Classification of Elements and Periodicity in Properties Study Guide

Unit 3: Classification of Elements and Periodicity in Properties

• The Periodic Table is considered the most important concept in chemistry, serving as a support for students and a guide for research professionals.

• It provides a succinct organization of the chemical elements, demonstrating that they are not a random cluster but instead display trends and exist in families.

• Understanding the Periodic Table is essential to understanding the fundamental building blocks of chemistry: the chemical elements.

• This unit covers:

  • Historical development of grouping elements based on properties.

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

  • IUPAC nomenclature for elements with Z > 100.

  • Classification into s, p, d, and f blocks.

  • Periodic trends in physical properties (atomic/ionic radii, ionization enthalpy, electron gain enthalpy, electronegativity).

  • Periodicity in chemical properties (valence, reactivity, and occurrence in nature).

  • The relationship between ionization enthalpy and metallic character.

Why Do We Need to Classify Elements?

• Elements are the basic units of all matter.

• Historical growth in known elements:

  • In $1800$, only $31$ elements were known.

  • By $1865$, the number identified doubled to $63$.

  • At present, $114$ elements are formally recognized, with recently discovered ones being man-made.

• Systematic classification rationalizes known chemical facts and helps predict new information for further study, as studying individual elements and their innumerable compounds would be too difficult.

Genesis of Periodic Classification

• Classification is the result of observations and experiments by various scientists aimed at systematizing knowledge.

Johann Dobereiner and the Law of Triads (Early 1800s)

• In $1829$, Dobereiner noted similarities in physical and chemical properties of groups of three elements (Triads).

• The Law of Triads: The middle element of each Triad had an atomic weight approximately halfway between the other two, and its properties were intermediate.

• Table 3.1: Dobereiner’s Triads:

  • Triad 1: Lithium (Li, At. wt. $7$), Sodium (Na, At. wt. $23$), Potassium (K, At. wt. $39$).

  • Triad 2: Calcium (Ca, At. wt. $40$), Strontium (Sr, At. wt. $88$), Barium (Ba, At. wt. $137$).

  • Triad 3: Chlorine (Cl, At. wt. $35.5$), Bromine (Br, At. wt. $80$), Iodine (I, At. wt. $127$).

• This was dismissed as coincidence because it only worked for a few elements.

A.E.B. de Chancourtois (1862)

• A French geologist who arranged known elements in order of increasing atomic weights.

• He created a cylindrical table to display the periodic recurrence of properties, but it did not receive much attention.

John Alexander Newlands and the Law of Octaves (1865)

• Arranged elements in increasing order of atomic weights.

• Law of Octaves: Every eighth element has properties similar to the first element, analogous to musical octaves.

• Table 3.2: Newlands’ Octaves:

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

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

  • Followed by: K ($39$), Ca ($40$).

• The law seemed true only up to Calcium ($Ca$). Newlands was later awarded the Davy Medal in $1887$ by the Royal Society, London.

Dmitri Mendeleev and Lothar Meyer (1869)

• Working independently, both proposed that physical and chemical properties are periodic functions of atomic weights.

• Lothar Meyer: Plotted physical properties (atomic volume, melting point, boiling point) against atomic weight. By $1868$, he developed a table resembling the Modern Periodic Table, but published after Mendeleev.

• Dmitri Mendeleev: Credited with the first published Periodic Law.

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

• Mendeleev’s Table Characteristics:

  • Elements arranged in horizontal rows and vertical columns (groups).

  • Relied on empirical formulas and properties of compounds formed by elements.

  • He prioritized similarity in properties over strict atomic weight order. Example: Iodine ($I$) has a lower atomic weight than Tellurium ($Te$) but was placed in Group VII with Fluorine and Chlorine.

  • Left gaps for undiscovered elements, such as Eka-Aluminium and Eka-Silicon.

• Success/Predictions (Table 3.3):

  • Eka-aluminium (predicted): At. wt. $≈ 68$, Density $5.9\,g/cm^3$, Oxide $E_2O_3$, Chloride $ECl_3$.

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

  • Eka-silicon (predicted): At. wt. $72$, Density $5.5\,g/cm^3$, Oxide $EO_2$, Chloride $ECl_4$.

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

Modern Periodic Law and Present Form

• Henry Moseley (1913): Observed characteristic X-ray spectra. A plot of $\sqrt{\nu}$ (frequency) against atomic number ($Z$) was a straight line, while $\sqrt{\nu}$ vs. atomic mass was 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."

• This law revealed analogies among naturally occurring elements, including those in pitchblende (Uranium ore) such as Neptunium ($Np$), Plutonium ($Pu$), Actinium ($Ac$), and Protactinium ($Pa$).

Features of the Long Form of the Periodic Table

• Horizontal Rows: Called Periods. There are $7$ periods. The period number matches the highest principal quantum number ($n$).

• Vertical Columns: Called Groups or Families. Elements have similar outer electronic configurations.

• IUPAC Group Numbering: $1$ to $18$ (replaces old IA-VIIA, VIII, IB-VIIB, and $0$).

• Period Membership:

  • Period 1: $2$ elements ($1s$ level).

  • Period 2 & 3: $8$ elements each ($2s, 2p$ and $3s, 3p$ levels).

  • Period 4 & 5: $18$ elements each ($4s, 3d, 4p$ and $5s, 4d, 5p$ levels).

  • Period 6: $32$ elements ($6s, 4f, 5d, 6p$ levels).

  • Period 7: Incomplete (theoretically $32$ elements max).

• Separate Panels: Lanthanoids ($Z = 58-71$) and Actinoids ($Z = 90-103$) are placed at the bottom. This reconfiguration was led by Glenn T. Seaborg (Nobel Prize $1951$).

Nomenclature of Elements with Z > 100

• Systematic names are derived directly from the atomic number until the name is officially recognized by IUPAC.

• Numerical Roots (Table 3.4):

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

• Rule: Roots are combined in order of digits, ending with "ium".

• Example (Problem 3.1): Element $120$ is unbinilium with symbol Ubn.

• Official Examples:

  • $101$: Unnilunium (Mendelevium, Md)

  • $104$: Unnilquadium (Rutherfordium, Rf)

  • $106$: Unnilhexium (Seaborgium, Sg)

  • $118$: Ununoctium (Oganesson, Og)

Electronic Configuration

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

Configurations in Periods

• Period 1 ($n=1$): Starts with Hydrogen ($1s^1$) and ends with Helium ($1s^2$).

• Period 2 ($n=2$): Lithium ($2s^1$) to Neon ($2s^2 2p^6$).

• Period 4 ($n=4$): Transition series ($3d$) starts at Scandium ($Z=21, 3d^1 4s^2$) and ends at Zinc ($Z=30, 3d^{10} 4s^2$).

• Period 6 ($n=6$): Lanthanoid series ($4f$) starts at Cerium ($Z=58$) and ends at Lutetium ($Z=71$).

• Period 7 ($n=7$): Actinoid series ($5f$) starts after Actinium ($Z=89$).

Groupwise Electronic Configurations

• Group 1 (Alkali metals) have $ns^1$ configuration: Li ($2s^1$), Na ($3s^1$), K ($4s^1$), Rb ($5s^1$), Cs ($6s^1$), Fr ($7s^1$).

• Similarity in properties arises from having the same number and distribution of electrons in outermost orbitals.

Blocks in the Periodic Table

s-Block Elements

• Comprise Group 1 (alkali metals) and Group 2 (alkaline earth metals).

• Outermost configurations: $ns^1$ and $ns^2$.

• Characteristics: Reactive metals, low ionization enthalpies, form $1+$ or $2+$ ions. Reactivity increases down the group. Compounds are predominantly ionic (except Li and Be).

p-Block Elements (Representative Elements)

• Comprise Groups $13$ to $18$.

• Outermost configuration: $ns^2 np^1$ to $ns^2 np^6$.

• Group 18: Noble gases; stable closed shells, very low reactivity.

• Group 17: Halogens; highly negative electron gain enthalpy.

• Group 16: Chalcogens.

• Metallic character increases down the group; non-metallic character increases left to right.

d-Block Elements (Transition Elements)

• Comprise Groups $3$ to $12$.

• General configuration: $(n-1)d^{1–10} ns^{0–2}$ (except Pd at $4d^{10} 5s^0$).

• Characteristics: All are metals, form coloured ions, variable oxidation states, paramagnetism, catalytic properties.

• Exceptions: Zn, Cd, and Hg ($(n-1)d^{10} ns^2$) do not show most transition properties.

f-Block Elements (Inner-Transition Elements)

• Lanthanoids ($Z=58-71$) and Actinoids ($Z=90-103$).

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

• All are metals. Actinoids are radioactive and exhibit complex oxidation states. Elements after Uranium ($Z=92$) are called Transuranium Elements.

Metals, Non-metals, and Metalloids

• Metals: Over $78\%$ of elements. Left side. Solids (except Hg; Ga and Cs have low M.P.). High conductivity, malleable, ductile.

• Non-metals: Top right side. Usually gases or solids with low M.P./B.P. (except Boron and Carbon). Poor conductors; brittle.

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

Periodic Trends in Physical Properties

Atomic Radius

• Calculated size based on distance between atoms in a combined state.

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

• Metallic Radius: Half the internuclear distance between metal cores (e.g., $Cu$ distance is $256\,pm$, radius is $128\,pm$).

• Trends:

  • Decreases across a period (effective nuclear charge increase pulls electrons closer).

  • Increases down a group (principal quantum number $n$ increases; inner levels shield outer electrons).

• Noble Gases: Radii are larger because they are non-bonded (van der Waals radii).

Ionic Radius

• Estimated from distances between ions in crystals.

• Cations: Smaller than parent atom (fewer electrons, same charge).

• Anions: Larger than parent atom (increased electron-electron repulsion, decreased effective nuclear charge).

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

Ionization Enthalpy ($\Delta_i H$)

• Energy to remove an electron from an isolated gaseous atom in ground state: $X(g) \rightarrow X^+(g) + e^-$.

• Units: $kJ\,mol^{-1}$.

• Successive enthalpies (\Delta_i H_1 < \Delta_i H_2 < \Delta_i H_3).

• Trends:

  • Increases across a period due to increased nuclear charge.

  • Decreases down a group due to increased shielding and distance from nucleus.

• Specific Anomalies:

  • Be > B: Removing $2s$ electron (Be) requires more energy than $2p$ electron (B) because $s$ orbitals penetrate closer to the nucleus.

  • N > O: Nitrogen has half-filled stable $2p^3$ subshell. Oxygen ($2p^4$) has increased electron-electron repulsion in one orbital, making electron removal easier.

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

• Enthalpy change when an electron is added to neutral gaseous atom: $X(g) + e^- \rightarrow X^-(g)$.

• Trends:

  • Becomes more negative across a period (easier to add to smaller atoms).

  • Becomes less negative down a group (outer shell further from nucleus).

  • Halogens: Highly negative (attain noble gas configuration).

  • Noble Gases: Large positive values (electron must enter higher $n$ level).

  • O and F Exception: Less negative than S and Cl because the smaller $n=2$ level in O/F causes significant repulsion compared to the larger $n=3$ level in S/Cl.

Electronegativity

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

• Measured on Pauling scale (Fluorine = $4.0$, Cesium = $0.7$).

• Trends:

  • Increases across a period (related to decreasing atomic radius).

  • Decreases down a group (related to increasing atomic radius).

• Inverse relation to metallic character: Non-metals are electronegative; metals are electropositive.

Periodic Trends in Chemical Properties

Valence and Oxidation States

• Valence: Often equals the number of electrons in the outermost orbital or $8$ minus that number.

• Oxidation State: Charge an atom acquires based on electronegativity relative to other atoms.

  • Example: In $OF_2$, O is $+2$ (F is more electronegative). In $Na_2O$, O is $-2$.

• Formula prediction (Problem 3.8): Al (Gr 13) and S (Gr 16) form $Al_2S_3$.

Anomalous Properties of Second Period Elements

• The first members of groups (Li, Be, B to F) differ from lower members due to:

  • Small size.

  • High charge/radius ratio.

  • High electronegativity.

  • Lack of $d$ orbitals (limiting covalency to $4$).

• Diagonal Relationship: Similarity between Li and Mg, and Be and Al.

Chemical Reactivity Trends

• High reactivity at extreme ends of the period (Alkali metals lose electrons; Halogens gain electrons).

• Basic/Acidic Oxides:

  • Left side (Group 1): Basic oxides (e.g., $Na_2O$).

  • Right side (Group 17): Acidic oxides (e.g., $Cl_2O_{7}$).

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

• Amphoteric: Reacts with both acids and bases.

• Neutral: Neither acidic nor basic.