Classification of Elements and Periodicity in Properties

Introduction and Significance of the Periodic Table

• According to Glenn T. Seaborg, the Periodic Table is the most important concept in chemistry, serving both as a foundational principle and a practical tool for students and professionals.

• It organizes all known chemical elements and demonstrates that they are not a random cluster but follow specific trends and group into families.

• Understanding the Periodic Table is vital for understanding how the world is structured from fundamental building blocks known as chemical elements.

• This study unit covers several key learning objectives:

  • Appreciation of how grouping elements by properties led to the Periodic Table's development.

  • Understanding the Periodic Law and the significance of atomic number ($Z$) and electronic configuration.

  • Learning the International Union of Pure and Applied Chemistry (IUPAC) nomenclature for elements with Z > 100.

  • Classification of elements into $s$, $p$, $d$, and $f$ blocks and recognizing their characteristics.

  • Recognition of periodic trends in physical and chemical properties (atomic/ionic radii, ionization enthalpy, electron gain enthalpy, electronegativity, valence).

  • Correlation of reactivity with natural occurrence and ionization enthalpy with metallic character.

The Necessity of Element Classification

• Elements are the basic units of all matter.

• Historical growth of known elements:

  • In 1800: $31$ elements known.

  • By 1865: $63$ elements known (more than double).

  • At present: $114$ elements known (some are man-made through synthesis efforts).

• Classification is required because the sheer number of elements and their innumerable compounds makes individual study difficult. Its purpose is to rationalize chemical facts and predict new ones.

Genesis of Periodic Classification

• Johann Dobereiner (Early 1800s): The first to consider property trends. By 1829, he identified groups of three elements called ‘Triads’ where the middle element had an atomic weight approximately halfway between the other two. Properties of the middle element were also intermediate.

  • Table 3.1: Dobereiner’s Triads

    • Triad 1: $Li$ ($7$), $Na$ ($23$), $K$ ($39$)

    • Triad 2: $Ca$ ($40$), $Sr$ ($88$), $Ba$ ($137$)

    • Triad 3: $Cl$ ($35.5$), $Br$ ($80$), $I$ ($127$)

  • This was dismissed as coincidence because the Law of Triads only worked for a few elements.

• A.E.B. de Chancourtois (1862): A French geologist who arranged elements by increasing atomic weight into a cylindrical table to show periodic recurrence of properties. This attempt did not receive much attention.

• John Alexander Newlands (1865): The English chemist proposed the ‘Law of Octaves’. He arranged elements by increasing atomic weight and found that every eighth element resembled the first, like music octaves.

  • Table 3.2: Newlands’ Octaves

    • Row 1: $Li$ ($7$), $Be$ ($9$), $B$ ($11$), $C$ ($12$), $N$ ($14$), $O$ ($16$), $F$ ($19$)

    • Row 2: $Na$ ($23$), $Mg$ ($24$), $Al$ ($27$), $Si$ ($29$), $P$ ($31$), $S$ ($32$), $Cl$ ($35.5$)

    • Row 3: $K$ ($39$), $Ca$ ($40$)

  • Limitations: It only seemed accurate for elements up to Calcium. Newlands was later awarded the Davy Medal in 1887 by the Royal Society, London.

• Dmitri Mendeleev and Lothar Meyer (1869): Independently proposed that properties are periodic functions of atomic weights.

  • Lothar Meyer plotted physical properties (atomic volume, melting point, boiling point) against atomic weight, observing a repeating pattern with varying lengths. He developed a table by 1868 similar to the modern one, but Mendeleev published first.

  • Mendeleev is credited with the Modern Periodic Table's roots. He prioritized periodicity over strict atomic weight order, sometimes ignoring atomic measurements he believed were incorrect to group elements with similar properties together (e.g., placing Iodine after Tellurium).

  • Mendeleev left gaps for undiscovered elements, naming them ‘Eka-Aluminium’ and ‘Eka-Silicon.’

  • Table 3.3: Mendeleev’s Predictions (Eka-Aluminium vs. Gallium and Eka-Silicon vs. Germanium):

    • Eka-Aluminium (predicted): At. wt. $68$, Density $5.9\,\text{g/cm}^3$, low MP, Oxide $E_2O_3$, Chloride $ECl_3$.

    • Gallium (found): At. wt. $70$, Density $5.94\,\text{g/cm}^3$, MP $302.93\,K$, Oxide $Ga_2O_3$, Chloride $GaCl_3$.

    • Eka-Silicon (predicted): At. wt. $72$, Density $5.5\,\text{g/cm}^3$, high MP, Oxide $EO_2$, Chloride $ECl_4$.

    • Germanium (found): At. wt. $72.6$, Density $5.36\,\text{g/cm}^3$, MP $1231\,K$, Oxide $GeO_2$, Chloride $GeCl_4$.

Modern Periodic Law

• Mendeleev worked without knowledge of sub-atomic structures.

• Henry Moseley (1913): Observed regularities in characteristic X-ray spectra. A plot of $\sqrt{\nu}$ (frequency) against atomic number ($Z$) yielded a straight line, which it did not for atomic mass. This proved that atomic number is the more fundamental property.

• Modern Periodic Law: The physical and chemical properties of the elements are periodic functions of their atomic numbers.

• Significance of $Z$: Atomic number equals the nuclear charge (number of protons) and the number of electrons in a neutral atom.

• Periodicity is the consequence of the periodic variation in electronic configurations.

Present Form of the Periodic Table

• The ‘Long Form’ of the Periodic Table is most widely used.

• Horizontal Rows: Called ‘Periods.’ There are seven periods. The period number matches the highest principal quantum number ($n$) of the elements in that period.

  • Period 1: $2$ elements ($n=1$).

  • Subsequent: $8, 8, 18, 18$, and $32$ elements.

  • Period 7: Incomplete, theoretical max of $32$.

• Vertical Columns: Called ‘Groups’ or ‘Families.’ Elements in a group have similar outer electronic configurations.

  • IUPAC numbering: 1 to 18 (replacing old IA…VIIIA tags).

• Separate Panels: Lanthanoids (6th period) and Actinoids (7th period) are placed at the bottom to maintain the table's structure.

• Glenn T. Seaborg: His work starting with Plutonium (1940) and continuing through transuranium elements 94 to 102 led to the reconfiguration of the table. Nobel Prize 1951. Element 106, Seaborgium ($Sg$), is named in his honor.

Nomenclature for Elements with Z > 100

• High atomic number elements are unstable and produced in minimal quantities, often leading to discovery disputes (e.g., element 104: Americans named it Rutherfordium, Soviets named it Kurchatovium).

• IUPAC systematic nomenclature is used until an official name is ratified. It uses numerical roots for digits 0-9:

  • 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)

• The roots are joined and ‘ium’ is added at the end.

• Table 3.5: Examples include 101 (Unnilunium - Mendelevium), 104 (Unnilquadium - Rutherfordium), 112 (Ununbium - Copernicium), 118 (Ununoctium - Oganesson).

Electronic Configurations and the Periodic Table

• Periods: Successive periods fill higher energy levels ($n=1, 2,\dots$). The number of elements in a period is double the available orbitals in that energy level.

  • $n=1$: $1s$ orbital ($2$ elements: $H$, $He$).

  • $n=2$: $2s, 2p$ orbitals ($8$ elements: $Li$ to $Ne$).

  • $n=3$: $3s, 3p$ orbitals ($8$ elements: $Na$ to $Ar$).

  • $n=4$: $4s, 3d, 4p$ orbitals ($18$ elements: $K$ to $Kr$). Includes the ‘$3d$ transition series’ ($Sc$ to $Zn$).

  • $n=5$: $5s, 4d, 5p$ orbitals ($18$ elements: $Rb$ to $Xe$). Includes the ‘$4d$ transition series’ ($Y$ to $Cd$).

  • $n=6$: $6s, 4f, 5d, 6p$ orbitals ($32$ elements). Includes the ‘lanthanoid series’ ($Ce$ to $Lu$).

  • $n=7$: $7s, 5f, 6d, 7p$ orbitals. Includes the ‘actinoid series’ starting after Actinium ($Ac$, $Z=89$).

• Groups: Similar valence shell configurations lead to similar chemical behavior.

  • Group 1 (Alkali Metals): All have $ns^1$ configuration (e.g., $Li$ [$He$]$2s^1$, $Na$ [$Ne$]$3s^1$).

Block-wise Classification (s, p, d, f)

• s-Block: Groups 1 (alkali metals) and 2 (alkaline earth metals). Outer configuration $ns^1$ and $ns^2$. Reactive metals, low ionization enthalpies, form $1+$ and $2+$ ions. Predominantly ionic (except $Li$, $Be$).

• p-Block: Groups 13 to 18. Outer configuration $ns^2np^1$ to $ns^2np^6$. Includes metals, non-metals, and metalloids. Together with s-block, they are ‘Representative Elements.’ Group 17 are Halogens, group 16 are Chalcogens. Group 18 are Noble Gases with closed shells ($ns^2np^6$).

• d-Block: Groups 3 to 12. ‘Transition Elements.’ Characterized by filling of inner $d$ orbitals. General config: $(n-1)d^{1-10}ns^{0-2}$ (Exception: $Pd$ is $4d^{10}5s^0$). All are metals, form colored ions, variable valence, paramagnetism, and catalysts. $Zn, Cd, Hg$ ($(n-1)d^{10}ns^2$) lack typical transition properties.

• f-Block: ‘Inner-Transition Elements.’ Lanthanoids ($Ce$ to $Lu$) and Actinoids ($Th$ to $Lr$). Outer config: $(n-2)f^{1-14}(n-1)d^{0-1}ns^2$. All are metals. Actinoids are radioactive, many synthesized in tiny quantities. Elements after Uranium are ‘Transuranium Elements.’

• Exceptions: Helium strictly belongs to s-block ($1s^2$) but is placed in p-block (Group 18) due to its filled shell properties. Hydrogen ($1s^1$) can be in Group 1 (metal-like) or Group 17 (halogen-like) and is usually placed separately.

Metals, Non-metals, and Metalloids

• Metals: Over $78\%$ of elements. High MP/BP, malleable, ductile, good conductors. Located left/center.

• Non-metals: Located top-right. Low MP/BP (except $B$, $C$), poor conductors, brittle.

• Metalloids (Semi-metals): Border the zig-zag line. properties of both. Examples: $Si, Ge, As, Sb, Te$.

• Trend: Metallic character decreases left-to-right across a period and increases down a group.

Periodic Trends in Physical Properties

• Atomic Radius: Size of the atom ($\sim 1.2\,\text{Å}$ or $1.2 \times 10^{-10}\,m$). Measured via bond distance.

  • Covalent Radius: Half the bond distance in a covalent molecule (e.g., $Cl_2$ bond is $198\,pm$, radius is $99\,pm$).

  • Metallic Radius: Half the internuclear distance in a metallic crystal (e.g., $Cu$ distance is $256\,pm$, radius is $128\,pm$).

  • Trends: Decreases across a period (increasing nuclear charge attracting same-shell electrons). Increases down a group (increased principal quantum number $n$ and shielding by inner electrons).

  • Noble Gas radii are large because they represent non-bonded van der Waals radii.

• Ionic Radius: Estimated from distances in ionic crystals. Cations are always smaller than parent atoms (same charge, fewer electrons). Anions are larger (same charge, more electrons, increased repulsion).

  • Isoelectronic species: Species with the same number of electrons (e.g., $O^{2-}, F^-, Na^+, Mg^{2+}$). Size depends on nuclear charge; higher positive charge means smaller radius.

• Ionization Enthalpy ($\Delta_i H$): Energy required to remove an electron from an isolated gaseous atom in its ground state.

  • Reaction: $X(g) \rightarrow X^+(g) + e^-$ ($\Delta_i H$ is always positive).

  • Successive ionization enthalpies: \Delta_i H_1 < \Delta_i H_2 < \Delta_i H_3 because it's harder to remove electrons from positive ions.

  • Trends: Increases across a period (higher nuclear charge), decreases down a group (outer electron farther and more shielded).

  • Anomalies: Boron ($Z=5$) is lower than Beryllium ($Z=4$) because removing a $2p$ electron is easier than a more-penetrating $2s$ electron. Oxygen is lower than Nitrogen because of electron-electron repulsion between paired electrons in Nitrogen’s $2p$ orbital.

• 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 (attaining noble gas config). Noble gases have large positive values (forced into higher $n$ levels).

  • Trends: Becomes more negative left-to-right. Becomes less negative down a group.

  • Exception: $O$ and $F$ are less negative than $S$ and $Cl$ because the small size of $n=2$ results in high repulsion for the added electron.

• Electronegativity: Ability to attract shared electrons. Pauling Scale (Linus Pauling, 1922) gives Fluorine the highest value ($4.0$).

  • Trends: Increases across a period, decreases down a group. Follows same trend as ionization enthalpy.

Periodic Trends in Chemical Properties

• Valence/Oxidation State: Characteristic of electronic configuration. For representative elements, valence corresponds to the number of valence electrons or $8$ minus that number.

• Oxidation State: Charge acquired by an atom based on electronegativity (e.g., in $OF_2$, $F$ is $-1$, $O$ is $+2$; in $Na_2O$, $O$ is $-2$, $Na$ is $+1$).

• Anomalous Properties of 2nd Period: $Li, Be$ and $B-F$ differ from their groups due to small size, high charge/radius ratio, high electronegativity, and lack of $d$ orbitals (limiting max covalency to $4$).

• Diagonal Relationship: Similarities between $Li$ and $Mg$, or $Be$ and $Al$.

• Chemical Reactivity: Highest at extremes (loss/gain of electrons), lowest in the center.

• Oxides:

  • Extreme left: Basic (e.g., $Na_2O$).

  • Extreme right: Acidic (e.g., $Cl_2O_7$).

  • Center: Amphoteric (e.g., $Al_2O_3, As_2O_3$) or Neutral (e.g., $CO, NO, N_2O$).

Questions & Discussion

• Problem 3.1: What is the naming convention for element 120?

  • Response: Root for 1=un, 2=bi, 0=nil. Name: Unbinilium. Symbol: $Ubn$.

• Problem 3.2: Why are there 18 elements in the 5th period?

  • Response: For $n=5$, available orbitals are $5s$, $4d$, and $5p$. Total orbitals = $1 + 5 + 3 = 9$. Max electrons = $2 \times 9 = 18$.

• Problem 3.3: Where would $Z=117$ and $120$ be placed?

  • Response: $Z=117$ in Group 17 (Halogens); config: [$Rn$]$5f^{14}6d^{10}7s^27p^5$. $Z=120$ in Group 2 (Alkaline earth metals); config: [$Uuo$]$8s^2$.

• Problem 3.4: Order of metallic character for $Si, Be, Mg, Na, P$?

  • Response: P < Si < Be < Mg < Na.

• Problem 3.5: Largest and smallest of $Mg, Mg^{2+}, Al, Al^{3+}$?

  • Response: Largest is $Mg$ (parent atom). Smallest is $Al^{3+}$ (cation with highest positive charge among isoelectronics).

• Problem 3.6: Is $Al$'s first $\Delta_i H$ closer to $575$ or $760\,kJ\,mol^{-1}$?

  • Response: Closer to $575\,kJ\,mol^{-1}$. It should be lower than $Mg$ ($737$) due to shielding of $3p$ by $3s$.

• Problem 3.7: Most and least negative $\Delta_{eg} H$ for $P, S, Cl, F$?

  • Response: Most negative: $Cl$. Least negative: $P$.

• Problem 3.8: Predict formulas for (a) Silicon and Bromine (b) Aluminium and Sulphur.

  • Response: (a) $SiBr_4$ (b) $Al_2S_3$.

• Problem 3.9: Oxidation state and covalency of $Al$ in $[AlCl(H_2O)_5]^{2+}$?

  • Response: Oxidation state is $+3$; covalency is $6$.

• Problem 3.10: Chemical reactions showing $Na_2O$ as basic and $Cl_2O_7$ as acidic.

  • Response:

    • $Na_2O + H_2O \rightarrow 2NaOH$

    • $Cl_2O_7 + H_2O \rightarrow 2HClO_4$

• Exercise 3.15: Calculate $\Delta_i H$ of atomic hydrogen if ground state energy is $-2.18 \times 10^{-18}\,J$.

  • Response: Ionization requires supplying that energy. Per mole: $2.18 \times 10^{-18}\,J \times 6.022 \times 10^{23}\,mol^{-1} = 1.312 \times 10^6\,J\,mol^{-1}$.