Unit 3: Classification of Elements and Periodicity in Properties Notes

Introduction to Periodic Classification

• The Periodic Table provides a succinct organization of the entirety of chemistry, demonstrating that elements are not a random cluster of entities but instead display trends and lie together in families.
• Awareness of the Periodic Table is essential for understanding how the world is built from fundamental chemical elements.
• Classification is necessary because elements are the basic units of all matter. In 1800, only 31 elements were known. By 1865, this number more than doubled to 63. Currently, 114 elements are recognized, with the most recently discovered being man-made substances.
• Organizing elements into a systematic classification allows scientists to rationalize known chemical facts and predict new ones for future research.

Genesis of Periodic Classification

• Johann Dobereiner (Early 1800s): The German chemist was the first to identify trends among elements. By 1829, he noted similarities in physical and chemical properties of several groups of three elements, known as Triads.
• Law of Triads: He observed that the middle element of a triad had an atomic weight approximately halfway between the other two. Additionally, its properties were intermediate between the other two members.
• Dobereiner’s Triads Examples:
  • $Li$ (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 known elements in increasing order of atomic weights in a cylindrical table to display property recurrence.
• John Alexander Newlands (1865): The English chemist propounded the Law of Octaves. He arranged elements by increasing atomic weight and noted that every eighth element possessed properties similar to the first.
• Newlands' Law of Octaves: This relationship resembled musical octaves. It was found to be true only for elements up to calcium. He was later awarded the Davy Medal in 1887 by the Royal Society, London.
• Table of Newlands' Octaves:
  • $Li$ (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)

The Mendeleev and Lothar Meyer Periodic Periodicity

• Dmitri Mendeleev (1834–1907) and Lothar Meyer (1830–1895): Working independently in 1869, both proposed that physical and chemical properties of elements show periodic similarities when arranged by increasing atomic weight.
• Lothar Meyer's Plot: He plotted physical properties like atomic volume, melting point, and boiling point against atomic weight, obtaining a periodically repeated pattern with changing lengths of the repeating pattern. By 1868, he developed a table resembling the Modern Periodic Table.
• Mendeleev’s Periodic Law: ‘The properties of the elements are a periodic function of their atomic weights.’
• Mendeleev’s Systematic Arrangement: Elements were placed in horizontal rows and vertical columns (groups) such that elements with similar properties occupied the same column. He relied on empirical formulas and properties of compounds.
• Scientific Boldness: Mendeleev ignored strict atomic weight order if it didn't fit chemical similarity. For example, Iodine (lower atomic weight) was placed in Group VII after Tellurium (Group VI) to match Fluorine, Chlorine, and Bromine.
• Predictions of New Elements: Mendeleev left gaps for undiscovered elements, calling them Eka-Aluminium and Eka-Silicon.
• Comparison of Predicted and Found Properties:
  • Eka-Aluminium (predicted) / Gallium (found): Atomic weight: 68 / 70; Density: $5.9\,g/cm^3$ / $5.94\,g/cm^3$; Melting point: Low / $302.93\,K$; Oxide: $E_2O_3$ / $Ga_2O_3$; Chloride: $ECl_3$ / $GaCl_3$.
  • Eka-Silicon (predicted) / Germanium (found): Atomic weight: 72 / 72.6; Density: $5.5\,g/cm^3$ / $5.36\,g/cm^3$; Melting point: High / $1231\,K$; Oxide: $EO_2$ / $GeO_2$; Chloride: $ECl_4$ / $GeCl_4$.

The Modern Periodic Law

• Henry Moseley (1913): The English physicist observed regularities in characteristic X-ray spectra. A plot of $\sqrt{\nu}$ (where $\nu$ is the frequency) against atomic number ($Z$) produced a straight line, which was not the case for atomic mass.
• Atomic Number ($Z$): Moseley demonstrated that atomic number is a more fundamental property than atomic mass.
• Modern Periodic Law Statement: ‘The physical and chemical properties of the elements are periodic functions of their atomic numbers.’
• Atomic Number Definition: It is equal to the nuclear charge (number of protons) or the number of electrons in a neutral atom.
• Natural and Artificial Elements: 94 elements occur naturally (including Neptunium, Plutonium, Actinium, and Protactinium found in pitchblende). Modern classification also includes artificially produced short-lived elements.

Present Form of the Periodic Table

• The "Long Form" of the periodic table is the most widely used modern version.
• Groups: Vertical columns containing elements with similar outer electronic configurations. IUPAC numbers them 1 to 18 (replacing older IA-VIIB notation).
• Periods: Horizontal rows. There are 7 periods. The period number corresponds to the highest principal quantum number ($n$).
  • Period 1: 2 elements.
  • Period 2 and 3: 8 elements each.
  • Period 4 and 5: 18 elements each.
  • Period 6: 32 elements.
  • Period 7: Incomplete (would have theoretical max of 32 elements based on quantum numbers).
• Lanthanoids and Actinoids: 14 elements from both the 6th and 7th periods are placed in separate panels at the bottom. This reconfiguration was largely due to Glenn T. Seaborg’s work (awarded Nobel Prize in 1951).

Nomenclature of Elements with Atomic Numbers $Z > 100$

• To prevent disputes over naming (such as the American/Soviet rivalry over element 104, Rutherfordium vs. Kurchatovium), IUPAC recommended a systematic nomenclature.
• Numerical Roots: 0: nil; 1: un; 2: bi; 3: tri; 4: quad; 5: pent; 6: hex; 7: sept; 8: oct; 9: enn.
• Naming Rule: Roots are combined in order of digits plus the suffix ‘ium’. Symbols use the three root-initial letters.
• Example (Element 120): Roots are un (1), bi (2), and nil (0). The name is Unbinilium, and the symbol is $Ubn$.
• Permanent Names: These reflect the country of discovery or honor a scientist (e.g., Seaborgium for $Z = 106$).

Electronic Configuration in the Periodic Table

• Aufbau Principle: The distribution of electrons into orbitals (s, p, d, f) defines the element’s location.
• Filling of Shells: Each successive period involves filling the next higher principal energy level ($n = 1, n = 2, \dots$).
• Orbital Progression:
  • $n = 1$: 1s orbital fills (H to He).
  • $n = 2$: 2s and 2p orbitals fill (Li to Ne).
  • $n = 3$: 3s and 3p orbitals fill (Na to Ar).
  • $n = 4$: 4s starts, then 3d (transition series), then 4p (K to Kr, total 18 elements).
  • $n = 6$: Fills 6s, 4f, 5d, and 6p. The 4f-inner transition series (Lanthanoids) starts at Cerium ($Z = 58$) and ends at Lutetium ($Z = 71$).
  • $n = 7$: Fills 7s, 5f, 6d, and 7p. The 5f-inner transition series (Actinoids) starts after Actinium ($Z = 89$).
• Groups: Members of the same group have identical valence shell electronic configurations (e.g., alkali metals in Group 1 have $ns^1$ configuration).

Grouping Elements by Blocks (s, p, d, f)

• The s-Block Elements: Groups 1 (alkali metals) and 2 (alkaline earth metals) with $ns^1$ and $ns^2$ configurations. They are reactive metals with low ionization enthalpies and form predominantly ionic compounds (except Li and Be).
• The p-Block Elements: Groups 13 to 18 ($ns^2 np^1$ to $ns^2 np^6$). Together with s-block, these are Representative or Main Group elements. Group 18 (noble gases) has a stable closed shell. Group 17 (halogens) and Group 16 (chalcogens) are highly reactive non-metals.
• The d-Block Elements (Transition Elements): Groups 3 to 12. General configuration: $(n-1)d^{1-10} ns^{0-2}$. They are all metals, usually forming colored ions and exhibiting variable valency and paramagnetism. Zn, Cd, and Hg are exceptions as they lack typical transition properties.
• The f-Block Elements (Inner-Transition Elements): Lanthanoids ($Ce Z = 58$ to $Lu Z = 71$) and Actinoids ($Th Z = 90$ to $Lr Z = 103$). General configuration: $(n-2)f^{1-14} (n-1)d^{0-1} ns^2$. Late actinoids are radioactive and found only in trace amounts. Elements after Uranium are called Transuranium elements.

Metals, Non-metals, and Metalloids

• Metals: Over 78% of elements. Located on the left. Highly solid at room temperature (Mercury is the exception). They are malleable, ductile, and good conductors.
• Non-metals: Top right side. Usually solids or gases with low melting/boiling points (except Carbon and Boron).
• Metalloids (Semi-metals): Bordering the thick zigzag line (Si, Ge, As, Sb, Te). They exhibit properties of both metals and non-metals.

Periodic Trends in Physical Properties

• Atomic Radius: The size of an individual atom (~$1.2\,\text{…}$ or $1.2 \times 10^{-10}\,m$) cannot be measured precisely because boundaries are not sharp.
  • Covalent Radius: Half the distance between two atoms joined by a single bond (e.g., $Cl_2$ bond distance is $198\,pm$; radius = $99\,pm$).
  • Metallic Radius: Half the internuclear distance between metal cores in a crystal (e.g., Copper = $128\,pm$).
  • Period Trend: Decreases across a period due to increased effective nuclear charge pulling electrons closer.
  • Group Trend: Increases down a group as principal quantum number ($n$) increases and inner electrons shield the nucleus.
• Ionic Radius:
  • Cations: Smaller than parent atoms (fewer electrons, same charge).
  • Anions: Larger than parent atoms (increased electron repulsion, decreased effective nuclear charge). Fluorine ($64\,pm$) vs. Fluoride ion ($136\,pm$).
  • Isoelectronic Species: Atoms/ions with the same total electrons (e.g., $O^{2-}, F^-, Na^+, Mg^{2+}$). The radius decreases as positive nuclear charge increases.
• Ionization Enthalpy ($\Delta_i H$): Energy required to remove an electron from an isolated gaseous atom ($X(g) \rightarrow X^+(g) + e^-$).
  • Trend: Generally increases across a period; decreases down a group.
  • Anomalies: Boron has lower $\Delta_i H$ than Beryllium because 2p electrons are more easily removed than 2s. Nitrogen has higher $\Delta_i H$ than Oxygen due to the stability of half-filled p-orbitals in Nitrogen.
• Electron Gain Enthalpy ($\Delta_{eg} H$): Enthalpy change when adding an electron to a gaseous atom.
  • Trend: Becomes more negative across a period toward halogens. Noble gases have positive values.
  • Anomaly: O and F are less negative than S and Cl because of electron-electron repulsion in the smaller $n=2$ shell.
• Electronegativity: Ability to attract shared electrons. Pauling scale (Fluorine = 4.0, Caesium = 0.7).
  • Trend: Increases across a period; decreases down a group.

Periodic Trends in Chemical Properties

• Valence: Number of electrons in the outermost shell or $8$ minus that number.
• Oxidation State: Charge based on electronegativity rules. In $OF_2$, Fluorine is $-1$ and Oxygen is $+2$, but in $Na_2O$, Oxygen is $-2$ and Sodium is $+1$.
• Anomalous Second Period Properties: Lithium, Beryllium, and Boron to Fluorine differ from their heavier congeners due to small size, high charge/radius ratio, high electronegativity, and lack of d-orbitals.
  • Diagonal Relationships: Lithium resembles Magnesium; Beryllium resembles Aluminium.
• Chemical Reactivity Trends: High at extremes (alkali metals lose electrons easily; halogens gain electrons easily).
  • Oxides: Basic at left extreme ($Na_2O$), acidic at right extreme ($Cl_2O_7$), amphoteric/neutral in the center ($Al_2O_3$, $NO$, $CO$).

Questions & Discussion

• Question: What would be the IUPAC name and symbol for the element with Z = 120?
• Response: The roots for 1, 2, and 0 are un, bi, and nil. Thus, the name is unbinilium and the symbol is Ubn.
• Question: How would you justify the presence of 18 elements in the 5th period?
• Response: For $n=5$, the orbitals available are 5s, 4d, and 5p. Total orbitals = $1 + 5 + 3 = 9$. As each orbital holds 2 electrons, the total number of elements is 18.
• Question: Which has the largest and smallest size among $Mg, Mg^{2+}, Al, Al^{3+}$?
• Response: Magnesium ($Mg$) is the largest. Aluminium ion ($Al^{3+}$) is the smallest due to its higher nuclear charge among isoelectronic cations and period trends.
• Question: Predicting compound formulas for (a) silicon and bromine; (b) aluminium and sulphur.
• Response: (a) Silicon (Valence 4) and Bromine (Valence 1) form $SiBr_4$. (b) Aluminium (Valence 3) and Sulphur (Valence 2) form $Al_2S_3$.