Group 5A Elements PDF

Group 5A Elements  Group 15 elements - sometimes called “Pnictogens” ▪ Much of the important chemistry of the group 15 elements can be understood on the basis of their electronic structure. ▪ Since the elements have a [core]ns2 np3 electron configuration, neutral group 15 co...

Group 5A Elements  Group 15 elements - sometimes called “Pnictogens” ▪ Much of the important chemistry of the group 15 elements can be understood on the basis of their electronic structure. ▪ Since the elements have a [core]ns2 np3 electron configuration, neutral group 15 compounds can form up to five bonds. ▪ This provides for two common oxidation states (+3 and +5) electrons (with a complete octet) around the group 15 atom so such compounds are called “electron-rich”. ▪ The group 5A elements have the outer-shell electron configuration ns2np3, with n ranging from 2 to 6. ▪ More commonly, the group 5A element acquires an octet of electrons via covalent bonding and oxidation numbers ranging from –3 to +5. Occurrence Nitrogen is the real constituent of the air and records for 78% of it by volume. It is the primary member of this group and happens in a free state as a diatomic gas, N2. Phosphorus is a fundamental constituent of animal and plant matter. Phosphate groups are constituents of nucleic acids, that is, DNA and RNA. Around 60% of bones and teeth are made out of phosphates. Phosphoproteins are available in egg yolk, milk, and bone marrow. Arsenic, antimony, and bismuth, mostly happen as sulfides. For example, stibnite, arsenopyrite, and bismuth glance. Atomic properties 1) Atomic Radii Moving down the group, the ionic radii, and atomic radii increases. This is because of the expansion of another main energy level in each progressive element. 2) Ionization Enthalpy As we move down the group, the ionization enthalpy values keep on decreasing. This is because of the progressive increase in the size of the nucleus. 3) Electronegativity The electronegativity decreases gradually on moving down the group. The reason behind this is the increase in atomic radius. 4) Physical Properties Physical properties incorporate physical state, boiling and melting points, metallic character, allotropy, and density. Nitrogen is a diatomic gas, while the rest of the elements are solids in nature. Moving down a group, metallic character increases. On the other hand, the ionization enthalpy of the elements decreases due to an increase in their nuclear size. 5) Trends in Melting and Boiling Points The melting point increments from nitrogen to arsenic because of the continuous increment in nuclear size. The low melting point of nitrogen is because of its discrete diatomic particles. In spite of the fact that the nuclear size increments from arsenic to antimony, there is a decrease in their melting points. Despite the fact that antimony has a layered structure, it has a low melting point than arsenic on account of the generally free pressing of particles. The melting point of bismuth is not as much as antimony because of the packing of atoms loosely by metallic holding. The boiling point step by step increments from nitrogen to bismuth. The density of these elements increases from nitrogen to bismuth. 6) Allotropy All group fifteen elements, aside from bismuth, indicate allotropy. Nitrogen is found in two allotropic structures, alpha nitrogen and beta nitrogen. Phosphorus exists in numerous allotropic structures. Of these, the two critical allotropic structures are red phosphorus and white phosphorus Four allotropes of antimony are known: a stable metallic form (grey), and three metastable forms (explosive (white), black, and non-metallic (yellow)). Elemental Phosphorus The general properties of group fifteen elements Anomalous Properties ▪ The exceptional properties of nitrogen are credited to its small nuclear size, high ionization enthalpy or high electro negativity, the non–availability of d-orbitals and the possibility to shape various bonds. No one but nitrogen can shape nitride particles by picking up electrons because of its high electro negativity and small size. ▪ Nitrogen, being smaller in size, can successfully shape p – p bonds with different molecules of different elements with a small size and high electronegativity ▪ Dinitrogen is a diatomic particle with a triple bond between the two molecules The bond enthalpy of a triple bond is extremely high around 941.4 KJ/mole. Dinitrogen is stable under conventional conditions. ▪ The different elements in the group, for example, phosphorus, arsenic and antimony, exist as tetra atomic particles. In every one of these particles, just single bonds are available between the atoms. ▪ Bismuth in its elemental state shapes metallic bonds. ▪ The catenation inclination is less for nitrogen when contrasted with alternate elements of the group. This is on the grounds that there are higher inter- electronic repulsions amongst the lone pair of electrons present on the nitrogen atoms. ▪ The high inter- electronic aversions in dinitrogen are credited to its small bond length or little size of nitrogen particles. Nitrogen does not shape d – p bonds because of the missing d orbitals. ▪ Phosphorus can frame d – p bonds. Example: triethyl phosphate and phosphorus oxo chloride. ▪ Due to the accessibility of empty d orbitals in the rest of the elements of group 15, they frames compounds, for example, triphenyl arsine and triethyl phosphine, shape d – d bonds with transition element Chemical Reactivity Reactivity towards Hydrogen: Every one of the elements of group 15 forms EH3 type hydrides. Here E can be any element of group 15 such as nitrogen, phosphorus, arsenic, antimony or bismuth Stability: The inertness of hydrides decreases from ammonia to bismuth. This is on account of the fact that the central atom E increases in size down the group. With this increase in the central atom's size, the E – H bond gets to be distinctly weaker Reducing Character: Hydrides formed from group 15 elements are very strong reducing agents. There is an increase in the reducing character of hydrides from ammonia to bismuth because of a reduction in the quality and strength of the E – H bond down the group. Bismuth is the strongest reducing agent among every one of the hydrides of group 15 elements. Basic Nature: The hydrides of these elements are basic in nature. They go about as Lewis bases because of the accessibility of a lone pair of electrons present on the central atom. With the increasing size of the central atom, there is a decrease in the basic character as we move down the group. Boiling Point: The boiling point of hydrides decreases from ammonia to phosphine and afterward increments from phosphine to bismuth. A similar pattern is watched for their melting points Reactivity towards Oxygen: ✓Two types of oxides are formed in group 15 elements. They are E2O3 and E2O5. ✓ pπ-pπ bonding tendency with oxygen is very high in Nitrogen. This is the reason why nitrogen forms a variety of oxides. Nitrogen forms 5 stable oxides. ✓Due to the inert pair effect, bismuth is not able to form oxides in +5 oxidation. Oxides formed from elements when in higher oxidation state are more acidic than that of the lower oxidation state. Acidity The acidic strength of oxides of nitrogen increases from N2O3 to N2O5. As we move down the group, the acidic character diminishes. As such, the basic character of oxides increases on moving down the group Reactivity towards halogens: ✓On reaction with halogens, all the elements of group 15 form trihalides and pentahalides with the general formula EX3 and EX5. Example: NF3, PF3, AsF3, SbF3 and BiF3 are trihalides. ✓Phosphorus, arsenic and antimony shape pentahalides in light of the closeness of empty d orbitals in their valence shells. Nitrogen does not shape pentahalides as a result of the nonappearance of a d orbital in its valence shell. ✓Pentahalides are more covalent than the relating trihalides ✓the covalent character of halides decreases from nitrogen to bismuth. Reactivity towards metals: Each element of group 15 react with metals to frame their binary compounds demonstrating - 3 oxidation state with the general equation, M3E2. Here, M remains for metals while E remains for an element of group 15. Example: calcium phosphide, calcium nitride, and so forth. 3M + 2E → M3E 6Ca + P2 → 2Ca3P2 (calcium phosphide) 3Ca + N2 → Ca3N2 (calcium nitride) 6Zn + 4Sb → 2Zn3Sb2 (Zinc antimonide) 6Mg + 4Bi → 2Mg3Bi2 (Magnesium bismuthide) Important Compounds and Reactions Nitrogen ▪ Elemental nitrogen is an extremely stable molecule due to the triple bond. As a result, many nitrogen containing compounds decompose exothermically (and sometimes explosively) to form nitrogen gas. ▪ Nitrogen based explosives such as nitroglycerin, will rapidly decompose when ignited or exposed to a sudden impact Explosives C3H5(NO3)3(l) → 6 N2(g) + 12CO2(g) + 12H2O(g) + O2(g) + energy Note the large number of moles of gaseous products. Explosives typically involve a very large volume change, producing many moles of small gaseous molecules. Explosives Trinitrotoluene, TNT, is another nitrogen based explosive. 2C7H5(NO3)3(l) → 12 CO2(g) + 5 H2(g) + 3N2(g) + 2C(s) + energy Sodium Azide Sodium azide, NaN3(s), is used in air bags in automobiles. A small amount of sodium azide (100g) yields 56L of nitrogen gas at 25oC and 1 atm.  NaN3(s) → 2Na(l) + 3 H2(g) This reaction takes place in about 40ms. Other components are put in the air bag so that the molten sodium metal is deactivated into glassy silicates. 10 Na(l) +2KNO3(s) →K2O(s) +5Na2O(s)+ N2(g) 2 K2O(s) + SiO2(s) → K4SiO4(s) 2 Na2O(s) + SiO2(s) → Na4SiO4(s) Environmental Issues ▪ Nitrogen dioxide (NO2) and dinitrogen tetra-oxide (N2O4) are in equilibrium with each other: N2O4(g) ↔ 2 NO2(g) colorless red-brown ▪ The oxides of nitrogen are the result of high temperature combustion in jet engines and automobiles. ▪ They also react with moisture in the air to produce nitric acid and nitrous acid. ▪ This “acid rain” is a respiratory irritant, and destroys facades of buildings and statuary. Application of As, Sb, Bi, Compounds of As, Sb, and Bi with the metals of group III (Al, Ga, In, Tl) are important semiconductors Biological Aspects - Nitrogen All plant life requires nitrogen for growth and survival. Bacteria found in nodules on the roots of pea, bean, alder and clover plants convert nitrogen in the air to nitrogen compounds. Biological Aspects - Phosphorus Phosphorus is essential for life. The hydrogen phosphate ion and dihydrogen phosphate ions are involved in buffering blood. Phosphate units link the sugar esters of DNA and RNA, and also make up part of ATP, the energy storage unit in living things. Biological Aspects - Arsenic ▪ Arsenic, though generally considered toxic, is also essential to life. We only need trace amounts, and its role is still unknown. ▪ In the 19th century, before the discovery of antibiotics, arsenic was used as one of the first forms of chemotherapy to destroy the organism that causes syphilis. Summary ▪ The reactivity of group 15 elements decreases down the group, as does the stability of their catenated compounds. ▪ In group 15, nitrogen and phosphorus behave chemically like nonmetals, arsenic and antimony behave like semimetals, and bismuth behaves like a metal. ▪ Nitrogen forms compounds in nine different oxidation states. ▪ The stability of the +5 oxidation state decreases from phosphorus to bismuth because of the inert-pair effect. ▪ Due to their higher electronegativity, the lighter pnictogens form compounds in the −3 oxidation state. ▪ Because of the presence of a lone pair of electrons on the pnictogen, neutral covalent compounds of the trivalent pnictogens are Lewis bases. ▪ Nitrogen does not form stable catenated compounds because of repulsions between lone pairs of electrons on adjacent atoms, but it does form multiple bonds with other second-period atoms.  Nitrogen reacts with electropositive elements to produce solids that range from covalent to ionic in character. Reaction with electropositive metals produces ionic nitrides, reaction with less electropositive metals produces interstitial nitrides, and reaction with semimetals produces covalent nitrides. The reactivity of the pnictogens decreases with increasing atomic number. Compounds of the heavier pnictogens often have coordination numbers of 5 or higher and use dsp3 or d2sp3 hybrid orbitals for bonding. Because phosphorus and arsenic have energetically accessible d orbitals, these elements form π bonds with second-period atoms such as O and N. Phosphorus reacts with metals to produce phosphides. Metal-rich phosphides are hard, high-melting, electrically conductive solids with metallic luster, whereas phosphorus-rich phosphides, which contain catenated phosphorus units, are lower melting and less thermally stable. References Cotton F.A., G. Wilkinson, and Gauss, P.A, Basic inorganic Chemistry, Wiley & Sons., Rayner-Canham, G., Descriptive Inorganic Chemistry, 2nd edition, W.H Freeman & Co, New York, Shriver, D.F and Atkin, P.W., 2006, Inorganic Chemistry, 4th ed, WH Freeman & Co, New York Suyanta, 2013, Kimia Unsur, Gadjah Mada University Press, Yogyakarta https://chem.libretexts.org/Bookshelves/Inorganic_Chemistry/Modules_and_Websit es_(Inorganic_Chemistry)/Descriptive_Chemistry/Elements_Organized_by_Block/2_ p- Block_Elements/Group_15%3A_The_Nitrogen_Family/1Group_15%3A_General_Pro perties_and_Reactions https://www.toppr.com/guides/chemistry/the-p-block-elements/group-15-elements/ https://www.unf.edu/~michael.lufaso/chem4612/chapter15.pdf http://taladev.com/ebook/products/0-13-190443-4/ddref_chem05_eh07.pdf https://fns.uniba.sk/fileadmin/prif/chem/kag/Bakalar/vch_noga/GEN_INORG_CHEM 14b.pdf https://www.toppr.com/guides/chemistry/the-p-block-elements/group-15-elements/ The Group 6A Elements  Group 16 elements - the Chalcogens ▪ Much of the important chemistry of the group 16 elements can be understood on the basis of their electronic structure and electronegativity. ▪ Since the elements have a [core]ns2 np4 electron configuration, neutral group 16 compounds can form up to six bonds. ▪ This provides for common oxidation state from -2 to +6 electrons (with a complete octet) around the group 16 atom, so such compounds are also “electron-rich” but the high electronegativities of O and S make them good oxidizing agents.  The most common oxidation numbers for group 6A elements are +4, +6, and -2. The group 6A elements can be found in nature in both free and combined states. These elements are intimately related to life. We need oxygen all the time throughout our lives. Sulfur is responsible for some of the protein structures in all living organisms. Many industries utilize sulfur, but emission of sulfur compounds is often seen more as a problem than the natural phenomenon. The metallic properties increase as the atomic number increases. The element polonium has no stable isotopes Atomic and Physical Properties and the Periodic Trends The electron configurations for each element are given below: Oxygen : 1s2 2s2 2p4 Sulfur : 1s2 2s2p6 3s2p4 Selenium: 1s2 2s2p6 3s2p6d10 4s2p4 Tellurium: 1s2 2s2p6 3s2p6d10 4s2p6d10 5s2p4 Polonium: 1s2 2s2p6 3s2p6d10 4s2p6d10f14 5s2p6d10 6s2p4  Atomic and Ionic Radii: The atomic and ionic radius increases as we move from Oxygen to Polonium. Ionization Enthalpy: Ionization enthalpy decreases with increase in the size of the central atom. Therefore, it decreases as we move from Oxygen to Polonium since the size of the atom increases as we move down. Electron Gain Enthalpy: The electron gain enthalpy decreases with increase in the size of the central atom moving down the group. Oxygen molecule has a less negative electron gain enthalpy than sulfur. This is on the grounds that Oxygen, because of its compressed nature encounter more repulsion between the electrons effectively present and the approaching electron  Electronegativity: The electronegativity decreases as we move down the group due to increase in size Nature of the Group 16 Elements: The metallic properties increase in the order oxygen, sulfur, selenium, tellurium, or polonium. Oxygen and Sulfur are non-metals, Selenium and Tellurium are metalloids and Polonium is a metal under typical conditions. Polonium is a radioactive element. Allotropy: Each one of the element of group 16 displays allotropy. ✓ Oxygen has two allotropes: Oxygen and Ozone. ✓ Sulphur exists as many allotropic forms but only two of them are stable, which are: Rhombic Sulphur and Monoclinic Sulphur. ✓ Selenium and Tellurium are found in both amorphous and crystalline forms.  The Melting and Boiling Points: As the atomic size increases from oxygen to tellurium, the melting and boiling points also increase. The huge distinction between the melting and boiling points of oxygen and sulfur might be clarified on the premise that oxygen exists as a diatomic atom (O2) while sulfur exists as a polyatomic particle (S8). Oxidation States: The group 16 elements have a configuration of ns2 np4 in their outer shell, they may accomplish noble gas configuration either by the gain of two electrons, framing M-2 or by sharing two electrons, in this manner shaping two covalent bonds. Thus, these elements indicate both negative and positive oxidation states. The regular oxidation states showed by the elements of group 16 incorporate -2, +2, +4 and + 6.  Oxygen differs from sulfur in chemical properties due to its small size. The differences between O and S are more than the differences between other members. Oxygen is paramagnetic because there are unpaired electrons in O2 molecules. Sources ▪ Large-scale production of oxygen is by fractional distillation of liquid air. Liquid oxygen is stored and shipped at its boiling point of -183oC in vacuum- walled bottles. ▪ The Frasch process is used to mine sulfur from underground deposits. A well is drilled into a sulfur bed and a set of concentric tubes installed. Superheated water melts the sulfur. Compressed air forces it to the surface. ▪ Sulfur is also produced from hydrogen sulfide, H2S, and sulfur dioxide, SO2. 2H2S(g) + SO2(g) → 2H2O(l) + 3S(s) ▪ Selenium and tellurium are by-products of the processing of sulfide ores for other metals. ▪ Polonium is formed by the radioactive decay of radium in minerals such as pitchblende. Oxidation states and trends in chemical reactivity: 1. Since electronegativity of oxygen is very high, it shows only negative oxidation state as –2 except in the case of OF2 where its oxidation state is +2 2. The stability of + 6 oxidation state decreases down the group and stability of + 4 oxidation state increase (inert pair effect) Anomalous behavior of oxygen: 1. The anomalous behavior of oxygen, like other members of p-block present in second period is due to its small size and high electronegativity. One typical example of effects of small size and high electronegativity is the presence of strong hydrogen bonding in H2O which is not found in H2S. 2. The absence of d orbitals in oxygen limits its covalency to four and in practice, rarely exceeds two. On the other hand, in case of other elements of the group, the valence shells can be expanded and covalence exceeds four. Reactivity with hydrogen All the elements of Group 16 form hydrides of the type H2E (E =O, S, Se, Te, Po). H2O, H2S, H2Se, H2Te : Bond Dissociation Energy (BDE) decreases Stability decreases Acidity increases Reducing nature increases Reactivity with oxygen: All these elements form oxides of the EO2 and EO3 types where E = S, Se, Te or Po. Ozone (O3) and sulfur dioxide (SO2) are gases while selenium dioxide (SeO2) is solid. Reducing property of dioxide decreases from SO2 to TeO2; SO2 is reducing while TeO2 is an oxidizing agent Both types of oxides are acidic in nature. Reactivity towards the halogens. The stability of the halides decreases in the order F– > Cl– > Br– > I–. Among the hexahalides, hexafluorides are the only stable halides. They have octahedral structure. Sulphur hexafluoride, SF6 is exceptionally stable for steric reasons. Tetrafluorides, have sp3d hybridization and thus, have trigonal bipyramidal structures in which one of the equatorial positions is occupied by a lone pair of electrons. This geometry is also regarded as see-saw geometry.  Dihalides have sp3 hybridization and thus, have tetrahedral structure. The well known monohalides are dimeric in nature. Examples are S2F2, S2Cl2, S2Br2, Se2Cl2 and Se2Br2. These dimeric halides undergo disproportionation as given below: 2Se2Cl2 → SeCl4 + 3Se Important compounds and reactions Oxygen reacts with almost all other elements to form oxides. Example: Ozone, O3, is produced directly from oxygen, O2, during lightening strikes. Oxygen is necessary for releasing energy from fuels, such as glucose, in organisms. Oxygen is used to produce steel and to oxidize hydrogen in fuel cells. Sulfur compounds often have unpleasant odors. Hydrogen sulfide, H2S, smells like a rotten egg. It forms when metallic sulfides and hydrochloric acid react.  Concentrated sulfuric acid, H2SO4, is a strong dehydrating agent. Example: So people will know when there is a natural gas leak, ethyl mercaptan, CH3CH2SH, is added to supplies of odorless natural gas. Sodium thiosulfate, Na2S2O3, also known as hypo, is used in the development of film. The addition of cadmium selenide, CdSe, gives glass a beautiful ruby color. DIOXYGEN Oxides OZONE Ozone is an allotropic form of oxygen. It is too reactive to remain for long in the atmosphere at sea level. At a height of about 20 kilometers, it is formed from atmospheric oxygen in the presence of sunlight. This ozone layer protects the earth’s surface from an excessive concentration of ultraviolet (UV) radiations. Properties of Ozone ▪ Ozone is thermodynamically unstable with respect to oxygen since its decomposition into oxygen results in the liberation of heat (∆H is negative) and an increase in entropy (∆S is positive). These two effects reinforce each other, resulting in large negative Gibbs energy change(∆G) for its conversion into oxygen. ▪ Due to the ease with which it liberates atoms of nascent oxygen (O3 → O2 + [O] ), it acts as a powerful oxidizing agent. For example, it oxidizes lead sulphide to lead sulphate and iodide ions to iodine.  Estimation of ozone: When ozone reacts with an excess of potassium iodide solution buffered with a borate buffer (pH 9.2), iodine is liberated which can be titrated against a standard solution of sodium thiosulphate. This is a quantitative method for estimating O3 gas. Nitrogen oxides (particularly nitric oxide) combine very rapidly with ozone and there is, thus, the possibility that nitrogen oxides emitted from the exhaust systems of supersonic jet aero planes might be slowly depleting the concentration of the ozone layer in the upper atmosphere. NO + O3 → NO2 + O2 Uses of Ozone: Ozone is used as a germicide, disinfectant and for sterilizing water. It is also used for bleaching oils, ivory, flour, starch, etc. It acts as an oxidizing agent in the manufacture of potassium permanganate. SULFUR Sulfur is a solid at room temperature and 1 atm pressure. It is usually yellow, tasteless, and nearly odorless. It exists naturally in a variety of forms, including elemental sulfur, sulfides, sulfates, and organosulfur compounds Sulfur is unique in its ability to form a wide range of allotropes, more than any other element in the periodic table. The most common state is the solid S8 ring, as this is the most thermodynamically stable form at room temperature. Sulfur exists in the gaseous form in five different forms (S, S2, S4, S6, and S8). In order for sulfur to convert between these compounds, sufficient heat must be supplied. ▪ Yellow rhombic sulfur = Monoclinic sulfur Transition temp = 369 K ▪ Both rhombic and monoclinic sulfur have S8 molecules. ▪ At elevated temperatures (~1000 K), S2 is the dominant species and is paramagnetic like O2 ▪ In vapor state sulfur partly exists as S2 molecule which has two unpaired electrons in the antibonding * orbitals like O2 and, hence, exhibits paramagnetism  Two common oxides of sulfur are sulfur dioxide (SO2) and sulfur trioxide (SO3). Sulfur dioxide is formed when sulfur is combusted in air, producing a toxic gas with a strong odor. These two compounds are used in the production of sulfuric acid, which is used in a variety of reactions. Sulfur also exhibits a wide range of oxidation states, with values ranging from -2 to +6. It is often the central ion in a compound and can easily bond with up to 6 atoms. In the presence of hydrogen it forms the compound hydrogen sulfide, H2S, a poisonous gas incapable of forming hydrogen bonds and with a very small dipole moment. Hydrogen sulfide can easily be recognized by its strong odor that is similar to that of rotten eggs, but this smell can only be detected at low, nontoxic concentrations.  A variety of sulfur-containing compounds exist, many of them organic. The prefix thio- in from of the name of an oxygen-containing compound means that the oxygen atom has been substituted with a sulfur atom. General categories of sulfur-containing compounds include thiols (mercaptans), thiophenols, organic sulfides (thioethers), disulfides, thiocarbonyls, thioesters, sulfoxides, sulfonyls, sulfamides, sulfonic acids, sulfonates, and sulfates SELENIUM Selenium appears as a red or black amorphous solid, or a red or grey crystalline structure; the latter is the most stable. Selenium has properties very similar to those of sulfur; however, it is more metallic though it is still classified as a nonmetal. It acts as a semiconductor and therefore is often used in the manufacture of rectifiers, which are devices that convert alternating currents to direct currents. Selenium is also photoconductive, which means that in the presence of light the electrical conductivity of selenium increases. Selenium is also used in the drums of laser printers and copiers. Selenium is now used in “xerox” machines and laser printers. They use a drum coated with selenium that is exposed to an electric field. The regions on the drum that are exposed to high light intensity lose their charge. Toner powder adheres only to the charged areas of the drum which correspond to the printed areas on the page. Photocopiers  It is rare to find selenium in its elemental form in nature; it must typically be removed through a refining process, usually involving copper. Selenium is often found in soils and in plant tissues that have bio- accumulated the element. In large doses, the element is toxic; however, many animals require it as an essential micronutrient. Selenium atoms are found in the enzyme glutathione peroxidase, which destroys lipid-damaging peroxides. In the human body it is an essential cofactor in maintaining the function of the thyroid gland. TELLURIUM ▪ Tellurium is the metalloid of the oxygen family, with a silvery white color and a metallic luster similar to that of tin at room temperature. ▪ Like selenium, it is also displays photoconductivity. ▪ Tellurium is an extremely rare element, and is most commonly found as a telluride of gold. ▪ Tellurium is often used in metallurgy in combination with copper, lead, and iron. ▪ Tellurium is used in solar panels and memory chips for computers. ▪ It is not toxic or carcinogenic; however, when humans are exposed to too much of it they develop a garlic-like smell on their breaths. POLONIUM Polonium is a very rare, radioactive metal. There are 33 different isotopes of the element and all of the isotopes are radioactive. Polonium exists in a variety of states, and has two metallic allotropes ( and ) It dissolves easily into dilute acids. Polonium does not exist in nature in compounds, but it can form synthetic compounds in the laboratory. It is used as an alloy with beryllium to act as a neutron source for nuclear weapons. Polonium is a highly toxic element. Crystal structure cubic α-Po Crystal structure rhombohedral -Po  The radiation it emits makes it very dangerous to handle. It can be immediately lethal when applied at the correct dosage, or cause cancer if chronic exposure to the radiation occurs. Methods to treat humans who have been contaminated with polonium are still being researched, and it has been shown that chelation agents could possibly be used to decontaminate humans. Summary The chalcogens have no stable metallic elements. The tendency to catenate, the strength of single bonds, and the reactivity all decrease moving down the group. Because the electronegativity of the chalcogens decreases down the group, so does their tendency to acquire two electrons to form compounds in the −2 oxidation state. The lightest member, oxygen, has the greatest tendency to form multiple bonds with other elements. Oxygen does not form stable catenated compounds, due to repulsions between lone pairs of electrons on adjacent atoms. Because of its high electronegativity, the chemistry of oxygen is generally restricted to compounds in which it has a negative oxidation state, and its bonds to other elements tend to be highly polar.  Metal oxides are usually basic, and nonmetal oxides are acidic The reactivity, the strength of multiple bonds to oxygen, and the tendency to form catenated compounds all decrease down the group, whereas the maximum coordination numbers increase. Because Te=O bonds are comparatively weak, the most stable oxoacid of tellurium contains six Te–OH bonds. The stability of the highest oxidation state (+6) decreases down the group. Double bonds between S or Se and second-row atoms are weaker than the analogous C=O bonds because of reduced orbital overlap. The stability of the binary hydrides decreases down the group. References Cotton F.A., G. Wilkinson, and Gauss, P.A, Basic inorganic Chemistry, Wiley & Sons., Rayner-Canham, G., Descriptive Inorganic Chemistry, 2nd edition, W.H Freeman & Co, New York, Shriver, D.F and Atkin, P.W., 2006, Inorganic Chemistry, 4th ed, WH Freeman & Co, New York Suyanta, 2013, Kimia Unsur, Gadjah Mada University Press, Yogyakarta https://chem.libretexts.org/Bookshelves/Inorganic_Chemistry/Modules_and_Websites_(Inorgani c_Chemistry)/Descriptive_Chemistry/Elements_Organized_by_Group/Group_16%3A_The_Oxyge n_Family/1Group_16%3A_General_Properties_and_Reactions http://taladev.com/ebook/products/0-13-190443-4/ddref_chem05_eh08.pdf https://www.unf.edu/~michael.lufaso/chem4612/chapter16.pdf http://wps.prenhall.com/wps/media/objects/3313/3392904/blb2206.html https://fns.uniba.sk/fileadmin/prif/chem/kag/Bakalar/vch_noga/GEN_INORG_CHEM14b.pdf https://www.toppr.com/guides/chemistry/the-p-block-elements/group-16-elements/