Atom and the Periodic Table

The Study of Chemistry

• Chemistry involves the study of observable changes in matter to understand their unobservable causes.
• Key focus areas include:
  • Matter ($\text{สสาร}$): Its composition ($\text{องค์ประกอบ}$), properties ($\text{สมบัติ}$), and states ($\text{สถานะ}$).
  • Change ($\text{การเปลี่ยนแปลง}$): Categorized into physical changes and chemical changes.
  • Energy ($\text{พลังงาน}$): The energy associated with transitions and states.
• Example: Brine Water ($\text{brine}$)
  • Mix of Table Salt ($NaCl$) and water ($H_2O$).
  • $NaCl(s) + H_2O(l) \rightarrow NaCl(aq)$
  • Composition includes compounds like Sodum ($Na$), Chlorine ($Cl$), Hydrogen ($H$), and Oxygen ($O$).
  • Brine is a mixture; the individual elements within the compounds belong to specific chemical identities.

Atomic Structure: Discovery and Timeline

• Ancient Greek Era: Early concepts of the atom.
• John Dalton (1803): Proposed the first modern atomic theory.
• J.J. Thomson (1897): Developed the "Plum-pudding" model using the Cathode Ray Tube (CRT). Discovered the electron ($e^-$).
• Robert Millikan (1909): Conducted the Oil Drop experiment to determine the charge of an electron.
• Ernest Rutherford (1911): Proposed the Nuclear model, establishing that the atom has a dense nucleus and mostly empty space.
• Niels Bohr (1913): Proposed circular orbits or shells for electrons with quantized energy levels.
• Erwin Schrödinger (1926): Developed the Modern Quantum model, introducing orbitals and the wave-like behavior of electrons.
• James Chadwick (1932): Discovered the neutron ($n$).

Subatomic Particles and Atomic Symbols

• Atomic Symbol Notation: $_Z^A X_n^{m\pm}$
  • $A$: Mass Number (Total number of protons and neutrons, $\#p + \#n$).
  • $Z$: Atomic Number (Total number of protons, $\#p$).
  • $m\pm$: Ion Charge (The difference between protons and electrons, $\#p - \#e$).
  • $n$: Subscript indicating the number of atoms.
• Properties of Subatomic Particles:
  • Proton ($p$):
    • Mass: $1.673 \times 10^{-27}\,kg$
    • Mass: $1.007\,u$
    • Charge: $+1\,e$ ($e = 1.602 \times 10^{-19}\,C$)
  • Neutron ($n$):
    • Mass: $1.675 \times 10^{-27}\,kg$
    • Mass: $1.009\,u$
    • Charge: $0$
  • Electron ($e$):
    • Mass: $9.109 \times 10^{-31}\,kg$
    • Mass: $\frac{1}{1823}\,u$
    • Charge: $-1\,e$
• Isotopes: Atoms of the same element with the same number of protons ($Z$) but different numbers of neutrons (and thus different mass numbers $A$). Examples given: Carbon-12 ($^{12}C$) and Carbon-13 ($^{13}C$).
• Unit Conversion: $1\,u = 1.661 \times 10^{-24}\,g$.

The First 30 Elements of the Periodic Table

• Main Group Elements (1-20):
  1. $H$ - Hydrogen
  2. $He$ - Helium
  3. $Li$ - Lithium
  4. $Be$ - Beryllium
  5. $B$ - Boron
  6. $C$ - Carbon
  7. $N$ - Nitrogen
  8. $O$ - Oxygen
  9. $F$ - Fluorine
  10. $Ne$ - Neon
  11. $Na$ - Sodium
  12. $Mg$ - Magnesium
  13. $Al$ - Aluminium
  14. $Si$ - Silicon
  15. $P$ - Phosphorus
  16. $S$ - Sulfur
  17. $Cl$ - Chlorine
  18. $Ar$ - Argon
  19. $K$ - Potassium
  20. $Ca$ - Calcium
• First-row Transition Metals (21-30):
  21. $Sc$ - Scandium
  22. $Ti$ - Titanium
  23. $V$ - Vanadium
  24. $Cr$ - Chromium
  25. $Mn$ - Manganese
  26. $Fe$ - Iron
  27. $Co$ - Cobalt
  28. $Ni$ - Nickel
  29. $Cu$ - Copper
  30. $Zn$ - Zinc

Electronic Configuration in Shells

• Bohr’s Model Principles:
  • Energy levels are quantized.
  • Maximum number of electrons in the $n^{th}$ shell is calculated by the formula: $2n^2$.
• Capacity by Shell:
  • $K$ shell ($n=1$): Max $2\,e^-$
  • $L$ shell ($n=2$): Max $8\,e^-$
  • $M$ shell ($n=3$): Max $18\,e^-$
  • $N$ shell ($n=4$): Max $32\,e^-$
• Rules for Filling Shells:
  • Electrons fill the innermost shells first (lowest energy levels).
  • The number of electrons in the outermost (valence) shell cannot exceed $8$.
  • No existing element has a shell containing more than $32\,e^-$.
  • Atoms prefer specific "magic numbers" of electrons for stability.
• Periodic Table Relation:
  • The Number of shells occupied corresponds to the Period (row).
  • The Number of valence electrons corresponds to the Group (column) for main group elements.
• Examples:
  • $_{11}Na$: Configuration $2, 8, 1$ (3 shells, Period 3; 1 valence electron, Group 1).
  • $_{20}Ca$: Configuration $2, 8, 8, 2$ (4 shells, Period 4; 2 valence electrons, Group 2).
  • $_{53}I$: Configuration $2, 8, 18, 18, 7$.
  • $_{82}Pb$: Configuration $2, 8, 18, 32, 18, 4$.

Electronic Configuration in Orbitals

• Quantum Model Principles:
  • Shells are divided into sub-shells known as "orbitals".
  • Each orbital accommodates a maximum of $2\,e^-$ with opposite spins.
• Orbital Types and Capacities:
  • s orbital: 1 orbital per level, max $2\,e^-$.
  • p orbital: 3 orbitals per level, max $6\,e^-$.
  • d orbital: 5 orbitals per level, max $10\,e^-$.
  • f orbital: 7 orbitals per level, max $14\,e^-$.
• General Rule: The $n^{th}$ shell contains $n$ types of orbitals.
• Filling Rules:
  • Fill the lowest-energy orbitals first (Aufbau Principle order).
  • Parallel spins are preferred in degenerate levels (Hund's Rule).
• Orbital Configurations Examples:
  • $_1H$: $1s^1$
  • $_2He$: $1s^2$
  • $_6C$: $1s^2\,2s^2\,2p^2$
  • $_{11}Na$: $1s^2\,2s^2\,2p^6\,3s^1$
  • $_{17}Cl$: $1s^2\,2s^2\,2p^6\,3s^2\,3p^5$
  • $_{20}Ca$: $1s^2\,2s^2\,2p^6\,3s^2\,3p^6\,4s^2$
  • $_{26}Fe$: $1s^2\,2s^2\,2p^6\,3s^2\,3p^6\,4s^2\,3d^6$ (also written as $2, 8, 14, 2$ in shell notation).

Terminology and Classification in the Periodic Table

• Blocks:
  • s-block: Groups 1 and 2, plus Helium.
  • p-block: Groups 13 through 18.
  • d-block: Transition metals (Groups 3 to 12).
  • f-block: Inner Transition Metals (Lanthanides and Actinides).
• Major Group Names:
  • Group 1 (IA): Alkali Metals (excluding Hydrogen).
  • Group 2 (IIA): Alkaline Earth Metals.
  • Group 15 (VA): Pnictogens.
  • Group 16 (VIA): Chalcogens.
  • Group 17 (VIIA): Halogens.
  • Group 18 (VIIIA): Noble Gases / Inert Gases.
• Metallicity:
  • Metals: Located on the left and center.
  • Nonmetals: Located on the upper right side (plus Hydrogen).
  • Metalloids: Border the "staircase" between metals and nonmetals (examples include $B, Si, Ge, As, Sb, Te, Po, At$).

Periodic Trends

• Effective Nuclear Charge ($Z_{eff}$): The net positive charge experienced by an electron in a multi-electron atom. It determines how tightly electrons are held.
• Atomic Radius ($r$): The mean or typical distance from the nucleus to the boundary of the surrounding electron cloud. Radius tends to decrease across a period and increase down a group.
• Ionic Radius ($r_{ion}$): The radius of an atom's ion. Cations (positive ions) are smaller than their neutral atoms; Anions (negative ions) are larger.
• Ionisation Energy ($IE$):
  • Process: $M(g) \rightarrow M(g)^+ + e^-$
  • Definition: The minimum energy required to remove an electron from an atom or molecule in the gaseous state.
  • Trend: Generally increases across a period and decreases down a group.
• Successive Ionization Energies (in $kJ\,mol^{-1}$):
  • $Li: IE1=520; IE2=7298; IE3=11815$
  • $Be: IE1=899; IE2=1757; IE3=14849; IE4=21006$
  • $B: IE1=800; IE2=2427; IE3=3660; IE4=25026$
  • $C: IE1=1086; IE2=2353; IE3=4620; IE4=6222; IE5=37830$
  • $N: IE1=1402, IE2=2856, IE3=4582, IE4=7475, IE5=9445, IE6=53266$
  • $O: IE1=1314, IE2=3388, IE3=5300, IE4=7469, IE5=10989, IE6=13326, IE7=71334$
  • $F: IE1=1681, IE2=3374, IE3=6050, IE4=8408, IE5=11023, IE6=15164, IE7=17868, IE8=92038$
  • $Ne: IE1=2080, IE2=3952, IE3=6122, IE4=9370, IE5=12178, IE6=15238, IE7=19999, IE8=23069, IE9=115379$
  • Note: A massive jump in energy indicates the removal of a "core electron" from a lower shell.
• Electron Affinity ($EA$):
  • Process: $X(g) + e^- \rightarrow X(g)^-$
  • Definition: The energy released when an electron is added to a neutral atom in the gas phase to form a negative ion.
• Electronegativity ($EN$): A measure of the tendency of an atom to attract a bonding pair of electrons. Fluorine ($F$) is the most electronegative element.