Lecture 4 - P, As, Sb, Bi Properties - PDF

Lecture 4 P, As, Sb and Bi:- Properties: - There are big differences between the chemistry of nitrogen and these elements in spite of the same electronic configuration of the outer shell they have. The reason of that is: 1. The nitrogen atom does not have d orbitals while the above elements have d...

Lecture 4 P, As, Sb and Bi:- Properties: - There are big differences between the chemistry of nitrogen and these elements in spite of the same electronic configuration of the outer shell they have. The reason of that is: 1. The nitrogen atom does not have d orbitals while the above elements have d orbitals. 2. The maximum coordination no. of nitrogen is (4) while others can use vacant d orbitals to form bonds which increase the number of electrons in their valance shells. Nitrogen can form double and triple bonds type (p-p). Phosphorous likes nitrogen (to be covalent character) in its compounds, while others like to form ionic compounds (increasing) from As to Bi, e.g. BiF3 is ionic. Acidic and basic properties change for the group elements (especially oxides) from acidic for P to basic for Bi. Oxidation state of these elements are trivalent or pentavalent, the stability of the trivalent oxidation state increases from P to As, Bi (III) is stable, but Bi 2O5 (Bi V) is difficult to prepare and is the least stable oxide in this group. Occurrence in Nature: - P is present as phosphate in many minerals like Apatite Ca5(PO4)3 (Cl, OH, F), As, Sb and Bi are rarely present in their elemental from mostly they present as sulfides. A- Phosphorous: - Prepared by reduction of phosphate rocks by coke and silica in an electrical furnace, it evaporates as P4 molecules, condensed under water forming white phosphorous: Phosphorous has three allotropic forms, white, red and black, each form has many shapes (at least 11 shapes). 1-White phosphorous presents in the liquid and solid states (P4), its structure is tetrahedral, so that the bonds of P4 molecule are weak and easy to break, which explain the activity of white phosphorus. 2-Black phosphorous has double layers (coupled) in which each P atom bonded to other three atoms. It can be prepared in its crystalline form by heating the white phosphorous under high pressure at 220-370 ˚c, in presence of Hg as catalyst. 3- Red phosphorous produced from heating the white type for many hours at 400˚c. The activity of phosphorous depends on its form, the white is more active it burns when exposes to air so it must keep under water, while the black and red phosphorous are stable in air, the black one is the less stable. Compounds:- 1- Hydrides:- All the group V elements form hydrides (MH3), each one prepared from the reaction of MCl3 and the metal hydride (to be wanted): - Phosphine and Arsine (PH3, AsH3) are prepared by the reaction of phosphides and arsenides of the metals with acids. SbH3 and BiH3 are unstable with temperature. Generally, the stability of hydrides decreases of the bond energy in the same direction. 2- Halides:- Group V elements form two types of halides, trihalides (MX3) and penta (MX5). a- MX3 MX3 of P, As, Sb and Bi (except PF3) are prepared from halogens with enough (or excess) quantity of the metals while PF3 is prepared by the reaction of ZnF2 with PCl3. Trihalides are mostly covalent from which one can conclude that they have relatively low boiling and melting points. The ionic character of these halides increases from P to Bi and for the central atom from I to F. 3- Oxides:- The group V elements have two types of oxides, tri and pentvalent oxides (+3 and +5), their basic properties increase with increasing the atomic no. P and As oxides are acidic, Sb oxides are amphoteric, while Bi oxides are basic. a- Phosphorous oxides: - They are prepared from the reaction of phosphorous vigorously with O2, their formation depends on the O2 quantity and P reacted. Increasing O2 give P2O5 while increasing P gives P2O3. Phosphorous atoms in P4O10 are occupying the corners of the tetrahedral shape while 6 oxygen atoms are at the sides (edges) of the tetrahedral, the other 4 oxygen atoms are bonded to P atoms along with the three axis. The 12 formed bonds between oxygen and phosphorous atoms are single ones, some others are double as in the structures: b-Oxo acids:- Phosphorous oxo acids: a- Hypophosphorous acid: It is a colorless crystal, m. p. 26.5cº, the structure is: b- Phosphorous acid H3PO3 Is prepared from PCl3 or P4O6 reaction with cold water. The pure acid melts at 70Cº (PKa=1.8) its structure is: Dibasic acid (can lose two H+ to form [HPO3]-2 ion), this acid and its salts are strong reducing agents. c-Phosphoric acid H3PO4 The most of the phosphorous compounds, prepared from H2SO4 with phosphate rocks or from P4O1O with water reaction. This acid is tribasic, has three types of salts, M3PO4, M2HPO4 and MH2PO4. d-Pyrophosphoric acid H4P2O7 produced as: Two molecules of phosphoric acid. Oxygen These elements have six electrons in the outer shell, e. g.:- 8O: 1S2 2S2 2P4 Oxygen forms compounds with all elements, except He, Ne, and may be Ar. It combines directly with other elements, except halogens and some Nobel gases, this takes place either at normal temperatures or high. Oxygen is the more found element, forms 50% by weight of the earth crust e.g. in water and silica which are the main components of earth. Oxygen is one of the second period, the outer shell saturated by 8 electrons by one of the following methods: 1. Gaining 2 es forming oxide ion (O2-). 2. Forming two single covalent bonds as in (R-O-R) or double bond as in (O=C=O). 3. Gaining electron in addition to form single covalent bond (OH-). 4. Forming 3 or 4 covalent bonds as in (R2-OH+). Oxides: The oxygen di compounds are called oxides, differ in their properties due to the nature of the bond bonding oxygen with other element, some of the compounds are ionic and covalent, others are between ionic and covalent properties. Hydroxide ion: It can be prepared in the aqueous solutions of oxides and peroxides of metals of high electropositive because of the hydrolysis reactions: In the solid-state hydroxide ion presents as a separated unit in the metal and alkaloid earth hydroxides, produced from the solvation of these ionic hydroxides in water. Lecture 5  Elements of Group VIA or 16 (The Oxygen Family, ns2np4) POSITION IN PERIODIC TABLE Group 16 or VIA of the extended form of periodic table consists of six elements oxygen (O), Sulphur (S). Selenium (Se), tellurium (Te), and polonium (Po). This family is known as oxygen family. These elements have six electrons in their valency shell and thus placed in the VIth group. 1- The elements oxygen and Sulphur are common while selenium, tellurium and polonium are comparatively rare. Oxygen is the most abundant element and is found both in free as well as in combined state. Oxygen makes up 20.9% by volume and 23% by mass of atmosphere. Most of the oxygen present in the atmosphere is produced by photosynthesis in plants. It also occurs in the form of ozone in the upper atmosphere which protects us from the harmful radiations of the sun. 2- Oxygen makes up 46.6% by mass of the earth's crust. Sulphur is the sixteenth most abundant element and constitutes 0.034% by mass of the earth's crust. It occurs mainly in combined form. The member, polonium is radioactive in nature. The inclusion of these elements in the same subgroup is justified on the basis of same electronic configuration and similarities as well as gradation in their physical and chemical properties. 1. Electronic Configuration The distribution of electrons in various energy shells of the atoms of these elements is given as below: All have six electrons in their outermost shell, i.e., they have electronic structure s2p4. 2. Physical Characteristics (a) Physical state: Oxygen is a gas while others are solids. Oxygen molecule is diatomic while the molecules of other elements are more complex. Sulphur, selenium and tellurium exist as staggered 8-atom rings. However, the tendency to exist in 8-atom rings is maximum with Sulphur and decreases as we go down the group. In Sg molecule, every Sulphur atom is in sp³ hybridized. [Oxygen atom has the tendency to form multiple bonds (pπ-pπ interaction) with other oxygen atom on account of small size while this tendency is missing in Sulphur atom. The bond energy of oxygen-oxygen double bond (O=O) is quite large (about three times that of oxygen-oxygen single bond, O-O = 34.9 kcal mol) while sulphur-sulphur double bond (S=S) is not so large (less than double of sulphur- sulphur single bond, S-S = 63.8 kcal mol). As a result, -O-O-O- chains are less stable as compared to O=O molecule while –S-S-S- chains are more stable than S=S molecule. Therefore, at room temperature, while oxygen exists as a diatomic gas molecule, Sulphur exists as S8 solid. (b) Metallic and non-metallic character: Metallic character increases with the increase of atomic number. Oxygen and Sulphur are distinctly non-metallic. Selenium and tellurium show both non-metallic and metallic characters but polonium is definitely a metal. (c) There is a gradual gradation in physical properties: (i) Atomic and ionic radii: The atomic radii of the elements of group 16 are smaller than those of the corresponding elements of group 15. The atomic radii of the elements of this group increase gradually on moving down the group. Element O S Se Te Po Atomic radii (pm) 66 104 117 137 146 The smaller atomic radii of group 16 elements as compared with corresponding elements of groups 15 are due to the increased effective nuclear charge with increase in greater attraction towards nucleus. The attraction brings contraction in size. The gradual increase on moving down the group is due to the increase in the number of electron shells from member to member. In the formation of anions, two electrons are being added to an atom. Therefore, the effective nuclear charge is reduced and hence, the electron closed expands. Thus, the negative ions (M²) are bigger in size than the corresponding atoms. Like atomic radii, ionic radii also increase gradually on moving down the group. Element O S Se Te Ionic radii (pm) 140 184 198 221 (ii) Ionization energy: The ionization enthalpies are high and thus the elements do not lose the electrons to form positive ions easily. The values decrease as the atomic number increases from O to Po and thus the tendency to form positive ion increases gradually, i.e., metallic nature increases. (iii) Electronegativity: Electronegativity decreases gradually. Oxygen is second most electronegative element after fluorine. This decrease indicates a change from non-metallic to metallic character. (iv) Electron affinity: Group 16 elements have high electron affinities. On moving from oxygen to Sulphur, the EA₁ value increases and then decreases from S to Po. The electron density in 2p energy shell in oxygen is high due to small size of oxygen atom and thus, there is some resistance to the incoming electron and thereby the EA of oxygen is comparatively low. Sulphur has maximum value and on moving from S to Po, the EA1 values decrease due to increased size. The elements show allotropy. (d) Allotropy: All the Element show Allotropic Element Allotropic forms Oxygen Ordinary oxygen and ozone Sulphur Rhombic, monoclinic, plastic, amorphous Selenium Red form (non-metallic),grey form (metallic form) Tellurium Crystalline and amorphous Polonium α and β forms (Both are metallic forms) (e)Catenation: Oxygen and Sulphur show the property of catenation. The property is more pronounced in Sulphur. H-O-O- H, H-S-S-H, H-S-S-S-H, H-S-S S-S-H, the peroxides and polysulphides are fairly stable. (f) Oxidation states: As the configuration of outer most shell is ns² np4, these elements try to gain or share two electrons in order to attain inert gas configuration. Oxygen being highly electronegative shows (-2) oxidation state in its compounds except in oxygen fluorides and most of the metal oxides are ionic and contain oxygen as di-negative anion. O²- Since the electronegativity decreases, the tendency to exhibit -2 oxidation state decreases as we go down in the group. However, positive oxidation states are exhibited by S, Se, Te and Po. In addition to +2 oxidation state, +4 and +6 oxidation states are observed. This is due to the availability of d-orbitals in these elements. Oxygen has no d orbitals and hence cannot show +4 and +6 oxidation states while Sulphur can have 2, 4 or 6 unpaired orbitals forming 2, 4 or 6 covalent bonds. Thus, oxygen is never more than divalent while other members may be divalent, tetravalent and hexavalent. The compounds having +4 oxidation state show both oxidizing and reducing properties while compounds having +6 oxidation states are only oxidizing. (g) Multiple bonding: Oxygen atom has the tendency to form multiple bonds (pπ-pπ interaction) with the other oxygen atom on account of small size and high electronegativity. However, the rest of the elements do not form pπ-pπ multiple bonds due to their large size. Sulphur and higher members of group 16 possess vacant d-orbitals in their valence shell. They use these orbitals to form dπ-pπ bonds. However, this tendency is maximum and stronger in Sulphur and decreases from Sulphur onwards. 3. Trends in Chemical Reactivity Oxygen is the most reactive element of the group despite it has high bond dissociation energy of oxygen molecule (493.4 kJ mol-¹) as nearly all its reactions are exothermic. Once initiated, these reactions continue spontaneously. Oxygen directly combines with almost all the metals except noble metals, all the non-metals except noble gases and halogens and many compounds under suitable conditions. The oxides are generally stable compounds. The elements such as Pt, halogens and noble gases which do not directly combine with oxygen, form compounds with oxygen indirectly. After oxygen, sulphur is quite reactive element especially at high temperatures which help in breaking of S-S bonds. Sulphur burns in air and reacts directly with carbon, phosphorus, arsenic and many metals. Oxidizing acids oxidize it into SO2 and alkalies dissolve it to give sulphides and thiosulphates. It reacts with H₂ and halogens. The sulphides are stable compounds and many metals are found in nature in the form of sulphides. However, the reactivity of group 16 elements decreases from oxygen to polonium O > S > Se> Te > Po Selenium and tellurium combine with highly electropositive elements such as alkali and alkaline earth metals. Se and Te combine with oxygen, fluorine and chlorine. In general, the compounds of selenium and tellurium are less stable than oxygen and Sulphur. Ozone O3 O3 is prepared by electrical discharge on O2, in which 10% concentration O3 can be produced, also O3 in small amounts can be formed by electrical analysis of dilution H2SO4 acid and also in some reactions that give the atomic oxygen. Lastly O3 is formed by the action of UV radiation on O2 in the higher layer of the atmosphere. The higher concentration of O3 is reached at 25 Km. above the ground, so that the earth can be protected from the excess UV radiation. The formation reaction of O3 from O2 endothermic, although it decomposes slowly at 250 ºc without catalyst or UV. Chemical Properties of O2 &O3 The chemical activity of O3 differs than that of O2, it is well known that O2 combines with most element but at high temperature, in the meantime O3 reacts at normal temperature with materials that O2 doesn’t react with e.g.: The ability of some materials for simultaneous oxidation in aqueous solution belongs to O2 dissolved, e.g.; Cr+2 ion doesn’t oxidize in pure water, while it is rapidly oxidize if water is saturated with O2, also Fe+2 oxidizes slowly in acidic medium but oxidizes rapidly in basic medium in presence of O2. The average of simultaneous oxidizing of many bio materials increases with presence of transition elements ions ( ascorbic acid in presence of Cu+2) in which Cu+2 reduced to Cu+ which simultaneous oxidizes to Cu+2 in presence of dissolved O2 and so on … Oxygen Compounds: - Hydrogen peroxide H2O2 Analysis process takes place at low temperature to prevent hydrolysis of the produced acid at the moment of formation as:-