Comprehensive Study of Elements in Groups 16, 17, and 18

Group 16, 17, and 18 Element Introduction and Occurrence

The p-block of the periodic table consists of groups 13 to 18, where the differentiating or last filling electron enters the p-orbital of the outermost shell. Because a p-subshell can accommodate a maximum of six electrons, six groups are present. Group 16 consists of oxygen ($O$), sulfur ($S$), selenium ($Se$), tellurium ($Te$), and polonium ($Po$). This group is known as the oxygen family or chalcogens, a term derived from their role as ore-forming elements, as many metal ores exist as oxides or sulfides. Oxygen is the most abundant element on Earth, constituting $20.95\,\%$ of the atmosphere by volume and $46.6\,\%$ of the earth's crust by mass. Sulfur is less abundant at $0.034\,\%$ by mass and is found in sulfates such as gypsum ($CaSO_4 \cdot 2H_2O$), epsom salt ($MgSO_4 \cdot 7H_2O$), and baryte ($BaSO_4$), or sulfides including galena ($PbS$), zinc blende ($ZnS$), and copper pyrites ($CuFeS_2$). Selenium and tellurium occur as metal selenides and tellurides in sulfide ores, while polonium is a radioactive decay product of thorium and uranium.

Group 17 includes fluorine ($F$), chlorine ($Cl$), bromine ($Br$), iodine ($I$), and astatine ($At$). These are collectively called halogens, meaning salt-producers. Due to their high electronegativity, they are never found in a free state but exist as compounds. Fluorine is found in insoluble fluorides like fluorspar ($CaF_2$), cryolite ($Na_3AlF_6$), and fluorapatite ($3Ca_3(PO_4)_2 \cdot CaF_2$). Sea water contains chlorides, bromides, and iodides of sodium, potassium, magnesium, and calcium, dominated by sodium chloride which is $2.5\,\%$ by mass. Marine life accumulates iodine; sea weed can contain up to $0.5\,\%$ iodine and chile saltpetre contains up to $0.2\,\%$ sodium iodate ($NaIO_3$). Astatine is radioactive with a short half-life of $8.1\,hours$.

Group 18 elements include helium ($He$), neon ($Ne$), argon ($Ar$), krypton ($Kr$), xenon ($Xe$), and radon ($Rn$), known as noble gases. Except for radon, all occur in the atmosphere, with dry air containing approximately $1\,\%$ noble gases by volume, largely argon. Natural gas is the primary commercial source of helium. Helium and neon are also found in radioactive minerals like pitchblende, monazite, and cleveite. Xenon and radon are the group's rarest elements, with radon produced from the decay of ${}^{226}Ra$.

Electronic Configuration and Atomic Properties

The general valence shell electronic configurations are $ns^2np^4$ for Group 16, $ns^2np^5$ for Group 17, and $ns^2np^6$ for Group 18. Both Group 16 and 17 elements have two and one fewer electrons, respectively, than the stable configuration of the nearest noble gas. The specific condensed configurations for Group 16 range from $[He]2s^22p^4$ for oxygen to $[Xe]4f^{14}5d^{10}6s^26p^4$ for polonium. In Group 17, they range from $[He]2s^22p^5$ for fluorine to $[Xe]4f^{14}5d^{10}6s^26p^5$ for astatine. For Group 18, configurations range from $1s^2$ for helium to $[Xe]4f^{14}5d^{10}6s^26p^6$ for radon.

Atomic and ionic radii increase down the group due to the addition of quantum shells but decrease across a period from left to right as the effective nuclear charge ($Z_{eff}$) increases. Halogens have the smallest atomic radii in their respective periods. Ionization enthalpy follows the inverse trend: it decreases down a group due to increased atomic size and increases across a period. Notably, Group 16 elements have lower first ionization enthalpies than Group 15 elements in the same period because Group 15 has the extra stability of a half-filled p-subshell ($ns^2np^3$). For instance, the first ionization enthalpy for sulfur is $1000\,kJ\,mol^{-1}$ while chlorine is higher at $1256\,kJ\,mol^{-1}$ due to increased nuclear charge and smaller size.

Electronegativity decreases down the group. Fluorine is the most electronegative element in the entire periodic table, followed by oxygen. Electron gain enthalpy generally becomes less negative down the group, but there are exceptions. Oxygen has a less negative electron gain enthalpy ($-141\,kJ\,mol^{-1}$) than sulfur ($-200\,kJ\,mol^{-1}$) because its small size causes significant inter-electronic repulsion. Similarly, fluorine's electron gain enthalpy ($-333\,kJ\,mol^{-1}$) is less negative than chlorine's ($-349\,kJ\,mol^{-1}$) because the incoming electron experiences more repulsion in fluorine's compact $2p$ orbitals. Noble gases have large positive electron gain enthalpies ($48$ to $116\,kJ\,mol^{-1}$) because they have no tendency to accept electrons into their stable shells.

Physical and Chemical Characteristics

Oxygen is a gas, whereas other Group 16 elements are solids. Oxygen and sulfur are nonmetals, selenium and tellurium are metalloids, and polonium is a radioactive metal with a half-life of $13.8\,days$. Melting and boiling points increase with atomic number. There is a significant difference between oxygen and sulfur because oxygen exists as diatomic $O_2$ with weak van der Waals forces, while sulfur exists as polyatomic $S_8$ with much stronger forces. All Group 16 members exhibit allotropy. For Group 17, fluorine and chlorine are gases, bromine is a liquid, and iodine is a solid. Their colors vary: fluorine is yellow, chlorine is greenish-yellow, bromine is red, and iodine is violet. Bond dissociation enthalpies for halogens follow the order $Cl-Cl > Br-Br > F-F > I-I$. The lower value for $F_2$ ($158.8\,kJ\,mol^{-1}$) compared to $Cl_2$ ($242.6\,kJ\,mol^{-1}$) is due to the strong repulsion between lone pairs on the tiny fluorine atoms.

Noble gases are monoatomic and sparingly soluble in water, possessing very low melting and boiling points because they only interact via weak van der Waals forces. Helium has the lowest boiling point of any known substance at $4.2\,K$. All Group 18 elements except argon have extremely low atmospheric abundances. Xenon can exhibit higher oxidation states ($+2$, $+4$, $+6$) because its large size and lower ionization enthalpy allow it to promote paired electrons into empty d-orbitals, forming compounds like $XeF_2$, $XeF_4$, and $XeO_3$.

Anomalous Behavior and Reactivity

Oxygen and fluorine exhibit properties different from their respective group members due to their small atomic size, high electronegativity, and the absence of d-orbitals in their valence shells. Oxygen is diatomic and paramagnetic, whereas others are polyatomic and diamagnetic. Oxygen is limited to oxidation states of $-2$, $-1$, and $+2$ (in $OF_2$) because it cannot expand its octet. Oxygen's hydride ($H_2O$) is a liquid due to hydrogen bonding, whereas others are gases. Fluorine is similarly unique; it is the most electronegative, has the highest electrode potential, and its reactions are often highly exothermic. Hydrogen fluoride ($HF$) is a liquid (boiling point $293\,K$) due to strong hydrogen bonding, while other hydrogen halides are gases.

Chemical reactivity towards hydrogen results in hydrides like $H_2E$ for Group 16 and $HX$ for Group 17. The acidity of these hydrides increases down the group ($H_2O < H_2S < H_2Se < H_2Te$ and $HF < HCl < HBr < HI$) because bond dissociation enthalpy decreases as the central atom grows larger. Thermal stability follows the opposite trend, decreasing down the group. For Group 16, all hydrides except water act as reducing agents. Reactivity with oxygen leads to various oxides. Group 16 forms dioxide ($EO_2$) and trioxide ($EO_3$) types. Ozone ($O_3$) and $SO_2$ are gases, while $SeO_2$ is a solid. Halogens form multiple oxides, though many are unstable; for example, iodine oxides like $I_2O_5$ are used for estimating carbon monoxide.

Group 16 elements react with halogens to form hexahalides ($EX_6$), tetrahalides ($EX_4$), and dihalides ($EX_2$). Hexahalides like $SF_6$ have octahedral structures and $sp^3d^2$ hybridization; $SF_6$ is remarkably stable due to steric factors. Tetrahalides like $SF_4$ have trigonal bipyramidal geometry with one equatorial lone pair. Halogens react with each other to form interhalogen compounds of types $XX'$, $XX'_3$, $XX'_5$, and $XX'_7$. Reactivity with metals produces metal oxides, sulfides, and halides. Ionic character in metal halides decreases in the order $MF > MCl > MBr > MI$. Metals in higher oxidation states tend to form more covalent halides, such as $SnCl_4$ being more covalent than $SnCl_2$.

Allotropy, Oxoacids, and Oxygen Compounds

Sulfur has several allotropes, the most significant being rhombic ($\alpha$-sulfur) and monoclinic ($\beta$-sulfur). Rhombic sulfur is pale yellow, possesses orthorhombic crystals, and is stable below $369\,K$. Above this transition temperature, it converts to bright yellow monoclinic sulfur, which consists of needle-shaped crystals. Both consist of $S_8$ molecules with a puckered ring structure. At high temperatures ($\sim 1000\,K$), sulfur exists as $S_2$ molecules which are paramagnetic like $O_2$. Selenium has red (non-metallic) and grey (metallic/photoconductor) forms, while polonium has $\alpha$ and $\beta$ metallic forms.

Sulfur forms several oxoacids, including sulfurous acid ($H_2SO_3$), sulfuric acid ($H_2SO_4$), pyrosulfuric acid ($H_2S_2O_7$), and peroxy acids. Halogens also form oxoacids where acid strength increases with the oxidation state of the halogen (e.g., $HClO < HClO_2 < HClO_3 < HClO_4$). Dioxygen ($O_2$) is prepared in the lab by heating salts like $KClO_3$ with a $MnO_2$ catalyst or by the thermal decomposition of metal oxides like $HgO$. Industrially, it is produced via the fractional distillation of liquefied air or the electrolysis of water. Dioxygen is a colorless, paramagnetic gas essential for respiration, steel manufacture, and oxyacetylene sparks for welding. Binary oxygen compounds are classified as acidic (forming acids with water like $SO_2$), basic (forming bases like $Na_2O$), amphoteric (reacting with both acids and bases like $Al_2O_3$), or neutral (like $CO$ or $NO$). Ozone ($O_3$) is an allotrope of oxygen formed by passing a silent electric discharge through oxygen in an ozoniser, an endothermic process where $3O_2 \rightarrow 2O_3$ with $\Delta H = +142\,kJ\,mol^{-1}$.

Questions & Discussion

Questions arise regarding the variation of electronic configurations and the identification of the first elements in groups 16, 17, and 18. The first elements are oxygen, fluorine, and helium, respectively. In the p-block, the configuration changes as electrons fill the p-subshell from $p^1$ to $p^6$. Discussions also cover why water must be changed in fish tanks (to maintain dissolved oxygen levels) and why oxygen has a less negative electron gain enthalpy than sulfur (its small size creates high electron density/repulsion). The paramagnetism of dioxygen, despite having an even number of electrons, is explained via Molecular Orbital Theory, which shows two unpaired electrons in the antibonding $\pi^*$ orbitals. Additionally, the large difference in melting and boiling points between oxygen and sulfur is attributed to their molecular atomicity ($O_2$ vs $S_8$) and the resulting difference in the magnitude of van der Waals forces.