Group 17 Elements PDF

Group 17 Elements  Group 17 Elements - the Halogens ▪ Much of the important chemistry of the group 17 elements can be understood on the basis of their electronic structure and electronegativity. ▪ Since the elements have a [core]ns2 np5 electron configuration, neutral group...

Group 17 Elements  Group 17 Elements - the Halogens ▪ Much of the important chemistry of the group 17 elements can be understood on the basis of their electronic structure and electronegativity. ▪ Since the elements have a [core]ns2 np5 electron configuration, neutral group 17 compounds can form up to seven bonds. This provides for several possible oxidation states (with a complete octet of electrons around the group 17 atom) although -1 is the most common. ▪ The structures of the poly halides are the typical examples used for VSEPR theory.  Group 17 Elements The group 17 elements include fluorine (F), chlorine (Cl), bromine (Br), iodine (I) and astatine (At) from the top to the bottom. They are called “halogens” because they give salts when they react with metals Group 7A(17): The Halogens Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display. Electronic Configuration of Group 17 Elements The valence shell electronic configuration of these electrons is ns2np5. Thus, there are 7 electrons in the outermost shell of these elements. The element misses out on the octet configuration by one electron. Thus, these elements look out to either lose one electron and form a covalent bond or gain one electron and form an ionic bond. Therefore, these are very reactive non-metals. SATP: Standard Ambient Temperature and Pressure. Astatine ▪ Astatine is a radioactive element with an atomic number of 85 and symbol At. ▪ Its possible oxidation states include: -1, +1, 3, 5 and 7. ▪ It is the only halogen that is not a diatomic molecule and it appears as a black, metallic solid at room temperature. ▪ Astatine is a very rare element, so there is not that much known about this element. ▪ In addition, astatine has a very short radioactive half-life, no longer than a couple of hours. ▪ It was discovered in 1940 by synthesis. ▪ Also, it is thought that astatine is similar to iodine. However, these two elements are assumed to differ by their metallic character. Atomic Properties 1) Ionic and Atomic Radii The nuclear and atomic radii of these elements keep on increasing as we move down the group. This happens because of the addition of an extra energy level. They have the minimal atomic radii compared to the other elements in the related periods. This can be attributed to the fact that their atomic charge is quite powerful. 2) Ionization Enthalpy These elements have higher ionization enthalpy. This value keeps on diminishing as we move down the group. This happens because of the increase in the size of the nucleus. However, it is interesting to note that fluorine has the highest ionization enthalpy than any other halogen. 3) Electron Gain Enthalpy The electron gain enthalpy of these elements becomes less negative upon moving down the group. Fluorine has lesser enthalpy than chlorine. We can attribute it to the small size and the smaller 2p sub-shell of the atom of fluorine. 4) Electronegativity The halogens exhibit high electronegativity values. However, it diminishes slowly on moving down the group from fluorine to iodine. This can be attributed to the increase in nuclear radii upon moving down the group. Physical Properties: Physical state: The group 17 elements are found in diverse physical states. For example, Fluorine and Chlorine are gases. On the other hand, Bromine is a liquid and Iodine is solid. Color: These elements have a variety of colors. For example, while Fluorine is pale yellow in color, Iodine is dark violet in color. Solubility: Florine and Chlorine are soluble in water. On the other hand, Bromine and Iodine are very less soluble in water. Melting and boiling points: Melting and boiling points of these elements increase as we move down the group from Fluorine to Iodine. Thus, Fluorine has the lowest boiling and melting points. Chemical Properties 1) Oxidizing Power All the halogens are great oxidizing agents. Of the list, fluorine is the most powerful oxidizing agent. It is capable of oxidizing all the halide particles to halogen. The oxidizing power reduces as we move down the group. The halide particles also act as reducing agents. However, their reducing capacity decreases down the group as well. 2) Reaction with Hydrogen All halogens react with hydrogen and produce acidic hydrogen halides. The acidity of these hydrogen halides increases from HF to HI. Fluorine reacts violently and chlorine requires the sunlight. On the other hand, bromine reacts upon heating and iodine needs a catalyst. 3) Reaction with Oxygen Halogens react with oxygen to form oxides. However, it has been found that the oxides are not steady. Beside oxides, halogens also form a number of halogen oxoacids and oxyanions. 4) Reaction with Metals As halogens are very reactive, they react with most of the metals instantly and form the resulting metal halides. For example, sodium reacts with chlorine gas and forms sodium chloride. This process is an exothermic one and gives out a bright yellow light and a lot of heat energy. 2 Na(s) + Cl2(g) → 2 NaCl(s) Metal halides are ionic in nature. This is because of the high electronegative nature of the halogens and high electro positivity of the metals. This ionic character of the halides reduces from fluorine to iodine. (MF > MCl > MBr > MI) Reactivity of the Halogens The reactivities of the halogens decrease down the group ( At < I < Br < Cl < F). This is due to the fact that atomic radius increases in size with an increase of electronic energy levels. This lessens the attraction for valence electrons of other atoms, decreasing reactivity. This decrease also occurs because electronegativity decreases down a group; therefore, there is less electron "pulling." In addition, there is a decrease in oxidizing ability down the group.  Reactivity of the Halogens A halogen atom needs only one electron to fill its valence shell. Halogens are therefore very reactive elements. The halogens display a wide range of electronegativities, but all are electronegative enough to behave as nonmetals. A halogen will either - gain one electron to form a halide anion or - share an electron pair with a nonmetal atom. The reactivity of the halogens decreases down the group, reflecting the decrease in electronegativity. Hydrogen Halides and Halogen Oxoacids Hydrogen Halides A halide is formed when a halogen reacts with another, less electronegative element to form a binary compound. Hydrogen, for example, reacts with halogens to form halides of the form HX: Hydrogen Fluoride: HF Hydrogen Chloride: HCl Hydrogen Bromide: HBr Hydrogen Iodide: HI Hydrogen halides readily dissolve in water to form hydrohalic (hydrofluoric, hydrochloric, hydrobromic, hydroiodic) acids. The properties of these acids are given below: The acids are formed by the following reaction: HX (aq) + H2O (l) → X- (aq) + H3O+ (aq) All hydrogen halides form strong acids, except HF The acidity of the hydrohalic acids increases as follows: HF < HCl < HBr < HI  Hydrofluoric acid can etch glass and certain inorganic fluorides over a long period of time. It may seem counterintuitive to say that HF is the weakest hydrohalic acid because fluorine has the highest electronegativity. However,​ the H-F bond is very strong; if the H-X bond is strong, the resulting acid is weak. A strong bond is determined by a short bond length and a large bond dissociation energy. Of all the hydrogen halides, HF has the shortest bond length and largest bond dissociation energy. Halogen Oxoacids A halogen oxoacid is an acid with hydrogen, oxygen, and halogen atoms. The acidity of an oxoacid can be determined through analysis of the compound's structure. The halogen oxoacids are given below: Hypochlorous Acid: HOCl Chlorous Acid: HClO2 Chloric Acid: HClO3 Perchloric Acid: HClO4 Hypobromous Acid: HOBr Bromic Acid: HBrO3 Perbromic Acid: HBrO4 Hypoiodous Acid: HOI Iodic Acid: HIO3 Metaperiodic Acid: HIO4; H5IO6 In each of these acids, the proton is bonded to an oxygen atom; therefore, comparing proton bond lengths is not useful in this case. Instead, electronegativity is the dominant factor in the oxoacid's acidity. Acidic strength increases with more oxygen atoms bound to the central atom. Molecular shapes of the main types of interhalogen compounds. ClF ClF3 linear, XY T-shaped, XY3 BrF5 IF7 Square Pentagonal pyramidal, XY5 bipyramidal, XY7  Chlorine oxides. lone e- dichlorine monoxide Cl2O dichlorine heptaoxide Cl2O7 chlorine dioxide ClO2 Applications of Halogens Fluorine: Although fluorine is very reactive, it serves many industrial purposes. For example, it is a key component of the plastic polytetrafluoroethylene (called Teflon- TFE by the DuPont company) and certain other polymers, often referred to as fluoropolymers. Chlorofluorocarbons (CFCs) are organic chemicals that were used as refrigerants and propellants in aerosols before growing concerns about their possible environmental impact led to their discontinued use. Hydrochlorofluorocarbons (HFCs) are now used instead. Fluoride is also added to toothpaste and drinking water to help reduce tooth decay. Fluorine also exists in the clay used in some ceramics. Fluorine is associated with generating nuclear power as well. In addition, it is used to produce fluoroquinolones, which are antibiotics. Table 1: Important Inorganic Compounds of Fluorine Compound Uses Na3AlF6 Manufacture of aluminum BF3 Catalyst CaF2 Optical components, manufacture of HF, metallurgical flux ClF3 Fluorinating agent, reprocessing nuclear fuels HF Manufacture of F2, AlF3, Na3AlF6, and fluorocarbons LiF Ceramics manufacture, welding, and soldering NaF Fluoridating water, dental prophylaxis, insecticide SF6 Insulating gas for high-voltage electrical equipment SnF2 Manufacture of toothpaste UF6 Manufacture of uranium fuel for nuclear reactors Chlorine: Chlorine has many industrial uses. It is used to disinfect drinking water and swimming pools. Sodium hypochlorite (NaClO) is the main component of bleach. Hydrochloric acid, sometimes called muriatic acid, is a commonly used acid in industry and laboratories. Chlorine is also present in polyvinyl chloride (PVC), and several other polymers. PVC is used in wire insulation, pipes, and electronics. In addition, chlorine is very useful in the pharmaceutical industry. Medicinal products containing chlorine are used to treat infections, allergies, and diabetes. The neutralized form of hydrochloride is a component of many medications. Chlorine is also used to sterilize hospital machinery and limit infection growth. In agriculture, chlorine is a component of many commercial pesticides: DDT (dichlorodiphenyltrichloroethane) was used as an agricultural insecticide, but its use was discontinued. Bromine: Bromine is used in flame retardants because of its fire-resistant properties. It also found in the pesticide methyl bromide, which facilitates the storage of crops and eliminates the spread of bacteria. However, the excessive use of methyl bromide has been discontinued due to its impact on the ozone layer. Bromine is involved in gasoline production as well. Other uses of bromine include the production of photography film, the content in fire extinguishers, and drugs treating pneumonia and Alzheimer's disease. Iodine: Iodine is important in the proper functioning of the thyroid gland of the body. If the body does not receive adequate iodine, a goiter (enlarged thyroid gland) will form. Table salt now contains iodine to help promote proper functioning of the thyroid hormones. Iodine is also used as an antiseptic. Solutions used to clean open wounds likely contain iodine, and it is commonly found in disinfectant sprays. In addition, silver iodide is important for photography development. Astatine: Because astatine is radioactive and rare, there are no proven uses for this halogen element. However, there is speculation that this element could aid iodine in regulating the thyroid hormones. Also, 211At has been used in mice to aid the study of cancer. References: Cotton F.A., G. Wilkinson, and Gauss, P.A, Basic inorganic Chemistry, Wiley & Sons., Cotton, F.A., Wilkinson, G., Murillo, C.A., Bochmann, M., 1999, Advanced Inorganic Chemistry, 6th edition, John Wiley & Sons, Inc., New York, 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 https://www.toppr.com/guides/chemistry/the-p-block-elements/group-17-elements/ https://www.vedantu.com/chemistry/group-17-elements  Group 18 Elements - the Noble Gases ▪ The chemistry of the group 18 elements seems to defy their electronic structure. ▪ Since the elements have a [core]ns2 np6 electron configuration with a complete octet, one would predict that there would be no chemistry for the noble gases. ▪ However, numerous group 18 compounds are known, although they may be very unstable and explosive! ▪ Understanding the reactivity of group 18 compounds requires an examination of their ionization potentials. Group 8A(18) Elements: The Noble Gases Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.  Noble Gases Noble gases have a full valence shell. The noble gases are the smallest elements in their respective periods, with the highest ionization energies. Atomic size increases down the group and IE decreases. Noble gases have very low melting and boiling points. Occurrence On account of their inert nature, the noble gases always occur in the free state. Argon is the most abundant noble gas in the atmosphere, while radon is not present in atmosphere. Helium, Argon and Neon are also found as constituents of dissolved gas of certain spring waters. The Nobel Gases Noble gases have a full valence shell. The noble gases are the smallest elements in their respective periods, with the highest ionization energies. Noble gases have very low melting and boiling points. Xe is the most reactive noble gas and exhibits all even oxidation states from +2 to +8. Very low electronegativities No colour, odor, or flavour under ordinary conditions Non flammable At low pressure, they will conduct electricity and fluoresce  All the noble gases occur in the atmosphere as monatomic gases. They are monatomic, which means they exist as individual atoms. Most other gases are diatomic. Together they make up 1% (by mass) of the atmosphere. Argon is the third most abundant gas in the atmosphere after N and O. Sources of the Noble Gases All of the noble gases except He and Rn are obtained by the fractional distillation of liquid air. The major source of helium is from the cryogenic separation of natural gas. Radon, a radioactive noble gas, is produced from the radioactive decay of heavier elements, including radium, thorium, and uranium. Element 118 (Oganesson, Og) is a man-made radioactive element, produced by striking a target with accelerated particles. The Atomic and Physical Properties Atomic mass, boiling point, and atomic radii INCREASE down a group in the periodic table. The first ionization energy DECREASES down a group in the periodic table. The noble gases have the largest ionization energies, reflecting their chemical inertness. Down Group 18, atomic radius and interatomic forces INCREASE resulting in an INCREASED melting point, boiling point, energy of vaporization, and solubility. The INCREASE in density down the group is correlated with the INCREASE in atomic mass. Because noble gases’ outer shells are full, they are extremely stable, tending not to form chemical bonds and having a small tendency to gain or lose electrons. Selected Properties of Group 8 A Elements Important Physical Properties Electronic configuration. These gases have highly stable ns2np6 configuration. Atomic and ionic radius On moving down the group, the atomic and ionic radii increases with increase in size. Boiling points On moving down the group, the boiling point increases with increase in size. Ionization energy and electron affinity. These gases have stable ns2np6 (fully filled) configuration, thus, have no tendency to add or lose electron. Hence, their ionization energy is very high and electron affinity is zero. Heat of vaporization and polarizability. ✓ These gases possess very low values of heat of vaporization due to the presence of very weak van der Waals’ force between their monoatomic molecules. However, the value of heat of vaporization increases with atomic number. ✓ On moving down the group, the polarizability increases.  Solubility Noble gases are slightly soluble in water and their solubility generally increases on moving down the group. Adsorption Except helium, all the noble gases are adsorbed by charcoal and the ease of adsorption increases down the group. Liquification Ease of liquification increases down the group from He to Rn due to increase in intermolecular forces. Conductivity Noble gases have high electrical conductivity at low pressure. These gases are monoatomic gases (Cp/Cv = 1.667). Chemical Properties In general, noble gases are not very reactive. Their inertness to chemical reactivity is attributed to the following reasons The noble gases have completely filled ns2np6 electronic configurations in their valence shells. The noble gases have very high ionization energies. The electron affinities of noble gases are almost zero. Preparation and properties of some compounds  Partial hydrolysis of XeF6 gives oxyfluorides, XeOF4 and XeO2F2. XeF2, XeF4 and XeF6 are colorless crystalline solids and sublime readily at 298K. They are powerful fluorinating agents. XeF6 is extremely reactive. It cannot be stored in glass or quartz vessels as it readily reacts with SiO2 present in glass. XeO3 is a colorless, explosive solid and has a pyramidal shape. XeOF4 is a colorless, volatile liquid and has a square pyramidal shape. XeF2 is linear and XeF4 is square planar. XeF6 has a distorted octahedral shape. Compounds of the Nobel Gases (Reactivity): The elements have a complete octet, predict that there would be no chemistry for the noble gases. However, numerous group 18 compounds are known, although they may be very unstable and explosive. He and Ne are chemically inert and they do not form any compounds. ✓ Their chemical inertness is due to very high ionization energy, zero electron affinity and the absence of vacant d-orbitals in valence shell. Ar, Kr and Xe will show some reactivity ✓ due to low ionization potentials and presence of vacant d-orbitals in valence shell. Xe is more reactive than Ar and Kr ✓ due to it's low ionisation energy. Radon is radioactive and it will not show chemical reactivity.  Krypton forms only one known stable neutral molecule KrF2. Xe shows tendency to lose electrons in many of it's reactions. Therefore, Xe combines with only more electronegative elements like F and O or electronegative groups. Xe does not combine with less electronegative elements like Cl2 or N2. Uses ▪ The noble gases are used to form inert atmospheres, to protect specimens, and to prevent chemical reactions (to prevent oxidation). ▪ Argon is also used to fill some types of light bulbs, where it conducts heat away from the filament. ▪ Krypton gives an intense white light when an electrical current is passed through it and it is used in airports for there runway lights used: In lasers for eye surgery, to stop bleeding on the retina. In lighthouses and other types of lamps. ▪ Xenon is used: In various types of electron tubes, lamps, lasers and in high speed photographic flash tubes ▪ Radon is used: To treat cancer by radiotherapy, because it is radioactive. However, because radon is radioactive, it is also an environmental hazard.  Xenon Xenon is unique for being the first noble gas element to be synthesized into a compound. Discovered on 1898 by Sir William Ramsay. Xenon is present to a small extent in the atmosphere (less than one ppm by volume). Metallic xenon is produced by applying several hundred kilobars of pressure. In 1962 the first noble gas compound was produced by Neil Bartlett, combining xenon, platinum and fluorine. It is now possible to produce xenon compounds in which the oxidation states range from +2 to +8. Most of the known xenon compounds contain the strongly reducing fluorine or oxygen atoms. Xenon Compounds and their Molecular Structure Fluorine is the only element that directly reacts with Xenon. Xenon-fluorine compounds XeF2, XeF4 and XeF6 Preparation : By the direct reaction of elements under appropriate experimental conditions. Properties: They are readily hydrolyzed even by traces of water. 673K, 1bar 873K, 7bar 573K, 60-70bar  Xenon Halides are reactive with other compounds such as water. XeF2+3H2O → XeO3+ 6HF The Xe has a total of 8 outside shell electrons while the Fluorine 7 valence electrons. Xe's outside shell electrons are very far away from the center, therefore Xenon cannot possibly attract all of the electrons. Fluorine is smaller, therefore is has a stronger positive attraction to the few electrons it has left. Fluorine is the only element that reacts with Xe because it is the most electronegative. In other words, it is the only element that is strong enough to pull electrons out of the stable xenon. Xenon-oxygen compounds XeO3 XeOF4 and XeO2F2 XeO3 Hydrolysis of XeF4 and XeF6 with water gives XeO3. 6XeF4 + 12 H2O → 4Xe + 2XeO3 + 24 HF + 3 O2 XeF6 + 3 H2O → XeO3 + 6 HF XeO3 is a colorless explosive solid and has a pyramidal sp3 molecular structure Hybridization  XeOF4 and XeO2F2 Partial hydrolysis of XeF6 gives oxyfluorides, XeOF4 and XeO2F2. XeF6 + H2O → XeOF4 + 2 HF Xenon oxytetrafluoride XeF6 + 2 H2O → XeO2F2 + 4HF Xenon dioxydifluoride XeOF4 is a colorless volatile liquid and has a square pyramidal molecular structure Square pyramidal  ↑↓ Xe F F O O sp3d sp3d2 Hybridization Hybridization XeO2F2 XeOF4 Structures of several known xenon compounds https://fac.ksu.edu.sa/sit es/default/files/10- _group_18- _nobel_gases.pptx References: Cotton F.A., G. Wilkinson, and Gauss, P.A, Basic inorganic Chemistry, Wiley & Sons., Cotton, F.A., Wilkinson, G., Murillo, C.A., Bochmann, M., 1999, Advanced Inorganic Chemistry, 6th edition, John Wiley & Sons, Inc., New York, 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 https://www.toppr.com/guides/chemistry/the-p-block-elements/group-18-elements/ 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/  Increase electro Increase electro negativity / positive character Increase ionic / Increase ionic Increase electro positive character / Increase ionic B is non metallic elements (covalent with small metallic character) Al, Ga, In dan Tl are fairly metal. The electropositive character increases from Al to Tl (Thallium).  Elements Elektonic Configiration Ion Radii Ionization Elektro Oxsidatio (M3+)(Ao) Potensial (eV) negativity n state 5B [He] 2s2 2p1 0,20 8,3 2,0 3 13Al [Ne] 3s2 3p1 0,52 6,0 1,5 (1),3 31Ga [Ar] 3d10 4s2 4p1 0,60 6,0 1,6 1, 3 49In [Kr] 4d10 5s2 5p1 0,81 5,8 1,7 1, 3 81Tl [Xe] 4f14 5d10 6s2 6p1 0,95 6,1 1,8 1, 3 q Boron is non metal, but Al, Ga, In, Tl are reactive metal q The all elements of III A group show an oxidation state of +3, oxidation state of +1 are also showing for Al, Ga, In dan Tl q Gallium as apparently divalen, Indium, Thalium stable at univalen q The increase in electropositivity from B to Al is the usual trend associated with increased size. Filling of electrovalen B and Al after the s-block elements while Ga, In dan Tl follow after the d-block. The extra d electrons do not shield the nuclear charge very effectively, so that the orbital electrons are firmly held, and the metals properties are less electropositive First ionization energies (kJ mol-1) for Group 3 elements 8,3 (2,241) Al, Ga, In dan Tl have 6,0 (4,06) 6,0 (6,22) 5,8 6, similar value of 1 ionization potensial 6,22 efective nuclear charge  8,3 ionization potensial Effective nuclear charge increased 3B [He] 2s2 2p1 13Al [Ne] 3s2 3p1 8 proton diffrences with B 31Ga [Ar] 3d10 4s2 4p1 18 proton diffrences with Al 49In [Kr] 4d10 5s2 5p1 18 proton diffrences with Ga 81Th [Xe] 4f14 5d10 6s2 6p1 32 proton diffrences with In Orbital d and f do not shield the nuclear charge very effectively, so that the orbital electrons are more firmly held and the character electropositive show similar value to the Aluminium eventhough the size of element are bigger  Properties of III A Group Elements Elements Melting Point Boiling Point ε P.I (eV) Atomic radii Ionic radii ( oC ) ( oC ) ( pm ) ( pm ) B 2300 2550 2,01 8,3 85 - Al 460 2467 1,47 6,0 143 54 Ga 30 2403 1,82 6,0 135 62 In 157 2080 1,49 5,8 142 80 Tl 303 1457 1,44 6,1 170 88 q The melting point of Boron is highest in group due to covalent bonding to form polymer structur of Boron q The melting point: Ga < In < Tl , related to the size of elements q Electronegativity Ga is high because of effective nuclear charge of Ga is higher q B and Al has one shell diffrences consist of 8 proton (shell number 2, 2s, 2p), while Al dgn Ga has one shell diffrences 18 proton (shell number 3, 3s, 3p, 3d)  III A Group Elements q Boron occurs naturally as Borax, Na2B4O7  4H2O and Na2B4O7  10H2O q Acidification of Borax to give Borax acid, H3BO3 or hidrated B2O3 q Crystal Boron is obatained by reduction of BCl3 and H2 over heated Tl wire q Amorphous Boron is obtained by reduction of B2O3 with Na and Mg q Crystal Boron formed in various structure, consist of unit B12 (icosahedral) and formed cluster polyhedra due to catenation properties q Boron has low oxidation number as hydride and halide compound Structure Icosahedral B12 q After oxygen, silicon, Al is the most common element in the Earth’s Crust q Aluminium naturally exist as Al2O3H2O Bauxite mineral soluble to base (NaOH) and reprecipated as Al2O33H2O and the metal form was formed by electrolysing of alumina in molten Na3AlF6 q Abundance of Ga, In and Tl elements are fairly low, occuring as minor components of a range of minerals, such as sphalerite (ZnS)  BORIDE COMPOUNDS q Boron has chemical properties like carbon dan silicon (polimeric properties) to form boride compound when reacted to metal q The formed structure of borida compounds are influenced by metal-boron ratio q Ratio 1 : 12 formed MB12, where boron has cluster of icosahedral Chain in MB as sheet at MB2 Cluster in MB6 Ratio 1 : 1 Rasio 1 : 2 Ratio 1 : 6  HALIDE COMPOUNDS B Al Melting Point Boiling Point Melting Point Boiling Point (oC) (oC) (oC) (oC) Fluoride - 127 - 101 - 1291 Chloride - 107 12,5 192(*) 180 Bromide - 46 91 97 255 Iodide 43 210 180 381 q Boron Halida is covalen and monomeric in gas formed, BCl3 exist is bensen solution q Aluminium fluorida, AlF3 is ionic, solid crystaline has high melting point As Chloride dan Bromide formed volatile compounds as dimer, at gas fase or non-polar solvent, halogen atom bonded tetrahedrally to the aluminium.  MONOMERIC BORON HALIDE (BX3) Planar molecule, empty orbital p perpendecular to the molecular planar BF3 and BCl3 can be occur -back donation from filled halogen p-orbital into the empty p orbital of boron can occur, does not work for Br and I. -back donation of BF3 is more effective than BCl3 due to less size of F. BCl3 is liquid and volatile, boiling point is 12,5oC How to reduce the volatility of BF3 ? (by reacted with Lewis base) q Trihalide, MX3, is strong Lewis acid, is easy to react with Lewis base to form new compound Lewis Base Lewis Base Lewis Base S BF3 Eter Structure Ion negative of BF4 structure Dietil sulfide BF3 Structure  AlX3 Compound q The structure of AlX3 is influenced to the halide type, AlF3 is a high-melting polymeric solid built from flouride-bridged of octahedral (AlF6) AlF6, similar to AlCl3 q Al2Cl6 are liquid and gas fase, as dimer compound q AlBr3 and AlI3 are all in dimer system. DIMER AlCl3 Electron deficient (2 electron and 3 centre) (2e,3c) B or Al X (halide) Electron deficient (2 electron and 3 centre) (2e,3c) Electron deficient compound AlCl3 becomes acid compound (Lewis Acid)  Lower Oxidation State Halides q All element of III A group form diatomic halides MX, except Tl, TlCl are unstable towards disproportionation to the metal and the trivalent halide q Boron halide with lower oxidation state exist as B2Cl4 consist of B-B bonding is made by passing electric discharge to BCl3 using mercury electrode or condensed Cu (Copper atoms) 2 BCl3 + 2 Hg B2Cl4 + Hg2Cl2 2 BCl3 + 2 Cu (atom) B2Cl4 + 2 CuCl2 Boron cluster of more halide are also formed as B4Cl4 as reaction aboved B8Cl8 and B9Cl9 (or B10Cl10, B11Cl11 and B12Cl12) Cluster Structure of boron chloride, B4Cl4 and B9Cl9 q AlCl, GaCl are formed by reacted Aluminium or Gallium metal with HCl gas at high temperature and low pressure to produce red color of AlCl or GaCl which are condensed at low temperature (77K) and following by heating to produce MCl3 q Gallium (II) as Ga2X62- (X = Cl, Br or I ) formed from metal electrolisis at strong acid solution. Bonding Ga-Ga has oxidation state of +2 and easy to oxidised by using X2 to be GaX4 Hydride dan Organometallic Compounds 1. Boron Hydrides q In group of gol III A, Boron produce more hydride compound than others by electron-deficient bonding of hydrogen  Table of Higher Boron hydride Compounds BnHn+4 BnHn+6 B2H6, B5H9, B6H10, B8H12, B10H14 B4H10, B5H11, B6H12, B8H14 q Boron hidride contain the same structural features as B2H6, with one or more B-B bonds present, as tetraborane (B4H10) or penta borane (B5H9) where have “open” polyhedral cluster structures, as shown above. q These higher boron hydrides can be formed by heating B2H6. Different reaction condition give different boron hydrides, as below o 100 120 C 2 B2 H 6     B4 H10  H 2 q There is an extensive series of boron hidride anions, which have closed polyhedral skeletons q [B6H6]2- and [B12H12]2- have octahedral B6 dan icosahedral B12 and terminal B-H bonds Carboranes are obtained by replacing one or more BH units with C atoms  The shapes of various borane clusters can be rationalized using Wade’s Rules which consider the total number of skeletal electron pair (SEP) available for cluster Wade’s Rules A Borane cluster can be built up from BH units (with terminal, i.e non- Bridging, hydrogens). Other hydrogen in the cluster bond in other ways The number of SEP is calculated as follows. 1. Determine the total number of valence electrons (three per B atom, one per H atom), adding and substracting electrons for any overall charge ( a C atom in a carborane controbutes four electron) 2. Substract two electrons for each BH (or CH) unit 3. The number of electrons, divided by 2, is the number of SEP, which determines the type of cluster 4. A cluster with n vertices occupied by n B atoms requires n+1 SEP for bonding. Such cluster are termed closo; examples are the octahedron and icosahedron with 7 and 13 SEP respectively Wade’s Rules 5. If there are n+1 SEP dan n-1 B atoms, the cluster is derived from a closo cluster, but one missing vertex, this called a nido cluster 6. If there are n+1 SEP and n-2 B atoms, the cluster has two missing vertices and is called an arachno cluster. The simple theory does not say unequivocally which vertices are missing 7. If there are more than n B atoms and n+1 SEP, the extra B atoms occupy capping positions over triangular faces 2. Hydride of other elements of III A group q For Aluminium hydride stabil as alane polimer, AlH3 q For Gallium, the reaction of the dimeric [H2GaCl]2 dgn Li[GaH4] gives [GaH3]n , mainly occurs as the rather thermally unstable diborane-like dimer Ga2H6 in the gas phase. q In some respects the chemistry of Gallane more resembles that of Borane than Alane, because the electronegativity of Gallane (1,8) is Higher than that of Aluminium (1,5) and closer to Boron (2,0). 3. Hydride Adducts q Hydride of Group IIIA elements form donor-stabilized monomers, MH3D (where D is the donor as illustrated by the formation of BH3 adducts B2H6 + 2D→ 2 BH3.D  Sodium tetrahydroborate (III), H H Na[BH4] is produced from H H trimethyl borate and sodium D H hydride, and is widely used as H Donor H reducing agent for carbonyl group. Acceptor MH4- Na[BH4] is soluble and relatively stable in water Oxide, Hydroxide and Oxyanion q The Oxide M2O3 of Group III A elements can be made by heating the elements in oxygen q Reaction of aqueous solutions of the metal trihalides with hydroxide gives the oxides in hydrated form. q Going down the group, there is a transition properties of acidic oxides, through amphoteric to basic, owing to increasing metallic character Oxide Properties B2O3 Weakly acidic 2O3 Amphoteric Increase electropositivity, Ga2O3 Amphoteric Increase metal properties In2O3 Weakly basic Tl2O3 Basic, oxidizing Layer Struktur of hydrogen bonding of Boric Acid [B(OH)3] Boric acid reacts with alcohols to form borate B(OH)3 + 3 MeOH→B(OMe)3 + 3H2O Structure of Borate Anion  Aluminium q There are two form of Al2O3 : the high temperatur as α (alpha) form and Lower as  (gamma) form has a more complex but open structure q Aluminium forms two hydrated oxides, MO(OH) and M(OH)3 q Aluminium oxide is amphoteric and dissolves in concentrated hydroxide solution to give aluminate solution containing the [Al(OH)4]- ion. Compounds with N : Borazine and Boron Nitride q B2H6 and NH3 initially form the expected H3NBH3 which gives borazine on warming to room temperature. q Borazine is isoelectronic with benzene (C6H6) (see the structure) q Borazine has also a delocalized structure, with all B-N bond lengths identical (144 pm). q Reactivity of borazine is rather different to that of benzene, due to polarized differences between B-N to C-C, and also there is a “well electronic“ due to empty orbital of boron q Borazine is sometimes called as “inorganic benzene“ q Reactivity of borazine is rather different to that of benzene, due to polarized differences between B-N to C-C, and also there is a “well electronic“ due to empty electron of Boron orbital q The reaction of NH4Cl with BCl3, gives B,B,B-trichloroborazine Structure of Borazine Structure of Benzene Borazine vs Benzene q Boron nitride (BN) is iso electronic with carbon. Both BN and Carbon form diamond and graphite structures. The graphite-like structure of BN has layers consisting of planar, hexagonal B3N3 rings (with BN distances similar to those in borazine. These layers stack on top each other : this contrasts with the staggered arrangement found in graphite Structure Boron Nitride Structure Graphite Boron Nitride vs Graphite THE END OF GROUP IIIA SLIDE