Atomic Structure and Chemical Bonding Study Notes (copy)

Historical Development and Atomic Theories

• Democritus (Ancient Greek Philosopher):

  • Studied the nature of matter and proposed the concept of basic constituents of all substances.

• John Dalton (1808):

  • Proposed the Atomic Theory of Matter.

• Postulates of Dalton's Atomic Theory:

  • Matter consists of extremely small, indivisible particles called atoms.

  • Atoms can neither be created nor destroyed in a chemical process.

  • Atoms of the same element are identical in mass, size, and all other properties, while atoms of different elements differ in their properties.

  • Atoms combine with other atoms in simple, whole-number ratios to form compound atoms or molecules.

  • An atom is the smallest unit of matter that takes part in a chemical reaction; all chemical changes result from the combination or separation of atoms.

• Comparison between Dalton's Atomic Theory and Modern Atomic Theory:

  • Indivisibility:

    • Dalton's Theory: Stated that atoms are indivisible and indestructible.

    • Modern Atomic Theory: Proves that atoms are divisible and destructible, consisting of subatomic particles including electrons, protons, and neutrons.

  • Identity of Atoms of the Same Element:

    • Dalton's Theory: Stated that all atoms of a given element are identical in all respects.

    • Modern Atomic Theory: Shows that atoms of the same element may not be identical in mass, as demonstrated by isotopes (atoms of the same element with the same atomic number but different mass numbers).

  • Points of Agreement / Correlation:

    • Atoms remain the smallest units of matter that participate directly in chemical reactions.

    • In a given compound, the relative number and kind of atoms remain constant.

Discovery of Subatomic Particles and Atomic Models

• Discovery of Electrons (Cathode Rays):

  • William Crookes (1878):

    • Discovered cathode rays using a discharge tube.

    • Observed that when an electric discharge is passed through a gas at extremely low pressure ($0.01\,mm\text{ of }Hg$), blue rays are emitted from the negative electrode (cathode).

    • These rays were named cathode rays.

  • Sir J. J. Thomson (1897):

    • Investigated the constituents and characteristics of cathode rays.

    • Characteristics of Cathode Rays:

      1. Travel in straight lines from the cathode to the anode, casting a sharp shadow of any opaque object placed in their path.

      2. Are deflected by electric and magnetic fields toward the positively charged plate, proving they carry negative charge.

      3. Possess kinetic energy and can raise the temperature of a thin metallic foil upon collision.

    • Conclusion: Cathode rays consist of negatively charged particles named electrons.

    • Electrons produced from any gas or electrode material are completely identical, carrying a unit negative charge and negligible mass.

  • Properties of an Electron:

    • Mass: $\frac{1}{1837}$ of the mass of a hydrogen atom (or proton), equal to $9.107 \times 10^{-28}\,g$

    • Charge: Unit negative charge, equal to $1.602 \times 10^{-19}\,C$

• Discovery of Protons (Positive Rays / Anode Rays):

  • Goldstein:

    • Discovered positive rays (anode rays) emitted from the anode using a discharge tube with a perforated cathode disc.

    • Observed stream of positively charged particles traveling in the direction opposite to cathode rays.

  • Properties of a Proton:

    • Mass: $1837$ times as heavy as an electron, equal to $1.670 \times 10^{-24}\,g$

    • Charge: Unit positive charge, equal to $1.602 \times 10^{-19}\,C$

• Discovery of the Atomic Nucleus:

  • Lord Rutherford (1911):

    • Directed positively charged alpha particles ($\alpha$-particles, $He^{2+}$ containing $2$ protons and $2$ neutrons) from a radium source toward a thin gold foil (thickness of approximately one millionth of a cm, i.e., $10^{-6}\,cm$).

  • Observations:

    1. Most of the alpha particles passed straight through the gold foil without any deflection.

    2. A small fraction of alpha particles were deflected slightly from their original path.

    3. A very small number of alpha particles were deflected by large angles or bounced back completely ($360^\circ$ trajectory deflection).

  • Conclusions:

    1. Most of the volume inside an atom is empty space.

    2. All the positive charge and nearly all the mass of the atom are concentrated in an extremely small region called the nucleus.

    3. Alpha particles passing close to this concentrated positive mass experience strong repulsion and deflect; those approaching head-on rebound backward.

  • Postulates of Rutherford's Atomic Model:

    • An atom consists of a dense, positively charged center called the nucleus.

    • Electrons revolve around the nucleus at high speeds in circular paths to balance the electrostatic attraction of the nucleus via centrifugal force.

  • Drawback of Rutherford's Model:

    • According to classical electromagnetic theory, a revolving charged particle must continuously emit electromagnetic radiation.

    • As the electron loses energy, its orbit would shrink in a spiral path until it ultimately collapses into the nucleus. Thus, Rutherford's model could not explain the stability of an atom.

• Bohr's Atomic Model:

  • Postulates:

    1. Electrons revolve around the nucleus only in certain stable, non-radiating circular paths called fixed orbits or energy levels.

    2. Energy levels are denoted by integers $n = 1, 2, 3, 4, \dots$ starting from the innermost shell, or by letters $K, L, M, N, \dots$

    3. As long as an electron remains in a specific orbit, it neither absorbs nor radiates energy.

    4. An electron moves to a higher energy level upon absorbing energy and drops to a lower energy level upon emitting energy.

    5. The energy difference between two orbits is given by the relation:            $E_2 - E_1 = h\nu$            where $E_2$ is the higher energy state, $E_1$ is the lower energy state, $h$ is Planck's constant, and $\nu$ is the frequency of radiated energy.

• Discovery of Neutrons:

  • James Chadwick (1932):

    • Investigated the mass anomaly of atoms: Nitrogen and Helium nuclei contained more mass than predicted solely by their proton count (e.g., Helium has $2$ protons with expected mass $2 \times 1.008\,a.m.u.$, but experimental mass is $4.003\,a.m.u.$).

    • Proved the existence of neutral subatomic particles located in the nucleus with mass nearly equal to that of a proton.

    • Protons and neutrons together are termed nucleons.

  • Properties of a Neutron:

    • Mass: $1.676 \times 10^{-24}\,g$ (approximately $1\,a.m.u.$, virtually equal to a proton)

    • Charge: Electrically neutral ($0$ charge)

Representation of Atoms, Subatomic Structure, and Electronic Configuration

• Basic Definitions:

  • Atom: The smallest particle of an element that can take part in a chemical reaction and represents the basic structural unit of matter (e.g., $O$, $Cl$).

  • Element: A pure substance made up of only one kind of atom, having the same atomic number, which cannot be broken down into simpler substances by physical or chemical means.

• Summary of Subatomic Particles:

  • Proton (${}_1^1 p$):

    • Charge: $+1$ ($+1.602 \times 10^{-19}\,C$)

    • Mass: $1\,a.m.u.$ ($1.670 \times 10^{-24}\,g$)

    • Location: Nucleus

  • Electron (${}_{-1}^0 e$):

    • Charge: $-1$ ($-1.602 \times 10^{-19}\,C$)

    • Mass: Negligible / $\frac{1}{1837}\,a.m.u.$ ($9.107 \times 10^{-28}\,g$)

    • Location: Shells/Orbits surrounding the nucleus

  • Neutron (${}_0^1 n$):

    • Charge: $0$

    • Mass: $1\,a.m.u.$ ($1.676 \times 10^{-24}\,g$)

    • Location: Nucleus

• Atomic Number ($Z$):

  • The total number of protons present in the nucleus of an atom.

  • For a neutral atom, Atomic Number ($Z$) = Number of Protons ($p$) = Number of Electrons ($e$).

  • Determines the total magnitude of positive nuclear charge.

• Mass Number ($A$):

  • The total number of protons and neutrons (nucleons) present in the nucleus of an atom.

  • Relationship formula: $A = p + n = Z + n$

  • Number of neutrons ($n$) = $A - Z$

  • Standard atomic symbol notation: ${}_Z^A X$

  • Worked Example: For ${}_9^{19} X$

    • Atomic Number $Z = 9$

    • Mass Number $A = 19$

    • Protons $p = 9$

    • Electrons $e = 9$

    • Neutrons $n = 19 - 9 = 10$

• Arrangement and Distribution of Electrons (Bohr-Bury Scheme):

  • Electrons occupy distinct energy levels named $K, L, M, N, \dots$

  • Electrons in the $K$ shell possess minimum energy; energy increases progressively as distance from the nucleus increases (K < L < M < N).

  • Rule 1 (Maximum Shell Capacity):

    • The maximum number of electrons in a given shell is determined by the formula $2n^2$, where $n$ is the shell number counting outward from the nucleus.

    • K\text{-shell } (n=1)$: 2 \times (1)^2 = 2\text{ electrons}\n * L\text{-shell } (n=2)$: $2 \times (2)^2 = 8\text{ electrons}$

    • M\text{-shell } (n=3)$: 2 \times (3)^2 = 18\text{ electrons}\n * N\text{-shell } (n=4)$: $2 \times (4)^2 = 32\text{ electrons}$

  • Rule 2 (Valence Shell Constraints):

    • The outermost (valence) shell cannot accommodate more than $8$ electrons.

    • The penultimate shell (second to last) cannot accommodate more than $18$ electrons.

    • A new outermost shell begins forming as soon as the previous outer shell attains $8$ electrons.

  • Worked Example: Calcium (${}_{20}^{40}Ca$):

    • $Z = 20, A = 40 \implies p = 20, e = 20, n = 20$

    • Applying $2n^2$ limits directly yields $2, 8, 10$, but the outermost shell cannot hold more than $8$ electrons.

    • Correct Electronic Configuration: $2, 8, 8, 2$ ($K=2, L=8, M=8, N=2$).

Atomic Structure and Properties of the First 20 Elements

• Hydrogen ($H$):

  • Atomic Number ($Z$): $1$

  • Mass Number ($A$): $1$

  • Protons: $1$, Electrons: $1$, Neutrons: $0$

  • Electronic Configuration: $1$

  • Valency: $\pm 1$

  • Classification: Metal / Nonmetal

• Helium ($He$):

  • Atomic Number ($Z$): $2$

  • Mass Number ($A$): $4$

  • Protons: $2$, Electrons: $2$, Neutrons: $2$

  • Electronic Configuration: $2$

  • Valency: $0$

  • Classification: Noble gas

• Lithium ($Li$):

  • Atomic Number ($Z$): $3$

  • Mass Number ($A$): $7$

  • Protons: $3$, Electrons: $3$, Neutrons: $4$

  • Electronic Configuration: $2, 1$

  • Valency: $+1$

  • Classification: Metal

• Beryllium ($Be$):

  • Atomic Number ($Z$): $4$

  • Mass Number ($A$): $9$

  • Protons: $4$, Electrons: $4$, Neutrons: $5$

  • Electronic Configuration: $2, 2$

  • Valency: $+2$

  • Classification: Metal

• Boron ($B$):

  • Atomic Number ($Z$): $5$

  • Mass Number ($A$): $11$

  • Protons: $5$, Electrons: $5$, Neutrons: $6$

  • Electronic Configuration: $2, 3$

  • Valency: $+3$

  • Classification: Metalloid

• Carbon ($C$):

  • Atomic Number ($Z$): $6$

  • Mass Number ($A$): $12$

  • Protons: $6$, Electrons: $6$, Neutrons: $6$

  • Electronic Configuration: $2, 4$

  • Valency: $4$

  • Classification: Nonmetal solid

• Nitrogen ($N$):

  • Atomic Number ($Z$): $7$

  • Mass Number ($A$): $14$

  • Protons: $7$, Electrons: $7$, Neutrons: $7$

  • Electronic Configuration: $2, 5$

  • Valency: $-3$ (also $3, 5$)

  • Classification: Nonmetal gas

• Oxygen ($O$):

  • Atomic Number ($Z$): $8$

  • Mass Number ($A$): $16$

  • Protons: $8$, Electrons: $8$, Neutrons: $8$

  • Electronic Configuration: $2, 6$

  • Valency: $-2$

  • Classification: Nonmetal gas

• Fluorine ($F$):

  • Atomic Number ($Z$): $9$

  • Mass Number ($A$): $19$

  • Protons: $9$, Electrons: $9$, Neutrons: $10$

  • Electronic Configuration: $2, 7$

  • Valency: $-1$

  • Classification: Nonmetal gas

• Neon ($Ne$):

  • Atomic Number ($Z$): $10$

  • Mass Number ($A$): $20$

  • Protons: $10$, Electrons: $10$, Neutrons: $10$

  • Electronic Configuration: $2, 8$

  • Valency: $0$

  • Classification: Noble gas

• Sodium ($Na$):

  • Atomic Number ($Z$): $11$

  • Mass Number ($A$): $23$

  • Protons: $11$, Electrons: $11$, Neutrons: $12$

  • Electronic Configuration: $2, 8, 1$

  • Valency: $+1$

  • Classification: Metal

• Magnesium ($Mg$):

  • Atomic Number ($Z$): $12$

  • Mass Number ($A$): $24$

  • Protons: $12$, Electrons: $12$, Neutrons: $12$

  • Electronic Configuration: $2, 8, 2$

  • Valency: $+2$

  • Classification: Metal

• Aluminium ($Al$):

  • Atomic Number ($Z$): $13$

  • Mass Number ($A$): $27$

  • Protons: $13$, Electrons: $13$, Neutrons: $14$

  • Electronic Configuration: $2, 8, 3$

  • Valency: $+3$

  • Classification: Metal

• Silicon ($Si$):

  • Atomic Number ($Z$): $14$

  • Mass Number ($A$): $28$

  • Protons: $14$, Electrons: $14$, Neutrons: $14$

  • Electronic Configuration: $2, 8, 4$

  • Valency: $4$

  • Classification: Metalloid

• Phosphorus ($P$):

  • Atomic Number ($Z$): $15$

  • Mass Number ($A$): $31$

  • Protons: $15$, Electrons: $15$, Neutrons: $16$

  • Electronic Configuration: $2, 8, 5$

  • Valency: $-3$ (also $3, 5$)

  • Classification: Nonmetal

• Sulphur ($S$):

  • Atomic Number ($Z$): $16$

  • Mass Number ($A$): $32$

  • Protons: $16$, Electrons: $16$, Neutrons: $16$

  • Electronic Configuration: $2, 8, 6$

  • Valency: $-2$

  • Classification: Nonmetal

• Chlorine ($Cl$):

  • Atomic Number ($Z$): $17$

  • Mass Number ($A$): $35$

  • Protons: $17$, Electrons: $17$, Neutrons: $18$

  • Electronic Configuration: $2, 8, 7$

  • Valency: $-1$

  • Classification: Nonmetal

• Argon ($Ar$):

  • Atomic Number ($Z$): $18$

  • Mass Number ($A$): $40$

  • Protons: $18$, Electrons: $18$, Neutrons: $22$

  • Electronic Configuration: $2, 8, 8$

  • Valency: $0$

  • Classification: Noble gas

• Potassium ($K$):

  • Atomic Number ($Z$): $19$

  • Mass Number ($A$): $39$

  • Protons: $19$, Electrons: $19$, Neutrons: $20$

  • Electronic Configuration: $2, 8, 8, 1$

  • Valency: $+1$

  • Classification: Metal

• Calcium ($Ca$):

  • Atomic Number ($Z$): $20$

  • Mass Number ($A$): $40$

  • Protons: $20$, Electrons: $20$, Neutrons: $20$

  • Electronic Configuration: $2, 8, 8, 2$

  • Valency: $+2$

  • Classification: Metal

Isotopes and Fractional Atomic Weight Calculation

• Definition of Isotopes:

  • Atoms of the same element having the same atomic number ($Z$) but different mass numbers ($A$).

  • Isotopes possess the same number of protons and electrons, but differ in the number of neutrons.

• Properties of Isotopes:

  • Chemical Properties:

    • Identical/similar among isotopes of the same element.

    • Reason: Chemical properties depend on the electronic configuration and atomic number ($Z$), which are identical for all isotopes of an element.

  • Physical Properties:

    • Different among isotopes of the same element.

    • Reason: Physical constants depend on mass number ($A$) and atomic mass, which differ due to different neutron counts.

• Examples of Isotopes:

  • Hydrogen (3 isotopes):

    • Protium (${}_1^1 H$): $p = 1, e = 1, n = 0$

    • Deuterium (${}_1^2 H$): $p = 1, e = 1, n = 1$

    • Tritium (${}_1^3 H$): $p = 1, e = 1, n = 2$

  • Carbon (3 isotopes):

    • ${}_6^{12} C$: $p = 6, e = 6, n = 6$

    • ${}_6^{13} C$: $p = 6, e = 6, n = 7$

    • ${}_6^{14} C$: $p = 6, e = 6, n = 8$

  • Oxygen (3 isotopes):

    • ${}_8^{16} O$: $p = 8, e = 8, n = 8$

    • ${}_8^{17} O$: $p = 8, e = 8, n = 9$

    • ${}_8^{18} O$: $p = 8, e = 8, n = 10$

  • Chlorine (2 isotopes):

    • ${}_{17}^{35} Cl$: $p = 17, e = 17, n = 18$

    • ${}_{17}^{37} Cl$: $p = 17, e = 17, n = 20$

  • Potassium (2 isotopes):

    • ${}_{19}^{39} K$, ${}_{19}^{41} K$

  • Uranium (2 isotopes):

    • ${}_{92}^{235} U$, ${}_{92}^{238} U$

• Fractional Atomic Mass of Chlorine:

  • Atomic mass is the ratio of the average mass of all naturally occurring isotopes of an element relative to $\frac{1}{12}\text{th}$ the mass of a ${}_6^{12} C$ atom.

  • Chlorine exists as a mixture of two isotopes, ${}_{17}^{35} Cl$ and ${}_{17}^{37} Cl$, in a natural abundance ratio of approximately $3 : 1$ (i.e., 3 out of 4 atoms are ${}_{17}^{35} Cl$ and 1 out of 4 atoms is ${}_{17}^{37} Cl$).

  • Calculation:

    • Mass of three ${}_{17}^{35} Cl$ atoms = $3 \times 35 = 105$

    • Mass of one ${}_{17}^{37} Cl$ atom = $1 \times 37 = 37$

    • Total mass of 4 atoms = $105 + 37 = 142$

    • Average Atomic Mass = $\frac{142}{4} = 35.5\,a.m.u.$

Octet Rule, Chemical Activity, and Electron Redistribution

• Electronic Configuration of Noble Gases:

  • Noble gases have fully filled outermost shells, making them chemically inert and unreactive.

  • Helium ($He$, $Z=2$): Electronic configuration $2$ (K shell complete; Duplet arrangement).

  • Neon ($Ne$, $Z=10$): Electronic configuration $2, 8$ (Octet arrangement).

  • Argon ($Ar$, $Z=18$): Electronic configuration $2, 8, 8$ (Octet arrangement).

  • Krypton ($Kr$, $Z=36$): Electronic configuration $2, 8, 18, 8$

  • Xenon ($Xe$, $Z=54$): Electronic configuration $2, 8, 18, 18, 8$

  • Radon ($Rn$, $Z=86$): Electronic configuration $2, 8, 18, 32, 18, 8$

• Driving Force for Chemical Activity:

  • Atoms of elements other than noble gases possess unstable, incomplete outer electron shells.

  • The primary driving force for chemical combination is the tendency of atoms to attain a stable electronic configuration corresponding to the nearest noble gas.

• Duplet and Octet Rules:

  • Duplet Rule: Atoms tend to achieve a stable electronic structure containing $2$ valence electrons in the first shell (like Helium).

  • Octet Rule: Atoms tend to achieve a stable electronic structure containing $8$ valence electrons in their outermost shell (like Neon, Argon, etc.).

• Modes of Achieving Stability:

  1. Transfer of Valence Electrons (Metal to Non-metal):

     • Forms positively charged cations and negatively charged anions.

     • Oxidation: Electronic process where an atom or ion loses electrons (e.g., $Fe^{2+} - e^- \rightarrow Fe^{3+}$).

     • Reduction: Electronic process where an atom or ion gains electrons (e.g., $Fe^{3+} + e^- \rightarrow Fe^{2+}$).

  2. Sharing of Valence Electrons (Non-metal to Non-metal):

     • Involves mutual sharing of electron pairs between two combining non-metallic atoms.

Electrovalent (Ionic) Bonding

• Definitions:

  • Electrovalent Bond: The chemical bond formed by the complete transfer of one or more valence electrons from a metallic atom to a non-metallic atom.

  • Electrovalent Compound: The compound formed as a result of electron transfer between metallic and non-metallic elements.

• Formation of Sodium Chloride ($NaCl$):

  • Sodium Atom (${}_{11}^{23} Na$):

    • Electronic configuration: $2, 8, 1$

    • Nearest noble gas: Neon ($2, 8$)

    • Loses $1$ valence electron to achieve stability, forming a sodium cation: $Na^+$

  • Chlorine Atom (${}_{17}^{35} Cl$):

    • Electronic configuration: $2, 8, 7$

    • Nearest noble gas: Argon ($2, 8, 8$)

    • Gains $1$ electron into its outer shell, forming a chloride anion: $Cl^-$

  • Reaction:         $Na\,(2,8,1) + Cl\,(2,8,7) \rightarrow [Na]^+\,(2,8) + [Cl]^-\,(2,8,8) \rightarrow NaCl$

  • Electrostatic forces of attraction bind the oppositely charged $Na^+$ and $Cl^-$ ions together in an electrovalent bond.

• Formation of Calcium Oxide ($CaO$):

  • Calcium Atom (${}_{20}^{40} Ca$):

    • Electronic configuration: $2, 8, 8, 2$

    • Nearest noble gas: Argon ($2, 8, 8$)

    • Loses $2$ valence electrons, forming a calcium cation: $Ca^{2+}$

  • Oxygen Atom (${}_8^{16} O$):

    • Electronic configuration: $2, 6$

    • Nearest noble gas: Neon ($2, 8$)

    • Gains $2$ electrons into its outer shell, forming an oxide anion: $O^{2-}$

  • Reaction:         $Ca\,(2,8,8,2) + O\,(2,6) \rightarrow [Ca]^{2+}\,(2,8,8) + [O]^{2-}\,(2,8) \rightarrow CaO$

  • Electrostatic forces bind $Ca^{2+}$ and $O^{2-}$ together in an electrovalent bond.

• Formation of Magnesium Chloride ($MgCl_2$):

  • Magnesium Atom (${}_{12}^{24} Mg$):

    • Electronic configuration: $2, 8, 2$

    • Nearest noble gas: Neon ($2, 8$)

    • Loses $2$ valence electrons, forming a magnesium cation: $Mg^{2+}$

  • Chlorine Atoms ($2$ atoms of ${}_{17}^{35} Cl$):

    • Electronic configuration: $2, 8, 7$

    • Each chlorine atom accepts $1$ electron to achieve Argon configuration ($2, 8, 8$), forming two chloride anions: $2\,Cl^-$

  • Reaction:         $Mg\,(2,8,2) + 2\,Cl\,(2,8,7) \rightarrow [Mg]^{2+}\,(2,8) + 2\,[Cl]^-\,(2,8,8) \rightarrow MgCl_2$

Covalent Bonding and Molecular Structures

• Definitions:

  • Covalent Bond: The chemical bond formed by the mutual sharing of one or more electron pairs between two non-metallic atoms.

  • Covalent Compound: The compound formed due to mutual electron sharing between non-metallic elements.

• Single Covalent Bond Formations:

  • Hydrogen Molecule ($H_2$):

    • Two hydrogen atoms ($H\,(1)$) each need $1$ electron to complete the Helium duplet ($2$).

    • Share $1$ pair of electrons to form a single covalent bond: $H - H$

  • Chlorine Molecule ($Cl_2$):

    • Two chlorine atoms ($Cl\,(2,8,7)$) each need $1$ electron to complete the Argon octet ($2,8,8$).

    • Share $1$ pair of electrons to form a single covalent bond: $Cl - Cl$

  • Hydrogen Chloride Molecule ($HCl$):

    • One hydrogen atom ($H\,(1)$) and one chlorine atom ($Cl\,(2,8,7)$) share $1$ pair of electrons.

    • Hydrogen completes its duplet ($2$) and chlorine completes its octet ($2,8,8$): $H - Cl$

• Multiple Covalent Bond Formations:

  • Oxygen Molecule ($O_2$):

    • Two oxygen atoms ($O\,(2,6)$) each need $2$ electrons to complete the Neon octet ($2,8$).

    • Share $2$ pairs of electrons to form a double covalent bond: $O = O$

  • Nitrogen Molecule ($N_2$):

    • Two nitrogen atoms ($N\,(2,5)$) each need $3$ electrons to complete the Neon octet ($2,8$).

    • Share $3$ pairs of electrons to form a triple covalent bond: $N \equiv N$

• Polyatomic Covalent Formations:

  • Water Molecule ($H_2O$):

    • One oxygen atom ($O\,(2,6)$) shares $1$ electron pair each with two separate hydrogen atoms ($H\,(1)$).

    • Hydrogen achieves stable duplet ($2$) and oxygen achieves stable octet ($2,8$).

    • Structure: $H - O - H$ (contains two shared pairs and two lone pairs of non-bonding electrons on oxygen).

  • Ammonia Molecule ($NH_3$):

    • One nitrogen atom ($N\,(2,5)$) shares $1$ electron pair each with three separate hydrogen atoms ($H\,(1)$).

    • Hydrogen achieves stable duplet ($2$) and nitrogen achieves stable octet ($2,8$).

    • Structure contains three shared pairs and one lone pair of non-bonding electrons on nitrogen.

  • Carbon Tetrachloride Molecule ($CCl_4$):

    • One carbon atom ($C\,(2,4)$) shares $1$ electron pair each with four separate chlorine atoms ($Cl\,(2,8,7)$).

    • Carbon achieves stable octet ($2,8$) and each chlorine achieves stable octet ($2,8,8$).

    • Structure consists of a central carbon atom with four single covalent bonds to chlorine atoms.

  • Methane Molecule ($CH_4$):

    • One carbon atom ($C\,(2,4)$) shares $1$ electron pair each with four separate hydrogen atoms ($H\,(1)$).

    • Carbon achieves stable octet ($2,8$) and each hydrogen achieves stable duplet ($2$).

    • Structure consists of a central carbon atom with four single covalent bonds to hydrogen atoms.

Radical Formation via Covalent Bond Fission

• Process of Radical Formation:

  • Occurs when a covalent bond breaks symmetrically such that each departing atom retains one of the two shared electrons.

  • This symmetrical cleavage process is termed symmetrical bond fission or homolytic fission.

• Initiating Factors:

  • Bond fission is initiated by light energy ($h\nu$), heat, or specific chemical agents.

• General Reaction Equation:     $A - B \xrightarrow{h\nu} A^\bullet + B^\bullet$     where $h\nu$ represents light energy, and $A^\bullet$ and $B^\bullet$ represent neutral free radicals carrying an unpaired electron.