Classification of Elements and Periodicity in Properties Notes

Significance of the Periodic Table in Chemistry

• The Periodic Table is the most important concept in chemistry, both in principle and in practice.

• It serves as everyday support for students, suggests research avenues for professionals, and provides a succinct organization of the entirety of chemistry.

• It demonstrates that chemical elements are not random entities but display trends and exist in families.

• The table is essential for understanding how the world is constructed from fundamental building blocks called chemical elements.

Genesis of Periodic Classification

• Classification is the result of systematizing knowledge gained through observations and experiments by numerous scientists.

• Johann Dobereiner (Early 1800s): The German chemist was the first to consider trends among element properties. By 1829, he identified "Triads," which were groups of three elements with similar physical and chemical properties.

• Law of Triads: In these groups, the middle element had an atomic weight approximately halfway between the other two. The properties of the middle element were also intermediate.

• Examples of Dobereiner’s Triads:

  • $Li$ ($7$), $Na$ ($23$), $K$ ($39$)
  • $Ca$ ($40$), $Sr$ ($88$), $Ba$ ($137$)
  • $Cl$ ($35.5$), $Br$ ($80$), $I$ ($127$)

• A.E.B. de Chancourtois (1862): A French geologist who arranged elements by increasing atomic weight and created a cylindrical table to display the periodic recurrence of properties.

• John Alexander Newlands (1865): An English chemist who proposed the Law of Octaves. He arranged elements in increasing order of atomic weights and noted that every eighth element had properties similar to the first, mimicking musical octaves.

• Newlands' Octaves Example:

  • $Li$ ($7$), $Be$ ($9$), $B$ ($11$), $C$ ($12$), $N$ ($14$), $O$ ($16$), $F$ ($19$)
  • $Na$ ($23$), $Mg$ ($24$), $Al$ ($27$), $Si$ ($29$), $P$ ($31$), $S$ ($32$), $Cl$ ($35.5$)
  • $K$ ($39$), $Ca$ ($40$)

• Newlands’ Law of Octaves was effective only up to calcium. He was later awarded the Davy Medal in 1887 by the Royal Society, London.

• Dmitri Mendeleev (1834-1907) and Lothar Meyer (1830-1895): These scientists independently proposed in 1869 that arranging elements by atomic weight reveals periodic similarities in properties.

• Lothar Meyer's Approach: Plotted physical properties (atomic volume, melting point, boiling point) against atomic weight, observing a periodically repeated pattern with changing lengths.

• Mendeleev’s Periodic Law: The properties of the elements are a periodic function of their atomic weights.

• Mendeleev arranged elements in horizontal rows and vertical columns so that elements with similar properties occupied the same group. He relied heavily on the empirical formulas and properties of compounds (like oxides and hydrides).

• Mendeleev’s Bold Predictions: He left gaps for undiscovered elements, such as Eka-Aluminium and Eka-Silicon. When discovered later as Gallium ($Ga$) and Germanium ($Ge$), their properties closely matched his predictions.

• Comparison of Predicted vs. Found Properties:

  • Eka-aluminium (Predicted): Atomic weight: $68$; Density: $5.9\,g/cm^3$; Oxide: $E_2O_3$; Chloride: $ECl_3$.
  • Gallium (Found): Atomic weight: $70$; Density: $5.94\,g/cm^3$; Oxide: $Ga_2O_3$; Chloride: $GaCl_3$.
  • Eka-silicon (Predicted): Atomic weight: $72$; Density: $5.5\,g/cm^3$; Oxide: $EO_2$; Chloride: $ECl_4$.
  • Germanium (Found): Atomic weight: $72.6$; Density: $5.36\,g/cm^3$; Oxide: $Ge_2O_2$; Chloride: $GeCl_4$.

Modern Periodic Law and The Long Form Table

• Henry Moseley (1913): Observed regularities in the characteristic X-ray spectra of elements. He found that a plot of $\sqrt{\nu}$ (where $\nu$ is frequency) against atomic number ($Z$) produced a straight line, whereas atomic mass did not.

• Moseley concluded that the 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.

• The atomic number ($Z$) equals the nuclear charge (number of protons) or the number of electrons in a neutral atom.

• Structural Features:

  • Horizontal Rows: Known as Periods. There are 7 periods. The period number corresponds to the highest principal quantum number ($n$).
  • Vertical Columns: Known as Groups or Families. There are 18 groups, numbered 1 to 18 according to IUPAC recommendations (replacing IA…VIIA, VIII, IB…VIIB, and 0).
  • Period Lengths: 1st ($2$ elements), 2nd ($8$), 3rd ($8$), 4th ($18$), 5th ($18$), 6th ($32$). The 7th is incomplete but theorized to hold $32$.
  • Lanthanoids and Actinoids: 14 elements from the 6th and 7th periods, respectively, placed in separate panels at the bottom to maintain the table's structure.

Nomenclature for Elements with $Z > 100$

• High atomic number elements are often unstable and synthesized in minute quantities. Historically, discovery claims led to controversy (e.g., Element 104 was claimed as Rutherfordium by Americans and Kurchatovium by Soviets).

• IUPAC Systematic Nomenclature: Uses numerical roots based on digits:

  • $0$: nil (n)
  • $1$: un (u)
  • $2$: bi (b)
  • $3$: tri (t)
  • $4$: quad (q)
  • $5$: pent (p)
  • $6$: hex (h)
  • $7$: sept (s)
  • $8$: oct (o)
  • $9$: enn (e)

• Naming Rule: Roots are combined in digit order, ending with "-ium."

• Example for $Z = 120$: un + bi + nil + ium = Unbinilium (Symbol: $Ubn$).

Electronic Configurations and Orbital Blocks

• The location of an element in the table reflects the quantum numbers of the last orbital filled. Elements are classified into four blocks:

• s-Block Elements (Groups 1 and 2):

  • Outermost configurations: $ns^1$ (Alkali metals) and $ns^2$ (Alkaline earth metals).
  • Characteristics: Reactive metals, low ionization enthalpies, lose electrons readily, form predominantly ionic compounds (except $Li$ and $Be$).

• p-Block Elements (Groups 13 to 18):

  • Outermost configuration: $ns^2np^1$ to $ns^2np^6$.
  • Include Noble Gases (Group 18), Halogens (Group 17) and Chalcogens (Group 16).
  • Together with the s-block, they are called Representative Elements.

• d-Block Elements (Groups 3 to 12 - Transition Elements):

  • Characterized by filling inner d-orbitals.
  • General configuration: $(n-1)d^{1-10}ns^{0-2}$.
  • Properties: Metals, colored ions, variable oxidation states, paramagnetism, catalytic properties.

• f-Block Elements (Inner-Transition Elements):

  • Two rows: Lanthanoids ($Z=58-71$) and Actinoids ($Z=90-103$).
  • Configuration: $(n-2)f^{1-14}(n-1)d^{0-1}ns^2$.
  • All are metals. Actinoids are radioactive; elements after Uranium are Transuranium Elements.

Periodic Trends in Physical Properties

• Atomic Radius:

  • Size is approximately $1.2\,$ ($1.2 \times 10^{-10}\,m$).
  • Covalent Radius: Half the distance between two atoms bound by a single bond (e.g., $Cl_2$ bond distance is $198\,pm$, radius is $99\,pm$).
  • Metallic Radius: Half the internuclear distance between metal cores in a crystal (e.g., copper atoms are separated by $256\,pm$, radius is $128\,pm$).
  • Trends: Decreases across a period (due to increasing effective nuclear charge); increases down a group (due to increased principal quantum number and shielding).

• Ionic Radius:

  • Cations: Smaller than their parent atoms (fewer electrons, same nuclear charge).
  • Anions: Larger than their parent atoms (increased electron repulsion, decreased effective nuclear charge).
  • Isoelectronic Species: Atoms/ions with the same number of electrons (e.g., $O^{2-}$, $F^-$, $Na^+$, $Mg^{2+}$ all have 10 electrons). Size decreases with increasing nuclear charge.

• Ionization Enthalpy ($\Delta_i H$):

  • Energy required to remove an electron from an isolated gaseous atom in the ground state.
  • $X(g) \rightarrow X^+(g) + e^-$ ($kJ\,mol^{-1}$).
  • Trends: Increases across a period; decreases down a group.
  • Factors: Shielding/Screening effect (effective nuclear charge is less than actual charge); Penetration of orbitals ($s > p > d > f$).
  • Anomalies: Boron ($Z=5$) has lower $\Delta_i H$ than Beryllium ($Z=4$) because 2p electrons are easier to remove than 2s. Nitrogen has higher $\Delta_i H$ than Oxygen because Nitrogen has a stable half-filled p-shell ($2p^3$).

• Electron Gain Enthalpy ($\Delta_{eg} H$):

  • Enthalpy change when an electron is added to a neutral gaseous atom: $X(g) + e^- \rightarrow X^-(g)$.
  • Halogens have high negative values. Noble gases have large positive values.
  • Trends: Becomes more negative across a period; generally less negative down a group.
  • Anomalies: $F$ has a less negative value than $Cl$ because the addition of an electron to the small $n=2$ level in fluorine results in significant electron-electron repulsion.

• Electronegativity (\chi):

  • Tendency of an atom to attract shared electrons in a chemical compound. Not a measurable quantity.
  • Pauling Scale: Fluorine is assigned the highest value of $4.0$.
  • Trends: Increases across a period; decreases down a group.

Periodic Trends in Chemical Properties and Reactivity

• Valence/Oxidation State: Usually equal to the number of valence electrons or eight minus that number.

• Oxidation State Definition: The charge acquired by an atom based on electronegativity considerations from other atoms in a molecule.

• Example: In $OF_2$, oxygen is $+2$. In $Na_2O$, oxygen 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, high electronegativity, and absence of d-orbitals.

  • Diagonal Relationship: $Li$ resembles $Mg$; $Be$ resembles $Al$.
  • Covalency Limitation: First members have a maximum covalency of 4 (only $2s$ and $2p$ orbitals), while others can expand their valence shell (using $d$-orbitals).

• Chemical Reactivity:

  • Highest at extremes (Groups 1 and 17), lowest in the center.
  • Oxides: Elements on the far left form basic oxides (e.g., $Na_2O$). Elements on the far right form acidic oxides (e.g., $Cl_2O_7$). Elements in the center form amphoteric oxides (e.g., $Al_2O_3$) or neutral oxides (e.g., $CO, NO$).

Problem Solving and Numerical Examples

• Problem 3.1: Symbol and name for $Z = 120$.
  • Roots for 1, 2, 0 are un, bi, nil. Result: unbinilium, symbol Ubn.
• Problem 3.2: Why are there 18 elements in the 5th period?
  • At $n=5$, available orbitals are $5s, 4d, 5p$. Total orbitals = $1 + 5 + 3 = 9$. Maximum electrons = $18$, thus 18 elements.
• Problem 3.3: Locating undiscovered elements $Z=117$ and $120$.
  • $Z=117$: Halogen family (Group 17), configuration $[Rn] 5f^{14}6d^{10}7s^27p^5$.
  • $Z=120$: Alkaline earth metals (Group 2), configuration $[Og] 8s^2$.
• Problem 3.4: Metallic character order for $Si, Be, Mg, Na, P$.
  • Metallic character increases down a group and decreases across a period. Order: $P < Si < Be < Mg < Na$.
• Problem 3.5: Largest and smallest size among $Mg, Mg^{2+}, Al, Al^{3+}$.
  • Largest: $Mg$. Smallest: $Al^{3+}$ (isoelectronic with $Mg^{2+}$ but higher nuclear charge).
• Problem 3.8: Formulas of pairs $Si$/$Br$ and $Al$/$S$.
  • $Si$ (valence 4) + $Br$ (valence 1) $\rightarrow SiBr_4$.
  • $Al$ (valence 3) + $S$ (valence 2) $\rightarrow Al_2S_3$.
• Problem 3.10: Nature of oxides with water.
  • Basic: $Na_2O + H_2O \rightarrow 2NaOH$.
  • Acidic: $Cl_2O_7 + H_2O \rightarrow 2HClO_4$.