, H₂O
Ion	Atom or group of atoms with a charge	Na⁺, Cl⁻, SO₄²⁻
Atomic mass	Mass of one atom in u	C = 12 u
Molecular mass	Sum of atomic masses of all atoms in a molecule	H₂O = 18 u
Molar mass	Mass of 1 mole in grams	H₂O = 18 g/mol
Valency	Combining capacity of an element	Na = 1, O = 2, Al = 3
Atomicity	Number of atoms in one molecule	O₂ = 2, O₃ = 3

🔢 Key Numbers to Remember

🔮 Avogadro's Number

6.022 × 10²³ particles per mole

🔮 1 u (amu)

1.66 × 10⁻²⁴ g

🔮 1 Å (Angstrom)

10⁻¹⁰ m = 0.1 nm

🔮 Atomic Masses

H=1, C=12, N=14, O=16, Na=23, Mg=24, Al=27, S=32, Cl=35.5, Ca=40, Fe=56

🧠 MCQs (Multiple Choice Questions)

1. The law of conservation of mass was given by:
- Dalton
- Proust
- Lavoisier
- Avogadro
✔ Lavoisier (1774) established this law.
2. In water, hydrogen and oxygen combine in the ratio of 1:8 by mass. This is an example of:
- Law of conservation of mass
- Law of constant proportions
- Dalton's atomic theory
- Avogadro's law
✔ Fixed ratio by mass = Law of constant proportions (Proust).
3. Which of the following is NOT a postulate of Dalton's atomic theory?
- Atoms are indivisible
- Atoms of the same element are identical
- Atoms of the same element can have different masses
- Atoms combine in simple whole-number ratios
✔ Dalton said atoms of the same element are identical. Different masses (isotopes) was discovered later.
4. The chemical symbol for sodium is Na. This comes from the Latin name:
- Natronium
- Natrium
- Sodium
- Natricus
✔ Na comes from Natrium (Latin).
5. The atomicity of ozone (O₃) is:
- 1
- 2
- 3
- 4
✔ Ozone has 3 oxygen atoms, so atomicity = 3 (triatomic).
6. Which of the following is a polyatomic ion?
- Na⁺
- Cl⁻
- SO₄²⁻
- O²⁻
✔ SO₄²⁻ (sulphate) is made of multiple atoms acting as one ion.
7. The formula of aluminium oxide is:
- AlO
- Al₃O₂
- Al₂O₃
- AlO₃
✔ Al(3) + O(2) → criss-cross → Al₂O₃.
8. The molecular mass of H₂SO₄ is:
- 96 u
- 98 u
- 100 u
- 94 u
✔ 2(1) + 32 + 4(16) = 2 + 32 + 64 = 98 u.
9. Avogadro's number is:
- 6.022 × 10²²
- 6.022 × 10²³
- 6.022 × 10²⁴
- 6.022 × 10²⁵
✔ Avogadro's number = 6.022 × 10²³ particles per mole.
10. The mass of 2 moles of oxygen atoms is:
- 16 g
- 32 g
- 64 g
- 8 g
✔ Atomic mass of O = 16 u. 2 moles of O atoms = 2 × 16 = 32 g.
11. Which ion has a charge of -2?
- Na⁺
- Cl⁻
- O²⁻
- Al³⁺
✔ Oxide ion (O²⁻) has a charge of -2.
12. How many atoms are present in 1 mole of hydrogen gas (H₂)?
- 6.022 × 10²³
- 2 × 6.022 × 10²³
- 3 × 6.022 × 10²³
- 6.022 × 10²²
✔ 1 mole of H₂ has 6.022 × 10²³ molecules, each with 2 atoms = 2 × 6.022 × 10²³ = 1.2044 × 10²⁴ atoms.
13. What is the mass of 3 moles of water?
- 36 g
- 54 g
- 18 g
- 72 g
✔ Molar mass of H₂O = 18 g/mol. Mass = 3 × 18 = 54 g.
14. The formula unit mass of NaCl is:
- 23 u
- 35.5 u
- 58.5 u
- 46 u
✔ Na = 23 u, Cl = 35.5 u. Formula unit mass = 23 + 35.5 = 58.5 u.
15. Which of the following is a correct chemical formula?
- NaCl₂
- Mg₂O
- CaCl₂
- AlO
✔ Ca(2) + Cl(1) → CaCl₂ is correct. NaCl (not NaCl₂), MgO (not Mg₂O), Al₂O₃ (not AlO).

✏️ Fill in the Blanks

1. The law of conservation of mass was proposed by __________.

✔ Antoine Lavoisier

2. In a pure compound, elements are always present in a definite __________ by mass.

✔ proportion (ratio)

3. According to Dalton, the smallest particle of an element that maintains its identity is called a(n) __________.

✔ atom

4. 1 atomic mass unit (u) = __________ of the mass of a carbon-12 atom.

✔ 1/12

5. The number of atoms in a molecule is called its __________.

✔ atomicity

6. A positively charged ion is called a __________ and a negatively charged ion is called a(n) __________.

✔ cation; anion

7. The combining capacity of an element is called its __________.

✔ valency

8. One mole of any substance contains __________ particles.

✔ 6.022 × 10²³ (Avogadro's number)

9. The molecular mass of water (H₂O) is __________ u.

✔ 18 u (2 × 1 + 16 = 18)

10. The Latin name for iron is __________, so its symbol is __________.

✔ Ferrum; Fe

✓ True or False

1. In a chemical reaction, atoms are created and destroyed.

✘ False. According to the law of conservation of mass and Dalton's theory, atoms are only rearranged, never created or destroyed in a chemical reaction.

2. Water obtained from different sources always has the same ratio of hydrogen to oxygen by mass.

✔ True. This is the law of constant proportions. H:O ratio is always 1:8 by mass in water.

3. An atom is the smallest particle that can exist independently.

✘ False. An atom is the smallest particle of an element that takes part in a chemical reaction. It may or may not exist independently. A molecule is the smallest particle that can exist independently.

4. The symbol for cobalt is CO.

✘ False. The symbol for cobalt is Co (small 'o'). CO (both capital) represents carbon monoxide, a compound.

5. Ionic compounds are made up of ions, not molecules.

✔ True. Ionic compounds like NaCl consist of ions (Na⁺ and Cl⁻) arranged in a crystal lattice, not individual molecules.

6. One mole of carbon atoms has a mass of 12 grams.

✔ True. The molar mass of carbon = 12 g/mol. So 1 mole of carbon atoms = 12 g.

7. The molecular mass of a substance and its molar mass have the same numerical value.

✔ True. Molecular mass is expressed in 'u' and molar mass in 'g/mol', but they have the same numerical value. E.g., H₂O: molecular mass = 18 u, molar mass = 18 g/mol.

8. The formula of calcium chloride is CaCl because calcium has valency 2 and chlorine has valency 1.

✘ False. Using criss-cross: Ca(2) + Cl(1) gives Ca₁Cl₂ = CaCl₂, not CaCl.

✍️ Short Answer Questions

Q1. State the law of conservation of mass.

Mass can neither be created nor destroyed in a chemical reaction. The total mass of the reactants is always equal to the total mass of the products. This law was given by Antoine Lavoisier in 1774.

Q2. What is the difference between molecules of elements and molecules of compounds?

Molecules of elements are formed by atoms of the same element (e.g., H₂, O₂, O₃, P₄). Molecules of compounds are formed by atoms of different elements (e.g., H₂O, CO₂, NH₃). Molecules of compounds can be broken down into simpler substances by chemical methods.

Q3. What is an ion? Distinguish between a cation and an anion.

An ion is an electrically charged atom or group of atoms formed by the loss or gain of electrons. A cation is a positively charged ion formed by losing electrons (e.g., Na⁺, Ca²⁺). An anion is a negatively charged ion formed by gaining electrons (e.g., Cl⁻, O²⁻).

Q4. Write the chemical formulae of: (a) Magnesium chloride (b) Calcium oxide (c) Aluminium sulphate

(a) MgCl₂: Mg(2) + Cl(1) → criss-cross → MgCl₂
(b) CaO: Ca(2) + O(2) → criss-cross → Ca₂O₂ = CaO (simplify)
(c) Al₂(SO₄)₃: Al(3) + SO₄(2) → criss-cross → Al₂(SO₄)₃

Q5. Calculate the molecular mass of glucose (C₆H₁₂O₆).

Molecular mass = 6(12) + 12(1) + 6(16) = 72 + 12 + 96 = 180 u

Q6. What is a mole? What is Avogadro's number?

A mole is the amount of a substance that contains as many particles (atoms, molecules, or ions) as there are atoms in exactly 12 g of carbon-12. Avogadro's number (Nᵃ) = 6.022 × 10²³. It represents the number of particles in one mole of any substance.

Q7. Why are symbols of some elements derived from their Latin names?

Many elements were known since ancient times and had Latin or other classical language names. When Berzelius proposed a systematic naming system, he used abbreviations from these Latin names. For example: Gold (Aurum) = Au, Silver (Argentum) = Ag, Iron (Ferrum) = Fe. IUPAC continues this tradition.

Q8. Define atomicity. Give examples.

Atomicity is the number of atoms present in one molecule of an element. Examples: Monoatomic -- He, Ne, Ar (atomicity = 1); Diatomic -- H₂, O₂, N₂ (atomicity = 2); Triatomic -- O₃ (atomicity = 3); Tetratomic -- P₄ (atomicity = 4); Polyatomic -- S₈ (atomicity = 8).

Q9. Calculate the number of moles in 46 g of sodium (Na = 23 u).

Molar mass of Na = 23 g/mol. Moles = mass / molar mass = 46 / 23 = 2 moles.

Q10. What is the formula unit mass of CaCO₃? (Ca = 40, C = 12, O = 16)

Formula unit mass = 40 + 12 + 3(16) = 40 + 12 + 48 = 100 u.

📖 Long Answer Questions

Q1. State and explain Dalton's atomic theory. Mention its merits and limitations.

Dalton's Atomic Theory (1808):

Postulates:

All matter is made of tiny, indivisible particles called atoms.
Atoms cannot be created or destroyed.
Atoms of the same element are identical in mass, size, and properties.
Atoms of different elements are different in mass, size, and properties.
Atoms combine in small, whole-number ratios to form compounds.
In chemical reactions, atoms are rearranged but never created or destroyed.

Merits: (i) It explained the laws of chemical combination. (ii) It distinguished between elements and compounds. (iii) It laid the foundation of modern chemistry.

Limitations: (i) Atoms are divisible (contain electrons, protons, neutrons). (ii) Isotopes show that atoms of the same element can have different masses. (iii) Isobars show atoms of different elements can have the same mass. (iv) Atoms can be created/destroyed in nuclear reactions. (v) Simple whole-number ratios are not always true (e.g., sugar C₁₂H₂₂O₁₁).

Q2. Explain how to write the chemical formula of a compound using the criss-cross method. Give three examples.

Steps for the Criss-Cross Method:

Write the symbols of the two elements/ions side by side (metal/cation first).
Write the valency of each element below its symbol.
Cross over the valencies -- the valency of one element becomes the subscript of the other.
If the subscripts can be simplified (divided by a common factor), simplify them.
If a polyatomic ion needs a subscript > 1, enclose it in brackets.

Examples:

1. Sodium oxide: Na(1) + O(2) → Na₂O₁ = Na₂O

2. Calcium chloride: Ca(2) + Cl(1) → Ca₁Cl₂ = CaCl₂

3. Aluminium sulphate: Al(3) + SO₄(2) → Al₂(SO₄)₃ = Al₂(SO₄)₃

Q3. Explain the mole concept. Calculate: (a) the number of molecules in 9 g of water, (b) the mass of 3.011 × 10²³ molecules of CO₂.

Mole Concept: A mole is the SI unit for the amount of a substance. 1 mole of any substance contains 6.022 × 10²³ particles (Avogadro's number). The mass of 1 mole of a substance in grams equals its molecular/atomic mass in 'u'.

(a) Molecules in 9 g of water:

Molar mass of H₂O = 18 g/mol

Moles = 9/18 = 0.5 mol

Molecules = 0.5 × 6.022 × 10²³ = 3.011 × 10²³ molecules

(b) Mass of 3.011 × 10²³ molecules of CO₂:

Moles = 3.011 × 10²³ / 6.022 × 10²³ = 0.5 mol

Molar mass of CO₂ = 44 g/mol

Mass = 0.5 × 44 = 22 g

Q4. Explain the difference between an atom and a molecule. Give examples of monoatomic, diatomic, and polyatomic molecules.

Atom: The smallest particle of an element that takes part in a chemical reaction. It may or may not exist independently. Example: Na, H, O atoms.

Molecule: The smallest particle of a substance (element or compound) that can exist independently. It consists of one or more atoms bonded together. Example: H₂, O₂, H₂O.

Monoatomic molecules: Noble gases exist as single atoms -- He, Ne, Ar, Kr. Their atomicity is 1.

Diatomic molecules: H₂, O₂, N₂, Cl₂, F₂. Atomicity = 2.

Polyatomic molecules: O₃ (ozone, 3 atoms), P₄ (phosphorus, 4 atoms), S₈ (sulfur, 8 atoms).

Q5. State and explain the law of constant proportions. How does Dalton's atomic theory explain this law?

Law of Constant Proportions (Proust, 1799): In a chemical substance (pure compound), the elements are always present in definite proportions by mass, regardless of the source or method of preparation.

Example: Water (H₂O) always contains hydrogen and oxygen in the ratio 1:8 by mass. Whether water comes from a tap, river, rain, or is made in a lab, this ratio never changes.

Dalton's Explanation: According to Dalton, atoms of a given element have a fixed mass. When atoms combine, they do so in simple, fixed, whole-number ratios. Since each atom has a definite mass, the mass ratio of elements in a compound is always constant. For example, in water, 2 hydrogen atoms (each 1 u) always combine with 1 oxygen atom (16 u), giving a mass ratio of 2:16 = 1:8.

Q6. What are polyatomic ions? Give examples and explain how they are used in writing chemical formulae.

Polyatomic ions are groups of atoms that carry a net charge and behave as a single unit in chemical reactions. They contain two or more atoms covalently bonded together, with an overall positive or negative charge.

Examples:

Ammonium (NH₄⁺) -- a positive polyatomic ion (cation)
Hydroxide (OH⁻) -- charge -1
Sulphate (SO₄²⁻) -- charge -2
Carbonate (CO₃²⁻) -- charge -2
Nitrate (NO₃⁻) -- charge -1
Phosphate (PO₄³⁻) -- charge -3

In formulae: When a polyatomic ion has a subscript greater than 1, it must be enclosed in brackets. For example: Ca(OH)₂ (calcium hydroxide), Al₂(SO₄)₃ (aluminium sulphate), (NH₄)₂SO₄ (ammonium sulphate).

Q7. A 3.0 g sample of carbon is burned completely in oxygen. If 11.0 g of CO₂ is formed, verify the law of conservation of mass. How much oxygen was used?

Reaction: C + O₂ → CO₂

By the law of conservation of mass:

Mass of reactants = Mass of products

Mass of C + Mass of O₂ = Mass of CO₂

3.0 + Mass of O₂ = 11.0

Mass of O₂ used = 11.0 - 3.0 = 8.0 g

Verification: Total mass of reactants = 3.0 + 8.0 = 11.0 g = Total mass of products. The law is verified!

Also, C:O ratio = 3:8, which is consistent with the law of constant proportions for CO₂.

💡 HOTS (Higher Order Thinking Skills)

Q1. If 10 g of CaCO₃ is heated and produces 5.6 g of CaO, what mass of CO₂ is released?

Reaction: CaCO₃ → CaO + CO₂

By the law of conservation of mass:

Mass of CaCO₃ = Mass of CaO + Mass of CO₂

10 = 5.6 + Mass of CO₂

Mass of CO₂ = 10 - 5.6 = 4.4 g

Q2. A compound has the formula X₂O₃. If the molecular mass is 102 u, find the atomic mass of element X.

Molecular mass of X₂O₃ = 2(X) + 3(16) = 102

2X + 48 = 102

2X = 54

X = 27 u

Element X is Aluminium (Al), and the compound is Al₂O₃ (Aluminium oxide).

Q3. Calculate the number of aluminium ions present in 0.051 g of aluminium oxide (Al₂O₃). [Atomic masses: Al = 27, O = 16]

Molecular mass of Al₂O₃ = 2(27) + 3(16) = 54 + 48 = 102 u

Molar mass = 102 g/mol

Moles of Al₂O₃ = 0.051 / 102 = 5 × 10⁻⁴ mol

Each formula unit has 2 Al³⁺ ions.

Number of Al³⁺ ions = 2 × 5 × 10⁻⁴ × 6.022 × 10²³

= 6.022 × 10²⁰ aluminium ions

Q4. In a reaction, 5.3 g of sodium carbonate reacted with 6 g of ethanoic acid. The products formed were 2.2 g of carbon dioxide, 0.9 g of water, and 8.2 g of sodium ethanoate. Show that these results follow the law of conservation of mass.

Total mass of reactants:

= Sodium carbonate + Ethanoic acid = 5.3 + 6.0 = 11.3 g

Total mass of products:

= CO₂ + H₂O + Sodium ethanoate = 2.2 + 0.9 + 8.2 = 11.3 g

Since total mass of reactants (11.3 g) = total mass of products (11.3 g), the law of conservation of mass is verified.

Q5. Calcium chloride (CaCl₂) has a formula unit mass of 111 u. (a) What is the mass of 0.5 mole of CaCl₂? (b) How many formula units are in 0.5 mole? (c) How many chloride ions are present?

(a) Mass:

Molar mass of CaCl₂ = 111 g/mol

Mass = moles × molar mass = 0.5 × 111 = 55.5 g

(b) Formula units:

= 0.5 × 6.022 × 10²³ = 3.011 × 10²³ formula units

(c) Chloride ions:

Each formula unit of CaCl₂ has 2 Cl⁻ ions.

Total Cl⁻ ions = 2 × 3.011 × 10²³ = 6.022 × 10²³ chloride ions

🎮 Interactive Quiz

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⚛️ Atomic Foundations of Matter

🔬 Key Scientists

🎓 Why These Laws Matter

⚖️ Law of Conservation of Mass (Lavoisier, 1774)

🔮 Law of Constant Proportions (Proust, 1799)

💧 Water (H₂O)

🔥 Carbon Dioxide (CO₂)

⚡ Ammonia (NH₃)

Postulate 1: Matter is Made of Atoms

Postulate 2: Atoms are Indivisible

Postulate 3: Same Element = Identical Atoms

Postulate 4: Different Elements = Different Atoms

Postulate 5: Simple Whole-Number Ratios

Postulate 6: Conservation

✔ Merits

❌ Limitations (Later Discoveries)

📜 Dalton's Symbols (Hover to see modern!)

📚 Dalton Was Colour-Blind!

🎓 Ancient Wisdom

🔎 Size of an Atom

🔮 Hydrogen (H)

🔮 Carbon (C)

🔮 Oxygen (O)

🧠 Mind-Blowing Scale!

🌎 How Many Atoms?

🔮 Can Atoms Exist Independently?

✔ Exist Independently

❌ Cannot Exist Alone

📝 Symbol Rules

🌐 Latin Names

📑 First 20 Elements

🎓 Latin-Origin Symbols

🔭 Mini Periodic Table (First 20 Elements)

🎮 Match the Element to its Symbol!

⚛️ Molecules of Elements

Monoatomic

Diatomic

Triatomic

Polyatomic

🔮 Molecules of Compounds

💧 Water (H₂O)

💨 Carbon Dioxide (CO₂)

⚡ Ammonia (NH₃)

💧 Hydrochloric Acid (HCl)

🎬 Animated Water Molecule (H₂O)

🎬 Watch: H₂ Molecule Formation

➕ Cation (Positive Ion)

➖ Anion (Negative Ion)

🎬 Watch: Na → Na⁺ + Cl → Cl⁻ (Ion Formation)

🔮 Polyatomic Ions

📑 Common Ions Table

🔮 How Ions Form: Step by Step

⚡ Covalent Bond (Electron Sharing)

Single Bond

Double Bond

Triple Bond

🔮 Lewis Dot Structures (Electron-Dot Diagrams)

Lewis Dot Structures of Common Molecules

Hydrogen (H₂)

Chlorine (Cl₂)

Oxygen (O₂)

Nitrogen (N₂)

Water (H₂O)

Carbon Dioxide (CO₂)

📝 IUPAC Naming of Covalent Compounds

⚡ Ionic Bond (Electron Transfer)

📊 Ionic vs Covalent Compounds

🔗 Valency

📋 Common Valencies

📝 Rules for Formulae

🔮 Valency at a Glance

🔄 The Criss-Cross Method

🛠 Formula Builder

📝 Worked Examples: Practice Writing Formulae

📝 Formula Unit Mass

🔢 Common Molecular Masses

🔢 Common Formula Unit Masses

📊 Molecular Mass Calculator

🔮 Avogadro's Number

🔢 Key Relationship