Periodic classification notes

Importance and Objectives of Periodic Classification

• Conceptual Significance: The Periodic Table is considered the most important concept in chemistry, serving as a support for students, a guide for research avenues for professionals, and a succinct organization of the entire field. It demonstrates that chemical elements are not random entities but display trends and form families.

• Glenn T. Seaborg Quote: "An awareness of the Periodic Table is essential to anyone who wishes to disentangle the world and see how it is built up from the fundamental building blocks of the chemistry, the chemical elements."

• Learning Objectives:

  • Appreciate the historical development of grouping elements.

  • Understand the Periodic Law (both original and modern).

  • Recognize the significance of atomic number ($Z$) and electronic configuration.

  • Master IUPAC nomenclature for elements where Z > 100.

  • Classify elements into $s, p, d, f$ blocks and understand their characteristics.

  • Identify periodic trends in physical properties (radii, ionization enthalpy, electronegativity) and chemical reactivity (valence, metallic/non-metallic character).

Why the Need for Classification?

• Growth of Known Elements:

  • In 1800, only 31 elements were known.

  • By 1865, the number doubled to 63.

  • At present, 114 elements are known (recently discovered ones are man-made).

• Scientific Rationalization: Individual study of over 100 elements and their innumerable compounds is impractical. Classification allows scientists to organize knowledge, rationalize chemical facts, and predict properties of unknown elements.

Genesis of Periodic Classification

• Johann Dobereiner (Early 1800s):

  • The 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. The properties of the middle element were also intermediate.

  • 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).

  • Limitation: It worked for only a few elements and was dismissed as a coincidence.

• A.E.B. de Chancourtois (1862): A French geologist who arranged elements in order of increasing atomic weights in a cylindrical table. This did not attract significant attention.

• John Alexander Newlands (1865):

  • The Law of Octaves: Arranged elements in increasing order of atomic weights. He observed that every eighth element had properties similar to the first, analogous to musical octaves.

  • Limitations: The law seemed true only for elements up to Calcium ($Ca$). While not initially accepted, he later received the Davy Medal in 1887.

• Lothar Meyer and Dmitri Mendeleev (1869):

  • Both worked independently to propose that properties of elements repeat at regular intervals when arranged by atomic weight.

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

  • Dmitri Mendeleev: Generally credited for the modern periodic system as he published his work first and more elaborately.

Mendeleev’s Periodic Law and Table

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

• Structural Organization: Elements were arranged in horizontal rows (series) and vertical columns (groups) such that elements with similar properties occupied the same column.

• Recognized Periodicity: Mendeleev relied on empirical formulas and properties of compounds (oxides and hydrides). He prioritized property similarity over strict atomic weight order. For example, Iodine (lower weight) was placed after Tellurium (higher weight) to sit with Fluorine, Chlorine, and Bromine.

• Predictive Power: He left gaps for undiscovered elements, naming them "Eka-" (Sanskrit for one) after the element above them.

  • Eka-Aluminum: Later discovered as Gallium ($Ga$).

  • Eka-Silicon: Later discovered as Germanium ($Ge$).

• Comparison of Predicted vs. Experimental Values (Table 3.3):

  • Eka-aluminum (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: $GeO_2$, Chloride: $GeCl_4$.

• Biographical Note on Mendeleev: Born in Tobalsk, Siberia. Appointed Professor of General Chemistry at the University of St. Petersburg in 1867. His textbook "Principles of Chemistry" led to the Periodic Table. Element 101, Mendelevium ($Md$), was named in his honor by Glenn T. Seaborg.

Modern Periodic Law and the Long Form of the Periodic Table

• Henry Moseley (1913): Observed the X-ray spectra of elements. He found that a plot of $\sqrt{\nu}$ (where $\nu$ is frequency) against atomic number ($Z$) gave a straight line, unlike the plot against atomic mass. He proved that atomic number is a more fundamental property.

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

• Significance of Atomic Number: $Z = \text{nuclear charge (protons)} = \text{number of electrons in a neutral atom}$ . Periodicity is a consequence of periodic variation in electronic configurations.

• Structure of the "Long Form" Table:

  • Periods: Horizontal rows (7 in total). The period number corresponds to the highest principal quantum number ($n$).

  • Groups: Vertical columns (numbered 1-18 according to IUPAC, replacing IA-VIIA, VIII, IB-VIIB, 0). Groups contain elements with similar outer electronic configurations.

  • Capacity of Periods:

    1. First Period: $2$ elements ($1s$ shell filled).

    2. Second Period: $8$ elements.

    3. Third Period: $8$ elements.

    4. Fourth Period: $18$ elements (includes $3d$ transition series).

    5. Fifth Period: $18$ elements (includes $4d$ transition series).

    6. Sixth Period: $32$ elements (includes $4f$ Lanthanoid series).

    7. Seventh Period: Incomplete, theoretically $32$ elements (includes $5f$ Actinoid series).

• Lanthanoids and Actinoids: Placed in separate panels at the bottom to maintain the table's structure and group similarity.

Nomenclature of Elements with Atomic Numbers ($Z$) > 100

• Rationale: To resolve discovery disputes (e.g., Soviet vs. American scientists naming element 104 as Kurchatovium vs. Rutherfordium), IUPAC introduced a systematic naming method until a permanent name is voted on.

• Numerical Roots:

  • 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 Convention: Put the roots together in order of digits and add "ium" at the end.

• Examples:

  • $101$: Unnilunium ($Unu$) -> Official: Mendelevium ($Md$)

  • $104$: Unnilquadium ($Unq$) -> Official: Rutherfordium ($Rf$)

  • $106$: Unnilhexium ($Unh$) -> Official: Seaborgium ($Sg$)

  • $110$: Ununnillium ($Uun$) -> Official: Darmstadtium ($Ds$)

  • $120$: Unbinilium ($Ubn$)

Electronic Configurations and the Division into Blocks

• Foundational Principle: The distribution of electrons into orbitals (Aufbau Principle) determines an element's location.

• Periods ($n$):

  • The period indicates the value of $n$ for the outermost shell.

  • Number of elements in a period is $\text{twice}$ the number of available atomic orbitals.

• The Four Blocks ($s, p, d, f$):

3.6.1 The s-Block Elements

• Groups: 1 (alkali metals) and 2 (alkaline earth metals).

• Configuration: Outer configuration is $ns^1$ or $ns^2$.

• Properties: Reactive metals with low ionization enthalpies. Lose electrons easily to form $1+$ or $2+$ ions. Reactivity increases down the group. Mostly form ionic compounds (excluding Lithium and Beryllium).

3.6.2 The p-Block Elements

• Groups: 13 to 18. Together with s-block, they are called Main Group or Representative Elements.

• Configuration: Outer configuration from $ns^2 np^1$ to $ns^2 np^6$.

• Noble Gases (Group 18): Have closed shells ($ns^2 np^6$), leading to low chemical reactivity.

• Important Families: Group 16 (Chalcogens) and Group 17 (Halogens). High negative electron gain enthalpies.

• Trends: Non-metallic character increases across a period; metallic character increases down a group.

3.6.3 The d-Block Elements (Transition elements)

• Groups: 3 to 12.

• Configuration: General outer electronic configuration $(n-1)d^{1-10} ns^{0-2}$. Filling of inner $d$ orbitals.

• Properties: All are metals. Usually form colored ions, exhibit variable oxidation states, and paramagnetism. Used as catalysts. Zinc ($Zn$), Cadmium ($Cd$), and Mercury ($Hg$) are exceptions and do not show typical transition behavior due to completed $(n-1)d^{10}$ shells.

3.6.4 The f-Block Elements (Inner-Transition elements)

• Series: Lanthanoids (Ce $Z=58$ to Lu $Z=71$) and Actinoids (Th $Z=90$ to Lr $Z=103$).

• Configuration: $(n-2)f^{1-14} (n-1)d^{0-1} ns^2$.

• Properties: All are metals. Actinoids are radioactive; many are man-made (transuranium elements) and synthesized in minute quantities.

Metals, Non-metals, and Metalloids

• Metals: > 78% of elements. Located on the left. High melting/boiling points, good conductors, malleable, and ductile. Mercury ($Hg$) is liquid. Gallium ($Ga$) and Caesium ($Cs$) have very low melting points ($303\,K$ and $302\,K$).

• Non-metals: Located top right. Mostly solids or gases. Poor conductors, brittle. Boron and Carbon are exceptions (high melting points).

• Metalloids (Semi-metals): Border the zig-zag line. Examples: Silicon ($Si$), Germanium ($Ge$), Arsenic ($As$), Antimony ($Sb$), Tellurium ($Te$). Display properties of both metals and non-metals.

Periodic Trends in Physical Properties

(a) Atomic Radius

• Measurement: Covalent Radius (half the distance between two atoms in a single-bonded molecule) or Metallic Radius (half the distance between metal cores in a crystal).

  • Example: Chlorine ($Cl_2$) bond distance is $198\,pm$, atomic radius is $99\,pm$.

  • Example: Copper solid adjacent atom distance is $256\,pm$, metallic radius is $128\,pm$.

• Trends:

  • Across a Period: Radius decreases as effective nuclear charge increases, pulling electrons closer.

  • Down a Group: Radius increases as the principal quantum number ($n$) increases and inner electron shielding offsets nuclear charge.

(b) Ionic Radius

• Cations: Smaller than their parent atom ($Na$ is $186\,pm$, $Na^+$ is $95\,pm$ due to fewer electrons and same nuclear charge).

• Anions: Larger than their parent atom ($F$ is $64\,pm$, $F^-$ is $136\,pm$ due to electron repulsion and decreased effective nuclear charge).

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

(c) Ionization Enthalpy ($\Delta_i H$)

• Definition: Energy required to remove an electron from an isolated gaseous atom ($X$) in ground state.

  • $X(g) \rightarrow X^+(g) + e^-$

• Successive Enthalpies: \Delta_i H_1 < \Delta_i H_2 < \Delta_i H_3 because it is harder to remove an electron from a positive ion.

• Trends: Generally increases across a period (nuclear charge) and decreases down a group (shielding/distance).

• Anomalies:

  • Boron ($Z=5$) vs Beryllium ($Z=4$): $B$ has lower $IE$ because it is easier to remove a $2p$ electron (more shielded) than a $2s$ electron.

  • Oxygen ($Z=8$) vs Nitrogen ($Z=7$): $O$ has lower $IE$ because of electron-electron repulsion in the doubly occupied $2p$ orbital.

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

• Definition: Enthalpy change when an electron is added to a neutral gaseous atom ($X$).

  • $X(g) + e^- \rightarrow X^-(g)$

• Trends: Becomes more negative across a period. Usually becomes less negative down a group.

• Exceptions: $O$ and $F$ have less negative enthalpies than $S$ and $Cl$ because the small size of $n=2$ leads to significant electron-electron repulsion for the incoming electron.

(e) Electronegativity

• Definition: Ability of an atom in a compound to attract shared electrons.

• Scales: Pauling scale (most common), Mulliken-Jaffe, Allred-Rochow.

• Pauling scale: Fluorine is the most electronegative ($4.0$).

• Trends: Increases across a period, decreases down a group. Inversely related to metallic character.

Periodic Trends in Chemical Properties

• Valence / Oxidation States: Usually equals the number of valence electrons or eight minus that number.

  • Example: Fluorine always exhibits $-1$; Oxygen exhibits $+2$ in $OF_2$ but $-2$ in $Na_2O$.

• Anomalous Properties of Second Period Elements: The first member of each group (Li, Be, B, C, N, O, F) differs from its siblings due to small size, large charge/radius ratio, high electronegativity, and lack of $d$ orbitals (limiting covalency to 4).

• Diagonal Relationship: Similarity between Lithium and Magnesium, or Beryllium and Aluminium.

• Chemical Reactivity:

  • High at the extremes of a period, low in the center.

  • Extreme left (alkali metals) reactive via electron loss; extreme right (halogens) reactive via electron gain.

• Oxide Character:

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

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

  • Center elements form Amphoteric (e.g., $Al_2O_3$) or Neutral (e.g., $CO, NO, N_2O$) oxides.

Questions & Discussion

• Q: Why do $O^{2-}, F^-, Na^+, Mg^{2+}$ have different sizes despite being isoelectronic?

• A: They have different nuclear charges. $Mg^{2+}$ has the highest nuclear charge ($12$ protons), attracting the $10$ electrons most strongly, making it the smallest. $O^{2-}$ has the least ($8$ protons) and is the largest.

• Q: What is the IUPAC name for element $120$?

• A: Roots for 1, 2, and 0 are "un", "bi", and "nil". Name: Unbinilium. Symbol: Ubn.

• Q: Justify the presence of 18 elements in the 5th period.

• A: For $n=5$, the orbitals available for filling are $5s, 4d, \text{and } 5p$. The total number of orbitals is $1+5+3 = 9$. As each orbital holds $2$ electrons, total elements = $18$.

• Q: Show that $Na_2O$ is basic and $Cl_2O_7$ is acidic.

• A: By reacting with water:

  • $Na_2O + H_2O \rightarrow 2NaOH$ (Strong base)

  • $Cl_2O_7 + H_2O \rightarrow 2HClO_4$ (Strong acid)

• Henry Moseley (1913): Observed the X-ray spectra of elements. He found that a plot of $\sqrt{\nu}$ (where $\nu$ is frequency) against atomic number ($Z$) gave a straight line, unlike the plot against atomic mass. He proved that atomic number is a more fundamental property.

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

• Significance of Atomic Number: $Z = \text{nuclear charge (protons)} = \text{number of electrons in a neutral atom}$. Periodicity is a consequence of periodic variation in electronic configurations.

• Structure of the "Long Form" Table:

  • Periods: Horizontal rows (7 in total). The period number corresponds to the highest principal quantum number ($n$).

  • Groups: Vertical columns (numbered 1-18 according to IUPAC, replacing IA-VIIA, VIII, IB-VIIB, 0). Groups contain elements with similar outer electronic configurations.

  • Capacity of Periods:

    1. First Period: $2$ elements ($1s$ shell filled).

    2. Second Period: $8$ elements.

    3. Third Period: $8$ elements.

    4. Fourth Period: $18$ elements (includes $3d$ transition series).

    5. Fifth Period: $18$ elements (includes $4d$ transition series).

    6. Sixth Period: $32$ elements (includes $4f$ Lanthanoid series).

    7. Seventh Period: Incomplete, theoretically $32$ elements (includes $5f$ Actinoid series).

• Lanthanoids and Actinoids: Placed in separate panels at the bottom to maintain the table's structure and group similarity.

The lanthanide series consists of elements with atomic numbers from 58 (Cerium, Ce) to 71 (Lutetium, Lu). The actinide series includes elements with atomic numbers from 89 (Actinium, Ac) to 103 (Lawrencium, Lr). Transition elements, which are also known as d-block elements, cover atomic numbers 21 (Scandium, Sc) to 30 (Zinc, Zn) and extend from 39 (Yttrium, Y) to 48 (Cadmium, Cd), and from 57 (Lanthanum, La) to 80 (Zinc, Zn) including some elements in the f-block as well.