Classification of Elements and Periodicity in Properties Study Notes

The Importance and Purpose of the Periodic Table

• The Periodic Table is cited by Glenn T. Seaborg as the most important concept in chemistry, both in principle and in practice.
• It serves as everyday support for students and suggests research avenues for professionals by providing a succinct organization of the chemical sciences.
• It demonstrates that elements are not random entities but display trends and form families.

The Need for Chemical Classification

• Elements are the basic units of matter. The history of their discovery shows a rapid increase in numbers:
  • 1800: Only 31 elements were known.
  • 1865: Identified elements more than doubled to 63.
  • Present: 114 elements are known (recently discovered ones are man-made).
• Rationale for Classification: Individual study of over 100 elements and their compounds is impractical. Periodic classification rationalizes known chemical facts and facilitates the prediction of new ones for future study.

Historical Genesis of Periodic Classification

• Johann Dobereiner (Early 1800s): The first to consider trends in properties.
  • Law of Triads (1829): He noted groups of three elements (Triads) where the middle element had an atomic weight approximately halfway between the other two.
  • Examples of Triads:
    1. $Li$ (7), $Na$ (23), $K$ (39)
    2. $Ca$ (40), $Sr$ (88), $Ba$ (137)
    3. $Cl$ (35.5), $Br$ (80), $I$ (127)
  • Dismissal: The relationship worked for only a few elements and was dismissed as coincidence.
• A.E.B. de Chancourtois (1862): A French geologist who arranged elements in order of increasing atomic weights in a cylindrical table to show periodic recurrence. It did not attract significant attention.
• John Alexander Newlands (1865): Proposed the Law of Octaves.
  • He arranged elements by increasing atomic weight and noted every eighth element had properties similar to the first (like musical octaves).
  • Limitations: True only for elements up to Calcium ($Ca$).
  • Recognition: Despite initial lack of acceptance, he was awarded the Davy Medal in 1887 by the Royal Society, London.
• Lothar Meyer (1830–1895) and Dmitri Mendeleev (1834–1907): Working independently in 1869, both proposed that properties are periodic functions of atomic weights.
  • Lothar Meyer: Plotted physical properties (atomic volume, melting point, boiling point) against atomic weight, revealing periodic patterns and changes in repeating pattern lengths. His table (1868) closely resembled the modern one but was published after Mendeleev's.
  • Dmitri Mendeleev: Generally credited with the Periodic Table. He published the Periodic Law first: "The properties of the elements are a periodic function of their atomic weights."

Features of Mendeleev’s Periodic Table

• Organization: Arranged elements in horizontal rows (series) and vertical columns (groups) based on increasing atomic weight.
• Prioritizing Properties: Mendeleev ignored strict atomic weight order if it meant placing an element in the wrong group.
  • Example: Iodine ($126.90$) was placed in Group VII after Tellurium ($127.60$) because its properties matched Fluorine, Chlorine, and Bromine.
• Predictions of Undiscovered Elements: He left gaps for unknown elements, calling them Eka-Aluminium and Eka-Silicon. When Gallium and Germanium were discovered, their experimental properties matched his predictions almost perfectly.
• Comparison Table (Predicted vs. Found):
  • Eka-aluminium (Predicted): Atomic weight $68$, Density $5.9\,g/cm^3$, Formula of oxide $E_2O_3$, Formula of chloride $ECl_3$.
  • Gallium (Found): Atomic weight $70$, Density $5.94\,g/cm^3$, Formula of oxide $Ga_2O_3$, Formula of chloride $GaCl_3$.
  • Eka-silicon (Predicted): Atomic weight $72$, Density $5.5\,g/cm^3$, Formula of oxide $EO_2$, Formula of chloride $ECl_4$.
  • Germanium (Found): Atomic weight $72.6$, Density $5.36\,g/cm^3$, Formula of oxide $GeO_2$, Formula of chloride $GeCl_4$.
• Legacy: Element 101 was named Mendelevium ($Md$) by Glenn T. Seaborg in his honor.

Modern Periodic Law

• Discovery: Henry Moseley (1913) observed that the plot of $\sqrt{\nu}$ (where $\nu$ is the frequency of emitted X-rays) against atomic number ($Z$) gave a straight line, while a plot against atomic mass did not.
• Conclusion: Atomic number is a more fundamental property than atomic mass.
• Modern Periodic Law Statement: "The physical and chemical properties of the elements are periodic functions of their atomic numbers."
• Nature of Atomic Number: It equals nuclear charge (number of protons) or the number of electrons in a neutral atom.
• Significance: This law revealed analogies among 94 naturally occurring elements (including Neptunium and Plutonium found in pitchblende) and guided the creation of synthetic elements.

Presentation of the Modern Periodic Table

• Long Form: The most used version, emphasizing electronic configurations.
• Periods: Seven horizontal rows. The period number corresponds to the highest principal quantum number ($n$).
  • Period 1: 2 elements.
  • Period 2 & 3: 8 elements each.
  • Period 4 & 5: 18 elements each.
  • Period 6: 32 elements (includes Lanthanoids).
  • Period 7: Incomplete; theoretical maximum of 32 elements (includes Actinoids).
• Groups: 18 vertical columns. Elements in a group have similar outer electronic configurations.

IUPAC Nomenclature for Elements with $Z > 100$

• Purpose: To prevent disputes over discovery priority (e.g., Element 104 was contested by American and Soviet scientists).
• Rules: Roots for digits are combined and ended with "ium".
  • Digits: 0=nil, 1=un, 2=bi, 3=tri, 4=quad, 5=pent, 6=hex, 7=sept, 8=oct, 9=enn.
  • Example (Z = 120): Roots are un (1), bi (2), and nil (0). Name: Unbinilium. Symbol: Ubn.
• Table of Official Names: Element 101 ($Md$), 102 ($No$), 103 ($Lr$), 104 ($Rf$), 105 ($Db$), 106 ($Sg$), 107 ($Bh$), 108 ($Hs$), 109 ($Mt$), 110 ($Ds$), 111 ($Rg$), 112 ($Cn$), 114 ($Fl$), 116 ($Lv$).

Electronic Configurations and Blocks (s, p, d, f)

• s-Block: Groups 1 (alkali metals, $ns^1$) and 2 (alkaline earth metals, $ns^2$). They are reactive metals with low ionization enthalpies and form predominantly ionic compounds (except Lithium and Beryllium).
• p-Block: Groups 13 to 18 (Representative Elements). Configuration: $ns^2np^{1-6}$. Ends with Noble Gases ($ns^2np^6$). Includes Non-metals (top right), Halogens (Group 17), and Chalcogens (Group 16).
• d-Block (Transition Elements): Groups 3 to 12. General configuration: $(n-1)d^{1-10}ns^{0-2}$. Characterized by variable oxidation states, colored ions, paramagnetism, and catalytic activity. Zinc, Cadmium, and Mercury ($(n-1)d^{10}ns^2$) do not show typical transition properties.
• f-Block (Inner-Transition Elements): Two rows at the bottom.
  • Lanthanoids: $Ce (Z=58)$ to $Lu (Z=71)$.
  • Actinoids: $Th (Z=90)$ to $Lr (Z=103)$. All are radioactive. Elements after Uranium are called Transuranium Elements.
  • General configuration: $(n-2)f^{1-14}(n-1)d^{0-1}ns^2$.
• Exceptions:
  1. Helium: Chemically a p-block noble gas despite being $1s^2$ (s-block).
  2. Hydrogen: Unique position; placed separately due to properties similar to both alkali metals and halogens.

Classification of Metals, Non-metals, and Metalloids

• Metals: Over $78\%$ of elements. High melting points, great conductors, malleable, and ductile. Located on the left.
• Non-metals: Top right of the table. Usually gases or solids with low melting points (Boron and Carbon are exceptions). Poor conductors, brittle.
• Metalloids (Semi-metals): Silicon, Germanium, Arsenic, Antimony, Tellurium. Border the zig-zag line. Show properties of both metals and non-metals.

Periodic Trends in Physical Properties

• Atomic Radius:
  • Period Trend: Decreases from left to right as effective nuclear charge increases, pulling electrons closer.
  • Group Trend: Increases down a group as the principal quantum number ($n$) increases and inner electrons shield the nucleus.
  • Types: Covalent radius (half distance between atoms in a covalent bond) and Metallic radius (half distance between metal cores in a crystal).
• Ionic Radius:
  • Cations: Smaller than parent atoms (fewer electrons, same charge).
  • Anions: Larger than parent atoms (increased electron repulsion, decreased effective nuclear charge).
  • Isoelectronic Species: Species with the same number of electrons (e.g., $O^{2-}$, $F^-$, $Na^+$, $Mg^{2+}$). Size decreases as positive nuclear charge increases.
• Ionization Enthalpy ($\Delta_i H$): Energy required to remove an electron from a gaseous atom in ground state ($X(g) \rightarrow X^+(g) + e^-$).
  • Trends: Increases across a period; decreases down a group.
  • Anomalies:
    1. $B < Be$: Be has a stable $2s^2$ shell; boron's $2p$ electron is better shielded.
    2. $O < N$: Nitrogen has three $2p$ electrons in separate orbitals (stable); Oxygen has electron-electron repulsion in its doubly occupied $2p$ orbital.
• Electron Gain Enthalpy ($\Delta_{eg} H$): Enthalpy change when adding an electron ($X(g) + e^- \rightarrow X^-(g)$).
  • Trends: Becomes more negative across a period; becomes less negative down a group.
  • Exception: $F$ and $O$ are less negative than $Cl$ and $S$ due to high electron repulsion in the small $n=2$ shell.
• Electronegativity: Ability to attract shared electrons.
  • Pauling Scale: Fluorine is assigned the highest value of $4.0$.
  • Trends: Increases across a period (decreasing radius); decreases down a group.

Periodic Trends in Chemical Properties

• Valence / Oxidation States: Determined by outer electronic configuration. For representative elements, valence is electrons in outermost shell or $8$ minus that number.
• Anomalous Properties of Second Period Elements: Lithium, Beryllium, Boron to Fluorine differ from their group members due to small size, high charge/radius ratio, high electronegativity, and lack of d-orbitals.
• Diagonal Relationship: Similarity between elements in adjacent groups/periods (e.g., $Li-Mg, Be-Al$).
• Chemical Reactivity:
  • High at extremes of a period (Group 1 and 17).
  • Basic Oxides: Formed by elements on the extreme left (e.g., $Na_2O$).
  • Acidic Oxides: Formed by elements on the extreme right (e.g., $Cl_2O_7$).
  • Amphoteric/Neutral Oxides: Center elements (e.g., $Al_2O_3$ is amphoteric; $CO, NO, N_2O$ are neutral).