Grade 9||Atomic Structure and Chemical Bond|| Notes

1. Introduction to Matter and Elements

• Matter exists as elements (e.g., gold, silver, copper – metals; sulphur, iodine – non-metals) and compounds.
• 118 elements discovered; millions of compounds formed from limited elements.
• Atom: Smallest particle of an element that takes part in chemical reactions.
  • All atoms of same element are identical; different elements have different atoms.
  • Some atoms (inert gases) are chemically inert.

2. Structure of an Atom

• Average diameter of atom: 10⁻¹⁰ m.
• Subatomic particles:

ParticleChargeLocationProtonPositiveNucleusNeutronNeutralNucleusElectronNegativeOrbits/shells around nucleus

• Nucleus contains protons + neutrons; electrons revolve in fixed orbits (shells).

3. Electronic Configuration (Shells: K, L, M, N)

• Maximum electrons in a shell: 2n² (n = shell number).
  • K shell (n=1): 2 electrons
  • L shell (n=2): 8 electrons
  • M shell (n=3): 18 electrons
  • N shell (n=4): 32 electrons

4. Stability Rules

• Duplet rule: K-shell complete with 2 electrons → stable (e.g., He: 2 electrons).
• Octet rule: Outermost shell complete with 8 electrons → stable (e.g., Ne, Ar, Kr, Xe, Rn).
• Inert gases are stable, do not react, found free in nature.

5. Valence Shell, Valence Electrons, and Valency

• Valence shell: Outermost shell.
• Valence electrons: Electrons in valence shell (participate in reactions).
• Valency: Combining capacity = number of electrons an atom can lose, gain, or share to achieve duplet/octet.

Valency Table (Atomic No. 1–20)

ElementSymbolAtomic No.Config (K L M)ValencyHydrogenH111HeliumHe220LithiumLi32,11BerylliumBe42,22BoronB52,33CarbonC62,44NitrogenN72,53OxygenO82,62FluorineF92,71NeonNe102,80SodiumNa112,8,11MagnesiumMg122,8,22AluminiumAl132,8,33SiliconSi142,8,44PhosphorusP152,8,53,5SulphurS162,8,62,6ChlorineCl172,8,71ArgonAr182,8,80PotassiumK192,8,8,11CalciumCa202,8,8,22

• Some elements show variable valency (e.g., Fe: 2,3; Cu: 1,2; Au: 1,3).

6. Ions

• Atoms become charged by losing/gaining electrons → ions.
• Cation (+ve): Formed by loss of electrons (metals) → e.g., Na⁺, Mg²⁺, Al³⁺.
• Anion (–ve): Formed by gain of electrons (non-metals) → e.g., Cl⁻, O²⁻.

7. Chemical Bonds

• Force that holds atoms together in compounds.
• Types studied:
  1. Electrovalent/Ionic Bond: Transfer of electrons (metal + non-metal).
     • Forms ionic compounds (high MP/BP, soluble in water, conduct when dissolved).
     • Examples: NaCl, MgO, CaCl₂.
  2. Covalent Bond: Sharing of electrons (non-metal + non-metal).
     • Forms covalent compounds (low MP/BP, insoluble in water).
     • Single (-), double (=), triple (≡) bonds possible.
     • Examples: H₂O, CH₄, NH₃, HCl.

8. Molecular Formula

• Symbol representation of a molecule (shows types and number of atoms).
• Criss-cross method to write formula:
  • Write symbols → write valencies above → criss-cross → simplify.
  • Example: Aluminium oxide → Al³⁺ O²⁻ → Al₂O₃.

9. Atomic Mass and Molecular Mass

• Atomic mass ≈ protons + neutrons (in amu).
• Molecular mass = sum of atomic masses of all atoms in molecule.

Method of Writing Molecular Formula (Criss-Cross Method)

The textbook explains the criss-cross method as the standard way to write the molecular formula of compounds, especially ionic (electrovalent) compounds. Here are the step-by-step rules as given in the chapter:

Steps:

1. Write the name of the compound (e.g., Magnesium chloride).
2. Write the symbols of the elements/ions side by side (usually metal first). Example: Mg Cl
3. Write the valency of each element/ion above the symbols. Example: Magnesium Chloride 2 1 Mg Cl
4. Criss-cross the valencies (exchange them) and write them as subscripts (below) on the opposite sides. Example: 2 1 Mg Cl ↓ ↓ Mg₁ Cl₂ → MgCl₂ (The subscript 1 is not written.)
5. If the subscript is more than 1 for a polyatomic ion (like SO₄²⁻, HCO₃⁻, etc.), put the ion in brackets () and write the subscript outside. Example: Magnesium bicarbonate Mg²⁺ (HCO₃)⁻ ↓ ↓ Mg₁ (HCO₃)₂ → Mg(HCO₃)₂
6. Simplify if possible (divide subscripts by common factor), but this is rare.

Examples from the Textbook (Page 239)

CompoundSymbols with ValenciesCriss-Cross ResultFinal FormulaSodium chlorideNa¹ Cl¹Na₁Cl₁NaClCalcium chlorideCa² Cl¹Ca₁Cl₂CaCl₂Aluminium chlorideAl³ Cl¹Al₁Cl₃AlCl₃Carbon tetrachlorideC⁴ Cl¹C₁Cl₄CCl₄Boron oxideB³ O²B₂O₃B₂O₃Ammonium sulphateNH₄¹ SO₄²(NH₄)₂SO₄(NH₄)₂SO₄Calcium sulphateCa² SO₄²Ca₂SO₄₂ → divide by 2CaSO₄Aluminium nitrateAl³ NO₃¹Al₁(NO₃)₃Al(NO₃)₃Calcium bicarbonateCa² HCO₃¹Ca₁(HCO₃)₂Ca(HCO₃)₂

Important Points to Remember for Exams

• Valency 1 is never written as subscript.
• Polyatomic ions (e.g., SO₄, NO₃, HCO₃, PO₄, OH, CO₃, NH₄) must be in brackets if subscript > 1.
• For covalent compounds (like H₂O, CH₄, NH₃), we usually write them directly from the number of shared electrons, but criss-cross can still be used where valency applies.

10. Nuclear Stability and Radioactivity

1. Nuclear Stability

• The nucleus of an atom is made up of protons (p⁺) and neutrons (n⁰).
• For the nucleus to be stable, the ratio of neutrons to protons (n/p) should be approximately 1 (or slightly more than 1 for heavier elements).
• If the neutron-to-proton ratio is much more than 1, or if the atomic number is very high, the nucleus becomes unstable.
• Instability increases with increase in atomic number.
• Unstable nuclei try to become stable by emitting energetic radiations → this is called radioactive emission.

Condition for stable nucleus: n/p ≈ 1 (or slightly >1)

2. Radioactive Elements and Radioactive Emission

• Elements with high atomic number (mostly atomic number > 83) have unstable nuclei.
• Such elements are called radioactive elements.
• Examples: Uranium (U), Plutonium (Pu), Thorium, Radium, etc.
• To gain stability, they emit three types of powerful radiations:
  • Alpha (α) particles
  • Beta (β) particles
  • Gamma (γ) rays
• The process of emission of these radiations is called radioactive emission or radioactivity.
• These radiations are harmful to human health.
• Radioactivity was discovered by French scientist Henri Becquerel in 1896 AD.

3. Nuclear Fission

• Definition: The process of splitting a heavy unstable nucleus into two or more lighter nuclei is called nuclear fission.
• It is done artificially by bombarding the heavy nucleus (e.g., Uranium or Plutonium) with high-speed neutrons.
• When the nucleus splits:
  • Smaller/lighter nuclei are formed.
  • Some mass is lost (mass defect).
  • The lost mass is converted into a huge amount of energy (along with α, β, γ radiations).
• (Uranium-236 + neutron → Barium-141 + Krypton-92 + 3 neutrons + energy)
• This is the principle behind atomic bombs and nuclear power plants.

4. Nuclear Fusion

• Definition: The process of combining two or more light nuclei to form a heavier nucleus is called nuclear fusion.
• It requires very high temperature and pressure.
• Example (from textbook):
• (Four hydrogen nuclei → one helium nucleus + energy)
• This process continuously happens in the Sun and stars (reason for their heat and light).
• Abundant hydrogen, extreme heat, and pressure in stars make fusion possible.
• We cannot yet produce controlled nuclear fusion artificially on Earth for energy (very difficult to create required conditions).

5. Atomic Energy (Nuclear Energy)

• Energy obtained from nuclear reactions (mainly fission) is called atomic energy or nuclear energy.
• How electricity is produced (nuclear power plant):
  1. Controlled nuclear fission in a nuclear reactor (kiln) produces huge heat.
  2. Heat boils water → produces steam.
  3. Steam runs a steam engine/turbine.
  4. Turbine runs a generator → produces electricity.
• Used in most developed countries for electricity generation.

Beneficial uses:

• Sterilisation of drinking water, food, medical equipment (using radiations).
• Treatment of diseases like cancer (radiotherapy).

Risks and harmful effects:

• Radiation leakage from reactors.
• Misuse in nuclear weapons (atomic bombs).
• Very dangerous for human health and environment.

Key Points for Exam

• Stable nucleus → n/p ≈ 1
• Radioactive elements → atomic no. > 83 mostly
• Fission → splitting heavy nucleus → used for electricity and bombs
• Fusion → joining light nuclei → happens in Sun
• Atomic energy → boon (electricity, medicine) but also bane (radiation danger)
• Stable nucleus: neutron/proton ratio ≈ 1.
• High atomic number → unstable → radioactive elements (e.g., U, Pu).
• Emit α, β, γ rays.
• Nuclear fission: Heavy nucleus splits → energy (e.g., U → Ba + Kr + energy).
• Nuclear fusion: Light nuclei combine → energy (e.g., 4H → He + energy).
• Atomic energy used for electricity, medicine; harmful if misused.