Atomic Structure & Chemical Bonding – Comprehensive Study Notes

Page 1 – Scope, Fundamental Definitions & Daltonian Foundations

1. Latest Syllabus → Key Topics

• Structure of an atom – sub-atomic particles, atomic number $Z$, mass number $A$, isotopes, octet rule.
• Types of chemical bonding – electrovalent (ionic) vs covalent with full orbital structures.

2. Atom: Concept & Classical Definitions

• Etymology – from Greek “atomos” (indivisible).
• Democritus: qualitative proposal that matter is made of indivisible units.
• John Dalton (1808) – first quantitative atomic theory.

Postulates of Dalton’s Atomic Theory

1. Matter consists of small, indivisible, indestructible particles called atoms.

2. Atoms of the same element are identical; atoms of different elements differ.

3. Atoms neither created nor destroyed in chemical change.

4. Atoms combine in simple whole-number ratios to form molecules.

5. Chemical reactions = rearrangement of atoms.

Modern Corrections to Dalton

3. Discovery of the Electron

• 1878 – William Crookes: Electric discharge through low-pressure gas (≈$0.01$ mm Hg) produced “cathode rays.”
• Characteristics of cathode rays (J. J. Thomson, 1897):
– Travel straight from cathode to anode, cast sharp shadows.
– Deflected by magnetic & electric fields toward a positive plate ⇒ negative charge.
– Possess kinetic energy; heat objects they strike.
• Conclusion: cathode rays are streams of identical, negatively charged particles → electrons.
• Quantitative data:
– Mass $m<em>e = \dfrac{1}{1837}$ mass of H-atom $= 9.107\times10^{-28}\,\text{g}$. – Charge $q</em>e = -1.602\times10^{-19}\,\text{C}$.
• Electrical neutrality of atoms implied presence of equal positive charge → led to discovery of protons.

Page 2 – Protons, Nucleus, Rutherford & Bohr Models, Neutrons

1. Discovery of the Proton

• Eugen Goldstein: observed “positive rays” (anode rays) when cathode was perforated.
• Rays travelled opposite direction to cathode rays; consisted of positively charged particles → protons.
• Data:
– Mass $m<em>p = 1.670\times10^{-24}\,\text{g}$ $\approx1837\ m</em>e$.
– Charge $q_p = +1.602\times10^{-19}\,\text{C}$ (unit positive).

2. Rutherford’s Alpha-Scattering Experiment (1911)

• Alpha particles (He nuclei, $2p+2n$) fired at gold foil (≈$10^{-6}$ cm thick).
• Observations:

1. Majority passed undeflected → atom mostly empty.

2. Some deflected at small angles → encounter with concentrated positive mass.

3. Very few rebounded (>$90^{\circ}$) → existence of tiny dense nucleus.

Rutherford Atomic Model – Postulates

1. Nucleus: minute central region containing entire positive charge & almost all mass.

2. Electrons revolve around nucleus in circular paths; centrifugal force counterbalances electrostatic attraction.

Drawback

• Classical electrodynamics: an accelerating (orbiting) electron should radiate energy → spiral into nucleus → atom unstable. Model failed to explain stability & spectra.

3. Bohr’s Atomic Model (1913)

1. Electrons occupy fixed circular orbits (energy levels) with quantised angular momentum.

2. Principal quantum number $n = 1,2,3,\ldots$ labelled K, L, M, N….

3. Electron neither absorbs nor emits energy while in a given orbit.

4. Transition between levels involves discrete energy $E<em>2-E</em>1 = h\nu$ (line spectra).

5. Provided first explanation of atomic stability & hydrogen spectrum.

4. Discovery of the Neutron (James Chadwick, 1932)

• Mass of He atom (≈$4.003$ amu) > mass of 2 protons; electrons negligible → existence of neutral massive particles.
• Neutron: charge $0$, mass $m_n = 1.676\times10^{-24}\,\text{g}$ (≈mass of proton).
• Protons + neutrons collectively called nucleons.

Page 3 – Atomic Number, Mass Number & Sub-Atomic Accounting

1. Atom & Element

• Atom: smallest unit of an element participating in chemical reactions (e.g., $\text{O},\text{Cl}$).
• Element: pure substance of identical atoms with same $Z$; cannot be decomposed physico-chemically.

2. Definitions & Symbols

• Atomic number $Z$ = number of protons = number of electrons (for neutral atom) = nuclear positive charge.
• Mass number $A$ = total nucleons $= p+n$ (whole-number approximation of atomic mass).
• Standard nuclear notation: $^{A}_{Z}\text{Element}$.

Relationships

\begin{aligned}
Z &= p = e,\
A &= p + n,\
n &= A - Z.
\end{aligned}

3. Sub-Atomic Particle Summary (relative to $^{12}\text{C}$ standard)

Page 4 – Electronic Configuration, 2$n^2$ Rule & First 20 Elements

1. Energy Levels (Shells)

• Electrons revolve in discrete shells K, L, M, N… (n = 1, 2, 3, 4…). Energy increases with distance from nucleus (K < L < M < N).

2. Bohr–Bury Distribution Rules

Rule 1 – Maximum electrons in a shell: $2n^{2}$.
(K $=2$, L $=8$, M $=18$, N $=32$ …)

Rule 2 – Outermost shell never holds >$8$ e⁻; penultimate shell never >$18$ e⁻. A new shell forms once an octet is completed.

3. Electronic Configurations of $Z = 1 \text{ to } 20$ (representative)

• $^{1}<em>{1}\text{H}:\;1$  • $^{4}</em>{2}\text{He}:\;2$  • $^{7}<em>{3}\text{Li}:\;2,1$ • $^{20}</em>{10}\text{Ne}:\;2,8$ (stable octet)
• $^{23}<em>{11}\text{Na}:\;2,8,1$ • $^{40}</em>{20}\text{Ca}:\;2,8,8,2$ (penultimate shell filled to 8, obeying Rule 2).

Diagrams (planetary style) illustrate each shell and valence electrons for all first 20 elements.

Page 5 – Isotopes: Definition, Properties & Examples

1. Definition

Isotopes = atoms of the same element with identical $Z$ but different $A$ (different $n$).

2. Chemical vs Physical Behaviour

• Chemical properties depend on electronic configuration (same for isotopes) → essentially identical.
• Physical properties depend on mass → differ among isotopes (density, diffusion rate, melting point, etc.).

3. Common Isotopic Sets

1. Hydrogen: $^{1}<em>{1}\text{H}$ (Protium), $^{2}</em>{1}\text{H}$ (Deuterium), $^{3}_{1}\text{H}$ (Tritium).

2. Carbon: $^{12}<em>{6}\text{C},\,^{13}</em>{6}\text{C},\,^{14}_{6}\text{C}$ (radio-carbon dating).

3. Oxygen: $^{16}<em>{8}\text{O},\,^{17}</em>{8}\text{O},\,^{18}_{8}\text{O}$.

4. Chlorine: $^{35}<em>{17}\text{Cl}$ & $^{37}</em>{17}\text{Cl}$ – natural abundance 3 : 1, giving atomic mass $35.5$ amu.

5. Potassium ($^{39}<em>{19}\text{K},^{41}</em>{19}\text{K}$), Uranium ($^{235}<em>{92}\text{U},^{238}</em>{92}\text{U}$), etc.

4. Fractional Atomic Mass

Average atomic mass $= \dfrac{\sum (\text{isotopic mass} \times \text{abundance})}{100}$.
Chlorine example:
$\dfrac{(3\times35) + (1\times37)}{4} = 35.5\,\text{amu}.$

Page 6 – Chemical Activity, Octet/Duplet Rule & Bond Formation

1. Noble Gases: Paradigm of Stability

2. Octet Rule

Atoms of other elements are unstable; they undergo chemical change to achieve:
• Duplet (2 e⁻) in K-shell (Helium analogue)
• Octet (8 e⁻) in outer shell (Ne, Ar … analogue)

3. Routes to Stability

1. Electron Transfer (Ionic/Electrovalent Bond)
   – Metal loses e⁻ → cation; non-metal gains e⁻ → anion; electrostatic attraction forms ionic lattice.
   – Examples:
   • $^{23}<em>{11}\text{Na} \rightarrow \text{Na}^{+} + e^{-}$ ($2,8,1 \Rightarrow 2,8$) • $^{35}</em>{17}\text{Cl} + e^{-} \rightarrow \text{Cl}^{-}$ ($2,8,7 \Rightarrow 2,8,8$)
   • $\text{Na}^{+} + \text{Cl}^{-} \rightarrow \text{NaCl}$.
   • $\text{Ca}^{2+} + \text{O}^{2-} \rightarrow \text{CaO}$.

2. Electron Sharing (Covalent Bond)
   – Two atoms share one or more pairs of electrons to complete octet/duplet.
   – Examples (orbital/Lewis structures):
   • $\text{H}<em>2$: $H\cdot + \cdot H \Rightarrow H:H$. • $\text{O}</em>2$: $\cdot\cdot O::O \cdot\cdot$ (double bond).
   • $\text{N}<em>2$: triple bond $N\equiv N$. • $\text{H}</em>2\text{O}$: $H-O-H$ (2 lone pairs on O).
   • $\text{NH}<em>3$: $H-N-H$ with one lone pair. • $\text{CH}</em>4$, $\text{CCl}_4$ illustrate tetra-covalency of carbon.

4. Electronic Theory of Redox (Linkage)

• Oxidation = loss of electrons (e.g., $\text{Fe}^{2+} - e^{-} \rightarrow \text{Fe}^{3+}$).
• Reduction = gain of electrons (e.g., $\text{Fe}^{3+} + e^{-} \rightarrow \text{Fe}^{2+}$).

Ethical, Philosophical & Real-World Notes

• Evolution of atomic models shows scientific method: hypotheses → experimental evidence → model refinement.
• Isotopes underpin radiometric dating, medical tracers ((^{131}\text{I}), (^{99m}\text{Tc})), and nuclear power ((^{235}\text{U})).
• Octet concept rationalises periodicity: valence e⁻ configuration governs group chemistry.
• Stability considerations of nuclei (balance of protons and neutrons) inform nuclear medicine and environmental safety.