Classification of Elements and Periodicity in Properties Study Guide

General Principles of the Modern Periodic Table and Historical Context

• Principles of Classification:

  • The $4f$-series (lanthanides) and $5f$-series (actinides) are placed separately at the bottom of the periodic table. This is done specifically to preserve the principle of classification and maintain the structural integrity of the periodic table.

  • s-block elements are generally too reactive to be found in their pure (elemental) form in nature; they usually exist as compounds.

• Development of Periodic Law:

  • Mendeleev's Periodic Law: According to Mendeleev, the properties of elements are a periodic function of their atomic masses.

  • Modern Periodic Law: Proposed based on the work of Henry Moseley, stating that the variation in properties of elements is related to their atomic numbers.

  • Moseley's Experiment: Henry Moseley studied the characteristic X-ray spectra of elements and found a linear relationship between the square root of the frequency ($\sqrt{\nu}$) of emitted X-rays and the atomic number ($Z$). The correct graphical representation is a straight line for $\sqrt{\nu}$ versus $Z$.

• D.I. Mendeleev Specifics:

  • He authored the definitive textbook titled "Principles of Chemistry."

  • Element with atomic number $101$ (Mendelevium) is named in his honor.

  • Historical Note: At the time Mendeleev proposed his periodic table, the internal structure of the atom was not known.

• Lother Meyer’s Curve Details:

  • Plotted between atomic weight and atomic volume.

  • Alkali metals occupy the maxima (peaks) of the curve.

  • Transition metals typically occupy the bottom portions of the curve.

  • Halogens are found on the ascending portions of the curve (Note: The transcript identifies an incorrect choice regarding halogens on descending portions).

IUPAC Nomenclature for Superheavy Elements (Z > 100)

• Naming Conventions:

  • The names are derived directly from the atomic number using roots: 0 = nil, 1 = un, 2 = bi, 3 = tri, 4 = quad, 5 = pent, 6 = hex, 7 = sept, 8 = oct, 9 = enn.

• Specific Element Identifications:

  • Atomic Number 101: Unnilunium (Symbol: $Unu$).

  • Atomic Number 102: Unnilbium (Symbol: $Unb$).

  • Atomic Number 103: Unniltrium (Symbol: $Unt$). Electronic configuration: $[Rn] 5f^{14} 6d^{1} 7s^{2}$.

  • Atomic Number 104: Unnilquadium (Symbol: $Unq$).

  • Atomic Number 107: Unnilseptium (Symbol: $Uns$).

  • Atomic Number 108: Unniloctium (Symbol: $Uno$).

  • Atomic Number 109: Unnilennium (Symbol: $Une$).

  • Atomic Number 111: Unununium (Symbol: $Uuu$). This element belongs to Group 11 of the periodic table.

  • Atomic Number 117: Ununseptium (Symbol: $Uus$).

  • Atomic Number 119: Ununennium (Symbol: $Uue$).

Electronic Configuration and Periodic Position

• Valence Shell Configurations:

  • $5s^2 5p^4$: Corresponds to an element in Group 16 ($10 + 2 + 4$) and Period 5.

  • $3s^2 3p^3$: Corresponds to Group 15 and Period 3 (Phosphorus, $Z=15$).

  • Identification based on Atomic Numbers:

    • $Z=15$ (Phosphorus): Group 15, contains $5$ valence electrons, and has a valency of $3$.

    • $Z=33$ (Arsenic): Metalloid.

    • $Z=53$ (Iodine): Non-metal.

    • $Z=83$ (Bismuth): Metal.

• Group and Period Trends:

  • Similarity in chemical properties within a group is determined by the number of valence electrons.

  • Vertical Relationships: An element in Period 4 and Group 16 is Selenium ($Se$). The element directly above it in the group (Period 3, Group 16) has the configuration $3s^2 3p^4$.

  • Group 17 Characteristics: If the outermost configuration of a Group 17 element is $4s^2 4p^5$ (Bromine), the element directly below it (Iodine) will have the configuration $5s^2 5p^5$.

  • Diagonal Relationships: A Group 13 element in Period 4 is Gallium ($Ga$, $4s^2 4p^1$). The element placed diagonally to it in Period 5 of the p-block is Tin ($Sn$), with configuration $[Kr] 4d^{10} 5s^2 5p^2$.

Periodic Trends 1: Atomic and Ionic Radius

• Radius Trends:

  • General Order: Cations are always smaller than their parent neutral atoms, which are smaller than their respective anions (e.g., I^+ < I < I^-).

  • Atomic Radius Order: Example order for certain elements like Cerium ($Ce$), Europium ($Eu$), Holmium ($Ho$), and Nitrogen ($N$) is Eu > Ce > Ho > N.

• Isoelectronic Species:

  • Species having the same number of electrons (e.g., $N^{3-}, O^{2-}, F^-, Na^+, Mg^{2+}, Al^{3+}$).

  • Size in isoelectronic species is determined by the nuclear charge; as nuclear charge increases, the radius decreases.

  • Radius Comparison Examples:

    • Mg^{2+} < Na^+ < F^- < O^{2-} < N^{3-}.

    • The largest difference in size in the series $N^{3-}, O^{2-}, F^-, Na^+, Mg^{2+}$ occurs between $N^{3-}$ and $Mg^{2+}$.

  • Calculated Values: The ionic radii for $N^{3-}$, $O^{2-}$, and $F^-$ are approximately 1.71\,&A, 1.40\,&A, and 1.36\,&A respectively.

Periodic Trends 2: Ionization Enthalpy (IE) and Potential

• Ionization Enthalpy Trends:

  • Typically increases across a period and decreases down a group.

  • Exceptions (Half-filled/Full-filled shells):

    • IE_1(N) > IE_1(O) because Nitrogen has a stable half-filled $2p^3$ configuration, while Oxygen ($2p^4$) experiences more electron-electron repulsion.

    • IE_1(Be) > IE_1(B) because the electron is removed from a stable $2s^2$ orbital in Beryllium.

  • First Ionization Enthalpy Order: For second-period elements Be, B, N, O, the order is B < Be < O < N.

  • IE Order for Mg, Al, S, P: Al < Mg < S < P.

• Successive Ionization Energies:

  • The difference between $IE_1$ and $IE_2$ is greatest for Group 1 elements (like Potassium, $K$) because removing the second electron requires breaking a stable noble gas configuration.

  • Alkaline Earth Metals ($M^{2+}$): Form dipositive ions because the difference $IE_2 - IE_1$ is small enough (typically < 10\,eV) that the lattice energy or hydration energy can compensate for it.

• Numerical Data:

  • If the first ionization potential of $Na$ is $5.1\,eV$, the electron gain enthalpy of $Na^+$ would be $-5.1\,eV$.

  • Highest $IE_1$ among alkali metals ($Na, K, Rb$) and Scandium ($Sc$) is found in Sc.

Periodic Trends 3: Electron Gain Enthalpy and Electronegativity

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

  • Halogens: Chlorine ($Cl$) has a more negative electron gain enthalpy than Fluorine ($F$) due to less electron-electron repulsion in the larger $3p$ orbital compared to the compact $2p$ orbital of Fluorine.

  • Halogen Values ($kJ/mol$): Cl (-349) > F (-333) > Br (-325) > I (-296).

  • Successive Values: The first $\Delta_{eg}H$ of Oxygen is negative ($-141\,kJ/mol$), but the second $\Delta_{eg}H$ is always positive due to the high repulsion between the added electron and the already negative $O^{-}$ ion.

• Electronegativity (Pauling Scale):

  • Correct Trends:

    • P < S (Electronegativity increases across a period).

    • Ga < Se.

    • Br > C > At > P.

Miscellaneous Classifications and Relationships

• Metalloids: Elements exhibiting properties of both metals and non-metals. Examples include Tellurium ($Te$), Arsenic ($As$), and Silicon ($Si$). Bismuth ($Bi$) and Lead ($Pb$) are metals.

• Gallium ($Ga$): A metal notable for having a very low melting point and a periodic position close to the metalloids.

• Noble Gas Configurations: Ions such as $Sr^{2+} (Z=38)$ and $Cs^+ (Z=55)$ achieve noble gas configurations, whereas transition or post-transition ions like $Fe^{2+}$ or $Pb^{2+}$ do not.

• Anomalous Pairs (Mendeleev): Pairs that violated the atomic mass order, such as Argon ($Ar$) and Potassium ($K$), Cobalt ($Co$) and Nickel ($Ni$), and Tellurium ($Te$) and Iodine ($I$). The pair Aluminium ($Al$) and Silicon ($Si$) was not an anomalous pair.

• Dobereiner Triads: Groups of three elements where the middle element's atomic weight is the average of the other two. H, F, Cl is not a recognized Dobereiner triad.

• Representative Elements: These belong to the s-block and p-block (Groups 1, 2, and 13–18).

• Typical Elements: Elements belonging to the 3rd period are considered typical elements as they represent the properties of their respective groups.

• Diagonal Relationships:

  • $Li-Mg$

  • $Be-Al$

  • $B-Si$