Classification of Elements and Periodicity in Properties

Historical Development and Genesis of Periodic Classification

• Need for Classification:

  • Elements serve as the basic building blocks of all matter.

  • Historical growth in identified elements:

  • Year 1800: 31 elements known.

  • Year 1865: 63 elements known (more than double in 65 years).

  • Present day: 114 elements officially recognized (up to 118 synthesized or discovered).

  • Classification organizes chemical facts, reveals underlying family relationships, and allows accurate prediction of undiscovered elements.

• Early Classification Attempts:

  • Johann Wolfgang Döbereiner (1829):

  • Proposed the Law of Triads.

  • Grouped elements into triads (groups of three) sharing physical and chemical characteristics.

  • Observed that the atomic weight of the middle element was approximately the arithmetic mean of the atomic weights of the other two.

  • Döbereiner's Triads:

    • Lithium ($\text{Li}$, atomic weight $7$), Sodium ($\text{Na}$, atomic weight $23$), Potassium ($\text{K}$, atomic weight $39$). Mean of Li and K = $\frac{7 + 39}{2} = 23$.

    • Calcium ($\text{Ca}$, atomic weight $40$), Strontium ($\text{Sr}$, atomic weight $88$), Barium ($\text{Ba}$, atomic weight $137$). Mean of Ca and Ba = $\frac{40 + 137}{2} = 88.5$.

    • Chlorine ($\text{Cl}$, atomic weight $35.5$), Bromine ($\text{Br}$, atomic weight $80$), Iodine ($\text{I}$, atomic weight $127$). Mean of Cl and I = $\frac{35.5 + 127}{2} = 81.25$.

  • Dismissed as a coincidence because it applied only to a few elements.

  • Alexandre-Émile Béguyer de Chancourtois (1862):

  • Arranged elements in order of increasing atomic weights on a cylindrical table (telluric helix).

  • Displayed periodic recurrence of properties, but received little attention.

  • John Alexander Reina Newlands (1865):

  • Formulated the Law of Octaves.

  • Arranged elements by increasing atomic weight and noted that every eighth element exhibited properties similar to the first (analogous to musical octaves).

  • Newlands' Octaves:

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

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

    • Row 3: $\text{K}\, (39)$, $\text{Ca}\, (40)$

  • Limitation: Applicable only for elements up to calcium ($\text{Ca}$).

  • Recognized later by being awarded the Davy Medal in 1887 by the Royal Society, London.

• Development of the Periodic Law:

  • Lothar Meyer (1869):

  • Plotted physical properties (atomic volume, melting point, boiling point) against atomic weight, obtaining a repeating pattern of changing length.

  • Developed a table of elements by 1868 closely resembling the modern periodic table, published after Mendeleev.

  • Dmitri Ivanovich Mendeleev (1869):

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

  • Arranged elements in horizontal rows (series) and vertical columns (groups) in order of increasing atomic weight.

  • Relied heavily on empirical formulas and chemical behaviors of compounds (especially oxides and hydrides).

  • Adjusted strict atomic weight order when necessary to keep similar elements together (e.g., Iodine with atomic weight 127 was placed in Group VII after Tellurium with atomic weight 128 in Group VI).

  • Predicted undiscovered elements and left gaps in his table:

    • Eka-Aluminium (discovered as Gallium in 1875).

    • Eka-Silicon (discovered as Germanium in 1886).

  • Mendeleev's Predictions vs. Experimental Findings:

    • Eka-Aluminium (predicted): Atomic weight = $68$, Density = $5.9\,\text{g/cm}^3$, Melting point = Low, Oxide = $\text{E}_2\text{O}_3$, Chloride = $\text{ECl}_3$.

    • Gallium (found): Atomic weight = $70$, Density = $5.94\,\text{g/cm}^3$, Melting point = $302.93\,\text{K}$, Oxide = $\text{Ga}_2\text{O}_3$, Chloride = $\text{GaCl}_3$.

    • Eka-Silicon (predicted): Atomic weight = $72$, Density = $5.5\,\text{g/cm}^3$, Melting point = High, Oxide = $\text{EO}_2$, Chloride = $\text{ECl}_4$.

    • Germanium (found): Atomic weight = $72.6$, Density = $5.36\,\text{g/cm}^3$, Melting point = $1231\,\text{K}$, Oxide = $\text{GeO}_2$, Chloride = $\text{GeCl}_4$.

  • Mendeleev's Periodic Table published in 1905:       

    

Modern Periodic Law and Structure of the Periodic Table

• Discovery of Atomic Number:

  • Henry Moseley (1913):

  • Studied characteristic X-ray spectra of elements.

  • Plotted $\sqrt{\nu}$ (where $\nu$ is X-ray frequency) against atomic number ($Z$) and obtained a straight line, whereas plot against atomic mass did not.

  • Established that atomic number ($Z$) 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."

  • Significance: Atomic number ($Z$) equals nuclear charge (number of protons) or total electrons in a neutral atom. The periodic variation in properties follows from the periodic variation in electronic configurations.

  • Natural Elements: Explains analogies among 94 naturally occurring elements (e.g., Neptunium and Plutonium are found in pitchblende alongside Actinium and Protactinium).

• Architecture of the Long Form Periodic Table:   

  

  • Periods (Horizontal Rows):

  • 7 periods, where period number equals the highest principal quantum number ($n$) of the elements in that row.

  • Period 1 ($n=1$): 2 elements ($\text{H}$, $\text{He}$).

  • Period 2 ($n=2$): 8 elements ($\text{Li}$ to $\text{Ne}$).

  • Period 3 ($n=3$): 8 elements ($\text{Na}$ to $\text{Ar}$).

  • Period 4 ($n=4$): 18 elements ($\text{K}$ to $\text{Kr}$).

  • Period 5 ($n=5$): 18 elements ($\text{Rb}$ to $\text{Xe}$).

  • Period 6 ($n=6$): 32 elements ($\text{Cs}$ to $\text{Rn}$).

  • Period 7 ($n=7$): 32 elements ($\text{Fr}$ to $\text{Og}$).

  • Groups (Vertical Columns):

  • 18 vertical columns numbered 1 to 18 according to 1984 IUPAC recommendations (replacing older IA–VIIA, VIII, IB–VIIB, 0 notation).

  • Elements in a group share identical outer valence shell electronic configurations and similar chemical behaviors.

  • Bottom Separate Panels:

  • Lanthanoids (14 elements: $\text{Ce}, Z=58$ to $\text{Lu}, Z=71$).

  • Actinoids (14 elements: $\text{Th}, Z=90$ to $\text{Lr}, Z=103$).

  • Reconfigured by Glenn T. Seaborg (discovered Plutonium in 1940 and transuranium elements 94 to 102; Nobel Prize in 1951; Element 106 named Seaborgium, $\text{Sg}$).

Nomenclature of Elements with Atomic Numbers $Z > 100$

• Necessity for Systematic Naming:

  • Superheavy elements are highly unstable and produced in trace quantities (a few atoms).

  • Prevents discovery priority disputes (e.g., Element 104 was named Rutherfordium by Americans and Kurchatovium by Soviets).

• IUPAC Rules for Temporary Names:

  • Derived directly from atomic number digits using numerical roots:

  • $0 = \text{nil (n)}$

  • $1 = \text{un (u)}$

  • $2 = \text{bi (b)}$

  • $3 = \text{tri (t)}$

  • $4 = \text{quad (q)}$

  • $5 = \text{pent (p)}$

  • $6 = \text{hex (h)}$

  • $7 = \text{sept (s)}$

  • $8 = \text{oct (o)}$

  • $9 = \text{enn (e)}$

  • Roots are written sequentially and appended with "ium".

• Systematic and Official Names ($Z = 101$ to $118$):

  • $101$: Unnilunium ($\text{Unu}$) $\rightarrow$ Mendelevium ($\text{Md}$)

  • $102$: Unnilbium ($\text{Unb}$) $\rightarrow$ Nobelium ($\text{No}$)

  • $103$: Unniltrium ($\text{Unt}$) $\rightarrow$ Lawrencium ($\text{Lr}$)

  • $104$: Unnilquadium ($\text{Unq}$) $\rightarrow$ Rutherfordium ($\text{Rf}$)

  • $105$: Unnilpentium ($\text{Unp}$) $\rightarrow$ Dubnium ($\text{Db}$)

  • $106$: Unnilhexium ($\text{Unh}$) $\rightarrow$ Seaborgium ($\text{Sg}$)

  • $107$: Unnilseptium ($\text{Uns}$) $\rightarrow$ Bohrium ($\text{Bh}$)

  • $108$: Unniloctium ($\text{Uno}$) $\rightarrow$ Hassium ($\text{Hs}$)

  • $109$: Unnilennium ($\text{Une}$) $\rightarrow$ Meitnerium ($\text{Mt}$)

  • $110$: Ununnillium ($\text{Uun}$) $\rightarrow$ Darmstadtium ($\text{Ds}$)

  • $111$: Unununnium ($\text{Uuu}$) $\rightarrow$ Roentgenium ($\text{Rg}$)

  • $112$: Ununbium ($\text{Uub}$) $\rightarrow$ Copernicium ($\text{Cn}$)

  • $113$: Ununtrium ($\text{Uut}$) $\rightarrow$ Nihonium ($\text{Nh}$)

  • $114$: Ununquadium ($\text{Uuq}$) $\rightarrow$ Flerovium ($\text{Fl}$)

  • $115$: Ununpentium ($\text{Uup}$) $\rightarrow$ Moscovium ($\text{Mc}$)

  • $116$: Ununhexium ($\text{Uuh}$) $\rightarrow$ Livermorium ($\text{Lv}$)

  • $117$: Ununseptium ($\text{Uus}$) $\rightarrow$ Tennessine ($\text{Ts}$)

  • $118$: Ununoctium ($\text{Uuo}$) $\rightarrow$ Oganesson ($\text{Og}$)

• Worked Problem:

  • Problem: Determine the IUPAC name and symbol for the element with atomic number 120.

  • Solution: Digits 1, 2, 0 correspond to roots un, bi, nil. Systematic name = unbinilium, Symbol = Ubn.

Electronic Configuration and Division into Blocks

• Orbital Filling in Periods:

  • Period 1 ($n=1$): Fills $1s$ orbital. 2 elements: $\text{H}\, (1s^1)$, $\text{He}\, (1s^2)$.

  • Period 2 ($n=2$): Fills $2s$ and $2p$ orbitals. 8 elements: $\text{Li}\, (1s^2 2s^1)$ to $\text{Ne}\, (1s^2 2s^2 2p^6)$.

  • Period 3 ($n=3$): Fills $3s$ and $3p$ orbitals. 8 elements: $\text{Na}\, ([\text{Ne}]3s^1)$ to $\text{Ar}\, ([\text{Ne}]3s^2 3p^6)$.

  • Period 4 ($n=4$): Fills $4s$, $3d$, and $4p$ orbitals in energy order $4s < 3d < 4p$. 18 elements: Contains $3d$ transition series from $\text{Sc}\, (Z=21, 3d^1 4s^2)$ to $\text{Zn}\, (Z=30, 3d^{10} 4s^2)$. Ends at $\text{Kr}$.

  • Period 5 ($n=5$): Fills $5s$, $4d$, and $5p$ orbitals in energy order $5s < 4d < 5p$ (9 available orbitals accommodating 18 electrons). Contains $4d$ transition series from $\text{Y}\, (Z=39)$ to $\text{Cd}\, (Z=48)$. Ends at $\text{Xe}$.

  • Period 6 ($n=6$): Fills $6s$, $4f$, $5d$, and $6p$ orbitals. 32 elements: Includes $4f$ inner transition series (Lanthanoids: $\text{Ce}, Z=58$ to $\text{Lu}, Z=71$).

  • Period 7 ($n=7$): Fills $7s$, $5f$, $6d$, and $7p$ orbitals. 32 elements: Includes $5f$ inner transition series (Actinoids: $\text{Th}, Z=90$ to $\text{Lr}, Z=103$).

• Classification into Four Blocks:

  • s-Block Elements:

  • Groups 1 (alkali metals, $ns^1$) and 2 (alkaline earth metals, $ns^2$).

  • Group 1 Configurations:

    • Lithium ($\text{Li}, Z=3$): $1s^2 2s^1$ or $[\text{He}]2s^1$

    • Sodium ($\text{Na}, Z=11$): $1s^2 2s^2 2p^6 3s^1$ or $[\text{Ne}]3s^1$

    • Potassium ($\text{K}, Z=19$): $1s^2 2s^2 2p^6 3s^2 3p^6 4s^1$ or $[\text{Ar}]4s^1$

    • Rubidium ($\text{Rb}, Z=37$): $1s^2 2s^2 2p^6 3s^2 3p^6 3d^{10} 4s^2 4p^6 5s^1$ or $[\text{Kr}]5s^1$

    • Caesium ($\text{Cs}, Z=55$): $1s^2 2s^2 2p^6 3s^2 3p^6 3d^{10} 4s^2 4p^6 4d^{10} 5s^2 5p^6 6s^1$ or $[\text{Xe}]6s^1$

    • Francium ($\text{Fr}, Z=87$): $[\text{Rn}]7s^1$

  • Characteristics: Reactive metals with low ionization enthalpies; readily form $1+$ or $2+$ ions; reactivity increases down the group; compounds are predominantly ionic (except $\text{Li}$ and $\text{Be}$).

  • p-Block Elements:

  • Groups 13 to 18. Outer configuration ranges from $ns^2 np^1$ to ns^2 np^6$.\n * Together with s-block elements, called Representative Elements or Main Group Elements.\n * Group 18 (Noble Gases): Closed valence shell ns^2 np^6; extremely unreactive.\n * Group 17 (Halogens): Outer configuration ns^2 np^5; highly negative electron gain enthalpy.\n * Group 16 (Chalcogens): Outer configuration ns^2 np^4; highly negative electron gain enthalpy.\n * Non-metallic character increases across a period; metallic character increases down a group.\n * **d-Block Elements (Transition Elements)**:\n * Groups 3 to 12. Characterized by filling of (n-1)d orbitals.\n * Outer electronic configuration: (n-1)d^{1-10} ns^{0-2}$(Exception: Palladium,$\text{Pd}$, is$4d^{10} 5s^0).\n * Characteristics: All are metals; form colored ions; display variable oxidation states, paramagnetism, and catalytic activity.\n * Zinc (\text{Zn}$), Cadmium ($\text{Cd}$), and Mercury ($\text{Hg}$) with$(n-1)d^{10} ns^2 configurations do not show typical transition properties.\n * **f-Block Elements (Inner-Transition Elements)**:\n * Two rows at bottom: Lanthanoids (\text{Ce}, Z=58$to$\text{Lu}, Z=71$) and Actinoids ($\text{Th}, Z=90$to$\text{Lr}, Z=103).\n * Outer electronic configuration: (n-2)f^{1-14} (n-1)d^{0-1} ns^2$.

  • All are metals with uniform properties within each series. Actinoids are radioactive; elements after Uranium ($Z > 92$) are called Transuranium Elements.

• Exceptions in Block Assignment:

  • Helium ($\text{He}$): Outer configuration $1s^2$ (s-block strictly), but placed in p-block (Group 18) due to filled valence shell and noble gas behavior.

  • Hydrogen ($\text{H}$): Outer configuration $1s^1$; can lose 1 electron like alkali metals or gain 1 electron like halogens; placed separately at the top.

Classification into Metals, Non-metals, and Metalloids

• Metals:

  • Make up more than 78% of all known elements; positioned on the left and center of the table.

  • Physical state: Solids at room temperature (Exception: Mercury, $\text{Hg}$, is liquid; Gallium, $\text{Ga}$, MP = $303\,\text{K}$, and Caesium, $\text{Cs}$, MP = $302\,\text{K}$, have low melting points).

  • Characteristics: High melting and boiling points, high thermal and electrical conductivity, malleable, and ductile.

• Non-Metals:

  • Positioned at the top right side of the periodic table.

  • Physical state: Solids or gases at room temperature (Exception: Bromine, $\text{Br}$, is liquid).

  • Characteristics: Low melting and boiling points (Exceptions: Boron, Carbon), poor conductors of heat and electricity, solid non-metals are brittle.

• Metalloids (Semi-Metals):

  • Elements bordering the thick diagonal zig-zag line separating metals from non-metals.

  • Examples: Silicon ($\text{Si}$), Germanium ($\text{Ge}$), Arsenic ($\text{As}$), Antimony ($\text{Sb}$), Tellurium ($\text{Te}$).

  • Display properties characteristic of both metals and non-metals.

• Worked Problem:

  • Problem: Arrange $\text{Si}$, $\text{Be}$, $\text{Mg}$, $\text{Na}$, $\text{P}$ in order of increasing metallic character.

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

  • Solution: $\text{P} < \text{Si} < \text{Be} < \text{Mg} < \text{Na}$.

Periodic Trends in Physical Properties

• Atomic Radius:

  • Definition & Determination:

  • Atomic size cannot be measured directly for a single atom due to the absence of a sharp boundary (atomic radius $\sim 1.2\,\text{\AA} = 1.2 \times 10^{-10}\,\text{m}$).

  • Covalent Radius: Half the distance between two single-bonded atoms of the same non-metallic element. Example: Bond length in $\text{Cl}_2 = 198\,\text{pm} \rightarrow$ Covalent radius of $\text{Cl} = 99\,\text{pm}$.

  • Metallic Radius: Half the internuclear distance between adjacent metal cores in a crystal lattice. Example: Distance between copper atoms in solid $\text{Cu} = 256\,\text{pm} \rightarrow$ Metallic radius of $\text{Cu} = 128\,\text{pm}$.

  • Trends:

  • Across a Period: Atomic radius decreases left to right. Outer electrons enter the same principal quantum shell while effective nuclear charge ($Z_{\text{eff}}$ is Effective nuclear charge the net positive charge experienced by an electron in a multi-electron atom accounting for both the attractive pull of the nucleus and the repulsive shielding from inner-shell electrons,)increases, attracting electrons closer to the nucleus.

    • Period 2 Radii ($\text{pm}$): $\text{Li}\, (152)$, $\text{Be}\, (111)$, $\text{B}\, (88)$, $\text{C}\, (77)$, $\text{N}\, (74)$, $\text{O}\, (66)$, $\text{F}\, (64)$

    • Period 3 Radii ($\text{pm}$): $\text{Na}\, (186)$, $\text{Mg}\, (160)$, $\text{Al}\, (143)$, $\text{Si}\, (117)$, $\text{P}\, (110)$, $\text{S}\, (104)$, $\text{Cl}\, (99)$

  • Down a Group: Atomic radius increases down a column. Principal quantum number ($n$) increases and inner energy levels shield valence electrons from nuclear attraction.

    • Group 1 Radii ($\text{pm}$): $\text{Li}\, (152)$, $\text{Na}\, (186)$, $\text{K}\, (231)$, $\text{Rb}\, (244)$, $\text{Cs}\, (262)$

    • Group 17 Radii ($\text{pm}$): $\text{F}\, (64)$, $\text{Cl}\, (99)$, $\text{Br}\, (114)$, $\text{I}\, (133)$, $\text{At}\, (140)$

  • Noble Gases: Monoatomic noble gases display non-bonded van der Waals radii, which are significantly larger than covalent radii.

• Ionic Radius:

  • Cations: Smaller than parent atom due to loss of electron(s), leaving same nuclear charge acting on fewer electrons. Example: $\text{Na}\, (186\,\text{pm})$ vs. $\text{Na}^+\, (95\,\text{pm})$.

  • Anions: Larger than parent atom due to gain of electron(s), increasing electron-electron repulsion and decreasing $Z_{\text{eff}}$. Example: $\text{F}\, (64\,\text{pm})$ vs. $\text{F}^-\, (136\,\text{pm})$.

  • Isoelectronic Species:

  • Species with the same number of electrons (e.g., $\text{O}^{2-}$, $\text{F}^-$, $\text{Na}^+$, $\text{Mg}^{2+}$ all contain 10 electrons).

  • Radius decreases with increasing positive nuclear charge:       $\text{O}^{2-} > \text{F}^- > \text{Na}^+ > \text{Mg}^{2+}$

  • Worked Problem:

  • Problem: Find largest and smallest among $\text{Mg}$, $\text{Mg}^{2+}$, $\text{Al}$, $\text{Al}^{3+}$.

  • Solution: Largest = $\text{Mg}$; Smallest = $\text{Al}^{3+}$.

• Ionization Enthalpy ($\Delta_i H$):

  • Definition: Energy required to remove an electron from an isolated gaseous atom $\text{X}$ in its ground state:     $\text{X}(g) \rightarrow \text{X}^+(g) + e^-$     Expressed in $\text{kJ\,mol}^{-1}$. Values are always positive ($\Delta_i H > 0$).

  • Successive Ionization Enthalpies:     $\Delta_i H_1 < \Delta_i H_2 < \Delta_i H_3$     Removing an electron from a positively charged ion requires progressively higher energy.

  • Periodic Trends:

  • Across a Period: Ionization enthalpy increases due to higher $Z_{\text{eff}}$ and smaller atomic radius.

  • Down a Group: Ionization enthalpy decreases due to larger atomic radius and inner shell shielding.

  • Anomalies in Period 2:

  • Boron vs. Beryllium: $\Delta_i H(\text{B}) < \Delta_i H(\text{Be})$. Boron's valence electron resides in a $2p$ orbital (higher energy, more shielded), whereas Beryllium's resides in a $2s$ orbital.

  • Oxygen vs. Nitrogen: $\Delta_i H(\text{O}) < \Delta_i H(\text{N})$. Nitrogen has three singly occupied $2p$ orbitals ($2p^3$), whereas Oxygen has two paired $2p$ electrons in one orbital ($2p^4$), introducing electron repulsion that facilitates electron removal.

  • Worked Problem:

  • Problem: Given $\Delta_i H_1$ for $\text{Na}\, (496\,\text{kJ\,mol}^{-1})$, $\text{Mg}\, (737\,\text{kJ\,mol}^{-1})$, $\text{Si}\, (786\,\text{kJ\,mol}^{-1})$, predict if $\text{Al}$ is closer to $575$ or $760\,\text{kJ\,mol}^{-1}$.

  • Solution: Closer to $575\,\text{kJ\,mol}^{-1}$ due to effective shielding of the $3p$ electron by $3s$ electrons.

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

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

  • Exothermic vs. Endothermic:

  • Halogens (Group 17) have large negative values because adding one electron yields a stable noble gas configuration.

  • Noble gases (Group 18) have positive values because the electron enters the next higher principal quantum level.

  • Periodic Trends:

  • Across a Period: Becomes more negative left to right as $Z_{\text{eff}}$ increases.

  • Down a Group: Becomes less negative down a group as atomic size increases.

  • Anomalies for Oxygen and Fluorine: $\Delta_{\text{eg}} H$ for Oxygen and Fluorine is less negative than for Sulfur and Chlorine because the added electron enters a small $n=2$ quantum level, experiencing strong inter-electronic repulsion.

  • Values for Selected Elements ($\text{kJ\,mol}^{-1}$):

  • Group 1: $\text{H}\, (-73)$, $\text{Li}\, (-60)$, $\text{Na}\, (-53)$, $\text{K}\, (-48)$, $\text{Rb}\, (-47)$, $\text{Cs}\, (-46)$

  • Group 16: $\text{O}\, (-141)$, $\text{S}\, (-200)$, $\text{Se}\, (-195)$, $\text{Te}\, (-190)$, $\text{Po}\, (-174)$

  • Group 17: $\text{F}\, (-328)$, $\text{Cl}\, (-349)$, $\text{Br}\, (-325)$, $\text{I}\, (-295)$, $\text{At}\, (-270)$

  • Group 18: $\text{He}\, (+48)$, $\text{Ne}\, (+116)$, $\text{Ar}\, (+96)$, $\text{Kr}\, (+96)$, $\text{Xe}\, (+77)$, $\text{Rn}\, (+68)$

  • Worked Problem:

  • Problem: Compare $\text{P}$, $\text{S}$, $\text{Cl}$, $\text{F}$ for most and least negative \Delta_{\text{eg}} H$.\n * Solution: Most negative = Chlorine (\text{Cl}$); Least negative = Phosphorus ($\text{P}).\n\n* Electronegativity:\n * Definition: Qualitative measure of the ability of an atom in a compound to attract shared electron pairs to itself.\n * Pauling Scale: Formulated by Linus Pauling; arbitrarily assigned Fluorine = 4.0$.

  • Trends:

  • Increases across a period from left to right (e.g., Period 2: $\text{Li}\, 1.0$, $\text{Be}\, 1.5$, $\text{B}\, 2.0$, $\text{C}\, 2.5$, $\text{N}\, 3.0$, $\text{O}\, 3.5$, $\text{F}\, 4.0$).

  • Decreases down a group (e.g., Group 17: $\text{F}\, 4.0$, $\text{Cl}\, 3.0$, $\text{Br}\, 2.8$, $\text{I}\, 2.5$, $\text{At}\, 2.2$).

  • Directly proportional to non-metallic character; inversely proportional to metallic character.

Periodic Trends in Chemical Properties

• Valence / Oxidation States:

  • Valence of representative elements equals valence electrons (Groups 1–4) or $8 - \text{valence electrons}$ (Groups 15–18).

  • Hydrides and Oxides Across Period:

  • Group 1: Hydride $\text{LiH}$, Oxide $\text{Li}_2\text{O}$

  • Group 2: Hydride $\text{CaH}_2$, Oxide $\text{MgO}$

  • Group 13: Hydride $\text{B}_2\text{H}_6$, Oxide $\text{B}_2\text{O}_3$

  • Group 14: Hydride $\text{CH}_4$, Oxide $\text{CO}_2$

  • Group 15: Hydride $\text{NH}_3$, Oxide $\text{N}_2\text{O}_5$

  • Group 16: Hydride $\text{H}_2\text{O}$, Oxide $\text{SO}_3$

  • Group 17: Hydride $\text{HF}$, Oxide $\text{Cl}_2\text{O}_7$

  • Oxidation State Examples:

  • In $\text{OF}_2$: Fluorine is $-1$; Oxygen exhibits $+2$.

  • In $\text{Na}_2\text{O}$: Oxygen is $-2$; Sodium exhibits $+1$.

  • Worked Problem:

  • Problem: Predict formulas for compounds formed by (a) Silicon and Bromine, (b) Aluminium and Sulphur.

  • Solution: (a) $\text{SiBr}_4$ (valences 4 and 1); (b) $\text{Al}_2\text{S}_3$ (valences 3 and 2).

• Anomalous Properties of Second Period Elements:

  • First elements ($\text{Li}$, $\text{Be}$, $\text{B}$, $\text{C}$, $\text{N}$, $\text{O}$, $\text{F}$) differ significantly from heavier members in their respective groups.

  • Examples: $\text{Li}$ and $\text{Be}$ form covalent compounds, whereas heavier alkali/alkaline earth metals form ionic compounds.

  • Diagonal Relationship: Similarity in properties between $\text{Li}$ and $\text{Mg}$, and between $\text{Be}$ and $\text{Al}$.

  • Causes:

  • Small atomic and ionic radius.

  • High charge/radius ratio.

  • High electronegativity.

  • Availability of only four valence orbitals ($2s$, three $2p$); maximum covalency is 4 (e.g., $\text{BF}_4^-$ vs. $\text{AlF}_6^{3-}$).

  • Ability to form $p\pi-p\pi$ multiple bonds ($\text{C=C}$, $\text{C}\equiv\text{C}$, $\text{N}\equiv\text{N}$, $\text{C=O}$).

  • Worked Problem:

  • Problem: Compare oxidation state and covalency of $\text{Al}$ in $[\text{AlCl}(\text{H}_2\text{O})_5]^{2+}$.

  • Solution: Oxidation state = $+3$; Covalency = 6$.\n\n* Periodic Trends and Chemical Reactivity:\n * Reactivity Extremes:\n * Group 1 (extreme left): High reactivity by losing an electron (low ionization enthalpy), forming strongly basic oxides (e.g., \text{Na}_2\text{O}).\n * Group 17 (extreme right): High reactivity by gaining an electron (highly negative electron gain enthalpy), forming strongly acidic oxides (e.g., \text{Cl}_2\text{O}_7).\n * Center elements: Form amphoteric oxides (e.g., \text{Al}_2\text{O}_3$,$\text{As}_2\text{O}_3$) or neutral oxides (e.g.,$\text{CO}$,$\text{NO}$,$\text{N}_2\text{O}).\n * Reaction with Water:\n    \text{Na}_2\text{O} + \text{H}_2\text{O} \rightarrow 2\text{NaOH} \quad (\text{Basic})\n    \text{Cl}_2\text{O}_7 + \text{H}_2\text{O} \rightarrow 2\text{HClO}_4 \quad (\text{Acidic})$$

  • Reactivity Trends in Groups:

  • Group 1 metals: Reactivity increases down the group as ionization enthalpy decreases.

  • Group 17 non-metals: Reactivity decreases down the group as electron gain enthalpy becomes less negative.