Classification of Elements and Periodicity in Properties

Significance and Necessity of Element Classification

• The Periodic Table is the most important concept in chemistry, serving as a support for students, a guide for researchers, and a succinct organization of the entire field.
• The chemical elements are not a random collection but display clear trends and organized families. Understanding this structure is essential to seeing how the world is built from fundamental building blocks.
• Historical Growth of Known Elements:
  • In $1800$, only $31$ elements were identified.
  • By $1865$, the number more than doubled to $63$.
  • At present, over $114$ elements are known (with official names recently confirmed up to $118$).
• The classification system was developed to rationalize known chemical facts and predict new elements for future study, making the study of innumerable elements and their compounds manageable.

Genesis of Periodic Classification

• Johann Dobereiner (Early 1800s):
  • The first to suggest trends among element properties.
  • Law of Triads (1829): He identified groups of three elements (Triads) where the middle element had an atomic weight approximately halfway between the other two.
  • Examples of Triads:
    • Lithium ($Li$, weight $7$), Sodium ($Na$, weight $23$), Potassium ($K$, weight $39$).
    • Calcium ($Ca$, weight $40$), Strontium ($Sr$, weight $88$), Barium ($Ba$, weight $137$).
    • Chlorine ($Cl$, weight $35.5$), Bromine ($Br$, weight $80$), Iodine ($I$, weight $127$).
  • The properties of the middle element were also intermediate between the other two members. This was largely dismissed as coincidence as it worked for only a few elements.
• A.E.B. de Chancourtois (1862):
  • A French geologist who arranged elements in increasing order of atomic weights in a cylindrical table to display the periodic recurrence of properties.
• John Alexander Newlands (1865):
  • Law of Octaves: He arranged elements by increasing atomic weight and noted that every eighth element shared properties with the first, similar to the eighth note in a musical octave.
  • This law was only accurate for elements up to Calcium ($Ca$).
  • Despite initial lack of acceptance, Newlands was awarded the Davy Medal in $1887$ by the Royal Society, London.

Mendeleev’s Periodic Law and Predictions

• Development: Credited to Dmitri Mendeleev ($1834-1907$) and Lothar Meyer ($1830-1895$), who worked independently in $1869$.
• Lothar Meyer: Plotted physical properties (atomic volume, melting point, boiling point) against atomic weight, observing a periodically repeated pattern with changing lengths of repeating units.
• Dmitri Mendeleev: First to publish the Periodic Law, stating: "The properties of the elements are a periodic function of their atomic weights."
• Key Features of Mendeleev's Table:
  • Elements were arranged in horizontal rows and vertical columns (groups) based on increasing atomic weight.
  • He prioritized similarity in properties over strict weight order if necessary. For example, Iodine (lower weight) was placed in Group VII with Fluorine, Chlorine, and Bromine rather than after Tellurium.
  • He utilized empirical formulas of oxides and hydrides to classify elements.
  • Bold Predictions: Mendeleev left gaps for undiscovered elements. He named them using the prefix "Eka" (one) followed by the name of the preceding element in the same group.
  • Eka-Aluminium (Gallium): Predicted weight $68$, density $5.9\,g/cm^3$. Found weight $70$, density $5.94\,g/cm^3$.
  • Eka-Silicon (Germanium): Predicted weight $72$, density $5.5\,g/cm^3$. Found weight $72.6$, density $5.36\,g/cm^3$.

Modern Periodic Law and the Present Form of the Periodic Table

• Henry Moseley (1913): Observed regularities in the X-ray spectra of elements. A 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: "The physical and chemical properties of the elements are periodic functions of their atomic numbers."
• Structure of the Long Form:
  • Periods: Seven horizontal rows. The period number corresponds to the highest principal quantum number ($n$).
  • Groups: $18$ vertical columns. Elements in a group share similar outer electronic configurations.
  • Group Numbering: IUPAC recommends numbering from $1$ to $18$, replacing the old notation ($IA-VIIA$, $VIII$, $IB-VIIB$, and $0$).
  • Capacity:
    • Period $1$: $2$ elements ($1s$ subshell).
    • Periods $2$ and $3$: $8$ elements each ($2s, 2p$ and $3s, 3p$ subshells).
    • Periods $4$ and $5$: $18$ elements each (includes $d$ transition series).
    • Period $6$: $32$ elements (includes $4f$ Lanthanoids).
    • Period $7$: Incomplete, theoretically contains $32$ elements (includes $5f$ Actinoids).

Nomenclature of Elements with Atomic Numbers $Z > 100$

• Naming of new elements was historically the privilege of discoverers, but competition between laboratories (e.g., American vs. Soviet claims for Element $104$) led to controversy.
• IUPAC Systematic Nomenclature: Based on numerical roots for the digits of the atomic number:
  • $0 = nil, 1 = un, 2 = bi, 3 = tri, 4 = quad, 5 = pent, 6 = hex, 7 = sept, 8 = oct, 9 = enn$.
  • The name is formed by concatenating the roots and adding "ium".
• Example for $Z = 120$: Root for $1 (un)$ + Root for $2 (bi)$ + Root for $0 (nil)$ + "ium" = Unbinilium (Symbol: $Ubn$).
• Officially Recognized Elements (101-118): Examples include Mendelevium ($101, Md$), Seaborgium ($106, Sg$), and Oganesson ($118, Og$).

Electronic Configurations and the Four Blocks

• Division into Blocks: Elements are classified into $s, p, d, f$ blocks based on which orbital is being filled by the last electron.
• s-Block Elements (Groups 1 and 2):
  • Valence configuration: $ns^1$ (Alkali metals) or $ns^2$ (Alkaline earth metals).
  • Characteristics: Reactive metals, low ionization enthalpies, form $1+$ or $2+$ ions, predominantly ionic compounds.
• p-Block Elements (Groups 13-18):
  • Valence configuration: $ns^2 np^{1-6}$.
  • Includes Representative Elements (Groups $13-17$) and Noble Gases (Group $18$).
  • Preceding Noble gases are Chalcogens (Group $16$) and Halogens (Group $17$), which have high negative electron gain enthalpies.
• d-Block Elements (Groups 3-12):
  • General configuration: $(n-1)d^{1-10} ns^{0-2}$.
  • Known as Transition Elements because they bridge chemically active $s$-block metals and less active $p$-block elements.
  • Characteristics: All are metals, mostly form colored ions, variable oxidation states, paramagnetism, catalytic properties.
  • Exceptions: $Zn, Cd, Hg$ have $(n-1)d^{10} ns^2$ and do not show typical transition properties.
• f-Block Elements (Inner-Transition Elements):
  • General configuration: $(n-2)f^{1-14} (n-1)d^{0-1} ns^2$.
  • Lanthanoids: $Ce (Z=58)$ to $Lu (Z=71)$.
  • Actinoids: $Th (Z=90)$ to $Lr (Z=103)$. All are radioactive; those after Uranium are "Transuranium Elements".

Metals, Non-metals, and Metalloids

• Metals: Over $78\%$ of elements. Positioned on the left and center. Solids at room temperature (except Mercury; Gallium/Caesium have low melting points $303\,K$ and $302\,K$). High conductivity, malleable, and ductile.
• Non-metals: Top right of the table. Usually solids or gases with low melting/boiling points (except Boron, Carbon). Poor conductors; brittle in solid form.
• Metalloids (Semi-metals): Border the zig-zag line separating metals and non-metals. Show properties of both. Includes Silicon ($Si$), Germanium ($Ge$), Arsenic ($As$), Antimony ($Sb$), and Tellurium ($Te$).

Periodic Trends in Physical Properties

• Atomic Radius:
  • Defined as Covalent Radius (non-metals) or Metallic Radius (metals).
  • Trend across a period: Decreases because effective nuclear charge increases, pulling electrons closer.
  • Trend down a group: Increases because principal quantum number ($n$) increases and inner shells shield outer electrons.
  • Noble gases have large radii because they are measured as van der Waals radii (non-bonded).
• Ionic Radius:
  • Cations are smaller than parent atoms (fewer electrons, same charge).
  • Anions are larger than parent atoms (increased electron repulsion, decreased effective nuclear charge).
  • Isoelectronic Species: Atoms/ions with the same number of electrons (e.g., $O^{2-}, F^-, Na^+, Mg^{2+}$). The radius decreases as nuclear charge increases ($Mg^{2+}$ is smallest, $O^{2-}$ is largest).
• Ionization Enthalpy ($\Delta_i H$):
  • Enthalpy required to remove an electron from an isolated gaseous atom in the ground state: $X(g) \rightarrow X^+(g) + e^-$.
  • Unit: $kJ\,mol^{-1}$. Always positive ($IE_1 < IE_2 < IE_3$).
  • Trend: Increases across a period, decreases down a group.
  • Anomalies: Boron has lower $IE$ than Beryllium because removing a $2p$ electron is easier than a penetrating $2s$ electron. Oxygen has lower $IE$ than Nitrogen due to electron-electron repulsion in the doubly occupied $2p$ orbital.
• Electron Gain Enthalpy ($\Delta_{eg}H$):
  • Enthalpy change when an electron is added to a gaseous atom: $X(g) + e^- \rightarrow X^-(g)$.
  • Trend: Becomes more negative across a period. Becomes less negative down a group.
  • Negative values are high for halogens ($ns^2 np^5$). Noble gases have large positive values.
  • Anomaly: $F$ and $O$ are less negative than $Cl$ and $S$ because their small $n=2$ shell results in significant electron-electron repulsion.
• Electronegativity:
  • Ability of an atom in a compound to attract shared electrons. Not a measurable quantity; Pauling scale is used (Fluorine = $4.0$).
  • Trend: Increases across a period, decreases down a group. Inversely related to metallic character.

Periodic Trends in Chemical Properties

• Valence/Oxidation States:
  • Valence is usually equal to the number of electrons in the outermost shell or $8$ minus that number.
  • Oxidation state is the charge acquired based on electronegativity (e.g., in $OF_2$, $O$ is $+2$; in $Na_2O$, $O$ is $-2$).
• Anomalous Properties of Second Period Elements:
  • $Li, Be, B, C, N, O, F$ differ from their groups due to small size, high charge/radius ratio, and lack of $d$ orbitals.
  • Maximum covalency for the second period is $4$. Subsequent periods can expand covalency using $d$ orbitals.
  • Ability to form $p\pi - p\pi$ multiple bonds is unique to the second period.
• Chemical Reactivity:
  • High at the extremes (Group $1$ and Group $17$), lowest in the middle.
  • Group $1$ oxides are most basic ($Na_2O$). Group $17$ oxides are most acidic ($Cl_2O_7$).
  • Middle oxides are amphoteric ($Al_2O_3, As_2O_3$) or neutral ($CO, NO, N_2O$).

Questions & Discussion

• Question: What is the IUPAC name and symbol for element $120$?
  • Response: The roots for $1, 2, 0$ are un, bi, and nil. Thus, the name is unbinilium and the symbol is $Ubn$.
• Question: Why does the 5th period have $18$ elements?
  • Response: At $n=5$, the energy order is $5s < 4d < 5p$. These comprise $1, 5$, and $3$ orbitals respectively, totaling $9$ orbitals. Each orbital holds two electrons, allowing for $18$ elements.
• Question: Which is larger: $Mg$ or $Al^{3+}$?
  • Response: $Mg$ is larger. Atomic radii decrease across a period ($Mg > Al$), and cations are smaller than parent atoms. $Al^{3+}$ is much smaller due to higher nuclear charge relative to electrons.
• Question: Predict the formula of a compound between Aluminum and Sulfur.
  • Response: Aluminum (Group $13$) has valence $3$. Sulfur (Group $16$) has valence $2$. The resulting formula is $Al_2S_3$.
• Question: Demonstrate that $Na_2O$ is basic and $Cl_2O_7$ is acidic using water.
  • Response: $Na_2O + H_2O \rightarrow 2NaOH$ (Strong base). $Cl_2O_7 + H_2O \rightarrow 2HClO_4$ (Strong acid).