Notes Class 9 Science Exploration Chapter 8 Journey Inside the Atom

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Chapter Introduction: What Is an Atom?

Everything around us — your desk, your food, the air you breathe, and even your own body — is made of matter. And all matter is made of incredibly tiny particles called atoms (परमाणु). These atoms are so small that they cannot be seen with the naked eye!

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Ancient Roots: Indian & Greek Origins

More than 2,000 years ago, the Indian philosopherAcharya Kanadaproposed that matter (dravya) can be divided until you reach the smallest particle — calledparmanu(recorded in Vaisesika Sutras). Around the same time in Greece,Leucippus and Democrituscalled the same idea atomos (meaning indivisible in Greek). Both civilisations independently hit upon the same idea!

📅 Dalton’s Atomic Theory (1808)

In 1808, John Dalton proposed the first scientific atomic theory based on experiments:

All matter is composed of tiny, indivisible particles called atoms.
Atoms of the same element are identical; atoms of different elements are different.
Atoms cannot be created or destroyed in a chemical reaction.
Compounds are formed when atoms of different elements combine in fixed ratios.

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Key Point to Remember

The concept of ‘atom’ originally came from imagination and philosophy, NOT from experiments. It was Dalton who made it a scientific, experimental idea.

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Think It Over

After Dalton, three key questions remained: What are atoms made of? What do they look like? What makes atoms of one element different from another?

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Historical Journey Through Atomic Models

As scientists performed new experiments, they kept improving the model of the atom. Each model was a step forward, but none was perfect until more evidence arrived.

1808

Dalton

Atom = indivisible sphere

1897

Thomson

Plum pudding model

1911

Rutherford

Nuclear / Planetary model

1913

Bohr

Fixed energy shells

Modern

Quantum

Electron cloud model

⚡ Thomson’s Plum Pudding Model (1897)

In 1897, J. J. Thomson studied cathode rays by passing electricity through gas at very low pressure. He discovered electrons — tiny negatively charged particles.

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Cathode Ray Experiment

Thomson used a glass tube with two electrodes (cathode and anode) and applied high voltage. He observed rays going from the cathode (−) to the anode (+). These were calledcathode raysand were found to be streams of electrons. Crucially, the nature of these rays was the same regardless of the cathode material — proving electrons are present in ALL atoms!

Since atoms are neutral but contain negative electrons, Thomson asked: where is the positive charge? His answer was the Plum Pudding Model:

The atom is a sphere of positive charge.
Electrons are embedded (scattered) throughout this positive sphere like plums in a pudding.
Indian analogy: think of it like a watermelon 🍉 — the red pulp is the positive charge, and the seeds are the electrons!

Charge of an electron = −1.602 × 10⁻¹⁹ C (represented as −1 for convenience)

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J. J. Thomson (Nobel Prize 1906)

Discovered the electron — the first subatomic particle. He won the Nobel Prize in Physics in 1906 for his studies on electrical conductivity of gases. As head of the Cavendish Laboratory, Cambridge, he mentored many scientists including Ernest Rutherford.

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The Gold Foil Experiment & Rutherford’s Model

🧪 The Gold Foil Experiment (1911)

In 1911, Geiger and Marsden, working under Ernest Rutherford, tested Thomson’s model by firing a beam of positively charged alpha (α) particles at an extremely thin gold foil.

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What is an Alpha Particle?

An alpha particle is the nucleus of a helium atom — it contains2 protons + 2 neutronsand carries a positive charge. They are emitted from certain radioactive elements.

📊 Observations vs. Expectations

What Thomson’s Model Predicted	What Actually Happened
All α-particles should pass straight through or be deflected only slightly	Most particles DID pass straight through ✅
No sharp deflections expected	Some particles were deflected at very large angles ❌
Definitely no particles bouncing back	A few particles bounced STRAIGHT BACK ❌

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Thomson’s Model FAILED!

The gold foil experiment completely disproved the plum pudding model. If positive charge were spread evenly throughout the atom, large deflections would be impossible. A bounced-back alpha particle meant it hit something extremely dense and heavy — a concentrated nucleus!

🌟 Rutherford’s Conclusions — The Nuclear (Planetary) Model

Most of the atom is empty space — this explains why most α-particles passed through undeflected.
All the positive charge and most of the mass are concentrated in a tiny region at the center called the nucleus (केंद्रक).
Electrons revolve around the nucleus in orbits, just like planets orbit the Sun. Hence, this is called the Planetary Model.

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Size Comparison — Atom vs. Nucleus

The nucleus is about10⁵ (one lakh) times smallerthan the atom! If an atom were the size of acricket ground(about 100 m across), the nucleus would be just a tinyblack pepper grain(a few mm) at the centre. The rest is just empty space!

Diameter of atom ≈ 10⁻¹⁰ m | Diameter of nucleus ≈ 10⁻¹⁵ m

⚠️ Limitation of Rutherford’s Model

While it was a big step forward, Rutherford’s model couldn’t explain why atoms are stable. Here’s the problem:

A charged particle moving in a circular path is constantly accelerating (changing direction).
An accelerating charged particle should continuously lose energy.
If electrons lose energy, they should spiral inward and eventually crash into the nucleus.
But this DOESN’T happen — atoms are stable! So Rutherford’s model needed improvement.

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Discovery of the Proton — Rutherford Again!

Rutherford also discovered that the nucleus contains positively charged particles calledprotons (प्रोटॉन). Protons are much heavier than electrons and carry a charge equal and opposite (+1) to that of electrons (−1). For an atom to be electrically neutral, the number of protons must equal the number of electrons.

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Ernest Rutherford — Father of Nuclear Physics (Nobel Prize 1908)

Born in New Zealand, Rutherford discovered the atomic nucleus and explained radioactive decay. He won the 1908 Nobel Prize in Chemistry. His portrait appears on New Zealand’s $100 banknote!

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Bohr’s Model: Fixed Energy Shells (1913)

To solve the stability problem, Danish physicist Niels Bohr proposed a new model in 1913 with a revolutionary idea: electrons move in fixed paths and don’t lose energy while doing so!

📋 Key Postulates of Bohr’s Model

Electrons move in fixed circular paths around the nucleus called orbits, shells, or stationary states (स्थिर कक्षाएं).
Each shell has a definite, fixed energy, so they are also called energy levels.
Shells are represented by letters K, L, M, N, … or numbers n = 1, 2, 3, 4, …
While in a fixed shell, an electron does NOT lose energy.
The K-shell (n=1) is closest to the nucleus and has the least energy.
Energy of shells increases as we move farther from the nucleus.
An electron can jump to another shell by absorbing or releasing a fixed amount of energy equal to the energy difference between the two levels.

NucleusK (n=1)e⁻L (n=2)e⁻M (n=3)e⁻↑ Energy increases outward

Fig: Bohr’s model showing K, L, M shells (energy levels) around the nucleus

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Why are shells named K, L, M, N and not A, B, C?

The naming comes from early X-ray experiments by physicist Charles Barkla. He named the first X-ray line “K” and left letters A–J open in case earlier series were discovered (none were). Bohr adopted this same notation.

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Niels Bohr — Nobel Prize in Physics 1922

A professor at Copenhagen University, Denmark, Bohr was puzzled by why electrons don’t fall into the nucleus. His atomic model explained stability and earned him the Nobel Prize in 1922. His model was a major leap forward in understanding atomic structure.

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Beyond Bohr: The Quantum Mechanical Model

Even Bohr’s model was later found to have limitations. Today, we know electrons don’t follow neat circular paths — they exist in “electron clouds” or probability regions around the nucleus. This is the modern quantum mechanical model, which you’ll study in higher grades!

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Subatomic Particles: Protons, Neutrons & Electrons

An atom is made of three fundamental subatomic particles. Let’s understand each one!

Particle	Symbol	Charge	Location	Relative Mass	Discovered By
Electron (इलेक्ट्रॉन)	e⁻	−1	Shells (outside nucleus)	~1/1836 of proton	J. J. Thomson (1897)
Proton (प्रोटॉन)	p⁺	+1	Inside Nucleus	1 unit	Rutherford
Neutron (न्यूट्रॉन)	n⁰	0 (neutral)	Inside Nucleus	~1 unit (≈ proton)	James Chadwick (1932)

🔍 Discovery of the Neutron (1932)

Scientists noticed a puzzle: Helium has 2 protons, but its mass is about 4 times that of Hydrogen (which has 1 proton). So where is the extra mass coming from?

In 1932, James Chadwick (a student of Rutherford) solved this puzzle by discovering the neutron — a particle with nearly the same mass as a proton but no electrical charge. Neutrons are present in the nucleus of all atoms except hydrogen.

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Why do neutrons exist? The nuclear stability problem!

Protons are all positively charged — they should repel each other in the nucleus! Neutrons help by increasing the distance between protons and by contributing to thenuclear forcethat binds nucleons together. That’s why heavier atoms need more and more neutrons relative to protons.

⚡ Radioactivity — First Clue!

In the late 19th century, scientists discovered that certain elements emit invisible energy and particles — calledradiation. This phenomenon calledradioactivityproved that atoms are NOT indivisible; they contain even smaller particles.

🇮🇳 India’s Atomic Research

BARC (Bhabha Atomic Research Centre), Mumbai, leads advanced neutron-scattering experiments using reactors like Dhruva. Research here has revealed insights into superconductors, battery electrodes, and drug molecules!

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James Chadwick — Nobel Prize in Physics 1935

Working at Cavendish Laboratory, Cambridge, Chadwick discovered the neutron in 1932. This breakthrough explained atomic mass and opened the door to understanding nuclear energy and atomic weapons.

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Atomic Number, Mass Number & Notation

📌 Atomic Number (Z)

The atomic number (परमाणु संख्या) is the number of protons in the nucleus of an atom. It is denoted by the symbol Z.

Atomic number uniquely identifies an element. No two different elements have the same atomic number.
Since atoms are electrically neutral, number of protons = number of electrons.
Example: Hydrogen has 1 proton → Z = 1. Helium has 2 protons → Z = 2.

Atomic Number (Z) = Number of Protons = Number of Electrons (in neutral atom)

⚖️ Mass Number (A)

The mass number (द्रव्यमान संख्या) is the total number of protons + neutrons in the nucleus. It is denoted by the symbol A. Protons and neutrons together are called nucleons.

Mass Number (A) = Number of Protons (p⁺) + Number of Neutrons (n⁰)

Element	Protons (Z)	Neutrons	Mass Number (A)	Electrons
Hydrogen (H)	1	0	1	1
Helium (He)	2	2	4	2
Lithium (Li)	3	4	7	3
Carbon (C)	6	6	12	6
Sodium (Na)	11	12	23	11

📝 Standard Atomic Notation

The symbol, atomic number, and mass number of an element are written in a standard way:

C126↑ Mass Number (A)↑ Atomic Number (Z)← Element Symbol

Standard notation for Carbon: Mass number 12, Atomic number 6

🔍 Finding Neutrons from Atomic & Mass Numbers:

Given: Atomic number (Z) = 17 (Chlorine), Mass number (A) = 35

Number of neutrons = A − Z = 35 − 17

Number of Neutrons = 18

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Common Mistake!

Students often confuse atomic number and mass number. Remember: Atomic number (Z) = protons only. Mass number (A) = protons + neutrons. Number of neutrons = A − Z.

🔤 Symbols of Elements

In 1803, Dalton introduced pictorial symbols. In 1813, Berzelius suggested alphabetical symbols from Latin names. Today, IUPAC approves all element names and symbols.

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Rules for Writing Element Symbols

First letter is ALWAYS uppercase. Second letter (if present) is always lowercase. Examples: Al (not AL), Co (not CO). Some symbols come from Latin names: Iron = Fe (Ferrum), Gold = Au (Aurum), Silver = Ag (Argentum), Sodium = Na (Natrium), Potassium = K (Kalium), Mercury = Hg (Hydrargyros).

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Electronic Configuration: How Electrons Fill Shells

The arrangement of electrons in different energy levels/shells of an atom is called its electronic configuration (इलेक्ट्रॉनिक विन्यास).

📋 Bohr-Bury Rules for Filling Electrons

Maximum electrons in a shell = 2n², where n is the shell number.
K-shell (n=1): max 2 electrons | L-shell (n=2): max 8 electrons | M-shell (n=3): max 18 electrons
The maximum number of electrons in the outermost shell is always 8 (except K-shell which holds only 2).
Electrons fill in order K → L → M → N, moving to the next shell only after the current one is full.

Max electrons in shell = 2n² | K: 2×1²=2 | L: 2×2²=8 | M: 2×3²=18

Element	Symbol	Atomic No. (Z)	K	L	M	Config.
Hydrogen	H	1	1	–	–	1
Helium	He	2	2	–	–	2
Lithium	Li	3	2	1	–	2, 1
Carbon	C	6	2	4	–	2, 4
Neon	Ne	10	2	8	–	2, 8
Sodium	Na	11	2	8	1	2, 8, 1
Magnesium	Mg	12	2	8	2	2, 8, 2
Chlorine	Cl	17	2	8	7	2, 8, 7
Argon	Ar	18	2	8	8	2, 8, 8

📝 Example: Electronic Configuration of Phosphorus (P, Z=15)

Total electrons = 15

K-shell → fills with 2 electrons (max = 2) → K: 2

L-shell → fills with 8 electrons (max = 8) → L: 8

Remaining = 15 − 2 − 8 = 5 electrons go to M-shell

Electronic Configuration of P = 2, 8, 5

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Valency: Combining Capacity of an Atom

🔑 Key Terms

🔵 Valence Shell (संयोजकता कोश)

The outermost shell of an atom containing electrons is called the valence shell. Example: For Sodium (2,8,1), the M-shell with 1 electron is the valence shell.

⚡ Valence Electrons (संयोजकता इलेक्ट्रॉन)

The electrons present in the outermost shell (valence shell) are called valence electrons. They determine chemical properties and reactivity.

✅ The Octet Rule

Atoms are most stable when their outermost shell has 8 electrons (octet) — or 2 electrons in the case of helium (only has K-shell).

Elements with a complete octet (like Ne, Ar) are inert/unreactive (noble gases).
Elements with incomplete valence shells are reactive. They lose, gain, or share electrons to complete the octet.

Valency (संयोजकता) = Number of electrons gained, lost, or shared to complete the octet

Element	Electron Config.	Valence Electrons	Action	Valency
Sodium (Na)	2, 8, 1	1	Loses 1 electron	1
Oxygen (O)	2, 6	6	Gains 2 electrons	2
Carbon (C)	2, 4	4	Shares 4 electrons	4
Chlorine (Cl)	2, 8, 7	7	Gains 1 electron	1
Neon (Ne)	2, 8	8 (complete)	No gain/loss (stable)	0

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Exam Shortcut for Valency!

If valence electrons ≤ 4 → Valency = number of valence electrons (loses them)

If valence electrons > 4 → Valency = 8 − valence electrons (gains electrons)

If valence electrons = 4 → Shares; Valency = 4 (like Carbon)

Noble gases have valence electrons = 8 (or 2 for He) → Valency = 0

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Isotopes & Isobars: Special Atomic Pairs

🔬 Isotopes (समस्थानिक)

Atoms of the same element that have the same atomic number (Z) but different mass numbers (A) — i.e., different numbers of neutrons — are called isotopes.

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NCERT Definition: Isotopes

Isotopes are atoms of the same element with the same number of protons but different numbers of neutrons, and therefore different mass numbers.

💧 Isotopes of Hydrogen

Naturally occurring hydrogen is a mixture of three isotopes:

¹H — Protium (~99.98%)

1 proton, 0 neutrons, 1 electron. Most abundant form of hydrogen. What we use in water (H₂O).

²H — Deuterium (~0.015%)

1 proton, 1 neutron, 1 electron. Used in heavy water (D₂O) and nuclear reactors.

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³H — Tritium (traces)

1 proton, 2 neutrons, 1 electron. Radioactive and found only in traces in nature. Used in nuclear weapons research.

☢️ Important Uses of Isotopes

²³⁵U — Used as fuel in nuclear reactors to generate electricity.
⁶⁰Co — Radioactive cobalt used in cancer radiation therapy.
¹³¹I — Iodine isotope used to treat goitre and thyroid cancer.
¹⁴C — Carbon-14 used in carbon dating — determining the age of fossils and ancient artefacts in archaeology!

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Why do isotopes have SAME chemical properties?

Because chemical properties depend on the number ofvalence electrons. Isotopes have the same atomic number → same number of electrons → same electronic configuration → same chemical properties! Physical properties (like melting/boiling points) differ due to different masses.

⚖️ Average Atomic Mass

Since isotopes exist in nature in different proportions, the atomic mass of an element is calculated as the weighted average of its isotopes based on their abundance.

📐 Example: Weighted Average Atomic Mass of Chlorine

Chlorine has two isotopes: ³⁵Cl (75%) and ³⁷Cl (25%)

Weighted Average = (35 × 75/100) + (37 × 25/100)

= (105/4) + (37/4)

= 142/4

Average Atomic Mass of Cl = 35.5 u

Note: This doesn’t mean any single Cl atom has mass 35.5 u. It’s just a weighted average over millions of atoms!

🔄 Isobars (समभारिक)

Atoms of different elements that have the same mass number (A) but different atomic numbers (Z) are called isobars.

Element	Atomic Number (Z)	Mass Number (A)	Relation
Argon (Ar)	18	40	All are isobars — same mass number (40), different atomic numbers
Potassium (K)	19	40
Calcium (Ca)	20	40

📊 Isotopes vs. Isobars (Quick Comparison)

Isotopes: Same Z (atomic number), Different A (mass number). Same element, different neutrons.

Isobars: Different Z (atomic number), Same A (mass number). Different elements, different elements.

🇮🇳 Dr. Homi Jehangir Bhabha

Known as the Father of India’s Nuclear Programme, Dr. Bhabha founded TIFR and BARC. He pioneered the use of atomic energy in India for electricity, agriculture, and medicine.

📝 Quick Revision Summary

⚛️ Dalton (1808)Atoms are indivisible particles — the building blocks of matter.

⚡ Thomson (1897)Plum pudding model — electrons embedded in positive sphere.

🥇 Rutherford (1911)Nuclear model — dense nucleus at center, electrons orbit outside.

🔵 Bohr (1913)Electrons move in fixed energy shells K, L, M, N. Stability explained.

🔢 Atomic Number (Z)Number of protons. Uniquely identifies an element. Z = electrons.

⚖️ Mass Number (A)A = Protons + Neutrons. Neutrons = A − Z.

🔵 Shell Capacity: 2n²K=2, L=8, M=18. Outermost shell max = 8 electrons.

🔗 ValencyElectrons gained/lost/shared to complete octet. Noble gases: valency = 0.

🧬 IsotopesSame Z, different A. Same chemical properties. Used in medicine, energy, dating.

🔄 IsobarsDifferent Z, same A. Different elements with equal nucleon count.

⚖️ Average Atomic MassWeighted average based on abundance of isotopes in nature.

🧩 Neutron (1932)Chadwick discovered neutron — neutral particle adding mass to nucleus.

📋 Important Exam Questions

Q1. What were the main observations of Rutherford’s gold foil experiment and what conclusions did he draw? (CBSE — 5 Marks)

Observations: (1) Most alpha particles passed straight through the gold foil. (2) A few alpha particles were deflected at large angles. (3) A very few bounced straight back.

Conclusions: (1) Since most particles passed through undeflected, most of an atom is empty space. (2) Since some particles were sharply deflected and a few bounced back, there must be a very small, dense, positively charged nucleus at the center of the atom. (3) Electrons revolve around the nucleus in orbits (planetary model). The nucleus is about 10⁵ times smaller than the atom.

Q2. Why did Rutherford’s model fail to explain atomic stability? How did Bohr’s model solve this problem? (CBSE — 3 Marks)

Rutherford’s failure: According to classical physics, a charged particle moving in a circular path constantly accelerates. An accelerating charged particle continuously loses energy. If electrons lose energy, they would spiral inward and fall into the nucleus — making atoms collapse. But atoms are stable, so this model was incomplete.

Bohr’s solution: Bohr proposed that electrons move in fixed stationary states (shells) where they do NOT lose energy. Energy remains constant in a fixed shell. An electron only gains/loses energy when jumping between shells, releasing a fixed quantum of energy. This explained stability.

Q3. An element X has mass number 35 and 18 neutrons. Find its (i) Atomic number (ii) Electronic configuration (iii) Valency. (CBSE — 3 Marks)

Given: Mass number (A) = 35, Neutrons (n⁰) = 18

(i) Atomic number Z = A − Neutrons = 35 − 18 = 17 → Element X is Chlorine (Cl)

(ii) Total electrons = 17 → Electronic configuration: 2, 8, 7 (K=2, L=8, M=7)

(iii) Valence electrons = 7 → Needs to gain 1 electron to complete octet → Valency = 1

Q4. Define isotopes. Give two examples and state any two uses of isotopes in daily life. (CBSE — 3 Marks)

Definition: Isotopes are atoms of the same element that have the same atomic number (Z) but different mass numbers (A) due to different numbers of neutrons.

Examples: (1) Isotopes of Hydrogen: Protium (¹H), Deuterium (²H), Tritium (³H) — all have Z=1. (2) Isotopes of Carbon: ¹²C, ¹³C, ¹⁴C — all have Z=6.

Uses: (1) ¹⁴C (carbon-14) is used in carbon dating to determine the age of fossils and archaeological artefacts. (2) ²³⁵U (uranium-235) is used as fuel in nuclear reactors to generate electricity.

Q5. Differentiate between Isotopes and Isobars with one example each. (CBSE — 2 Marks)

Isotopes: Atoms of the same element with same atomic number (Z) but different mass numbers (A). They have different numbers of neutrons. Example: ¹⁶O and ¹⁸O — both have Z=8 (oxygen), but different mass numbers.

Isobars: Atoms of different elements with different atomic numbers (Z) but same mass number (A). Example: Argon (Z=18, A=40), Potassium (Z=19, A=40), Calcium (Z=20, A=40) — all have mass number 40 but are different elements.

Notes Class 9 Science Exploration Chapter 8 Journey Inside the Atom

Amresh Academy

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