AL Chemistry Resource Book Unit 6-EM PDF

G. C. E. (Advanced Level) CHEMISTRY Grade 12 Unit 6: Chemistry of s, p and d Block Elements Department of Science Faculty of Science and Technology National Institute of Education www.nie.lk...

G. C. E. (Advanced Level) CHEMISTRY Grade 12 Unit 6: Chemistry of s, p and d Block Elements Department of Science Faculty of Science and Technology National Institute of Education www.nie.lk I Chemistry Resource Book Grade 12 © National Institute of Education First Print – 2018 Department of Science Faculty of Science and Technology National Institute of Education Sri Lanka Published by: Press National Institute of Education Maharagama Sri Lanka II Message from the Director General The National Institute of Education takes opportune steps from time to time for the development of quality in education. Preparation of supplementary resource books for respective subjects is one such initiative. Supplementary resource books have been composed by a team of curriculum developers of the National Institute of Education, subject experts from the national universities and experienced teachers from the school system. Because these resource books have been written so that they are in line with the G. C. E. (A/L) new syllabus implemented in 2017, students can broaden their understanding of the subject matter by referring these books while teachers can refer them in order to plan more effective learning teaching activities. I wish to express my sincere gratitude to the staff members of the National Institute of Education and external subject experts who made their academic contribution to make this material available to you. Dr. (Mrs.) T. A. R. J. Gunasekara Director General National Institute of Education Maharagama. III Message from the Director Since 2017, a rationalized curriculum, which is an updated version of the previous curriculum is in effect for the G.C.E (A/L) in the general education system of Sri Lanka. In this new curriculum cycle, revisions were made in the subject content, mode of delivery and curricular materials of the G.C.E. (A/L) Physics, Chemistry and Biology. Several alterations in the learning teaching sequence were also made. A new Teachers’ Guide was introduced in place of the previous Teacher’s Instruction Manual. In concurrence to that, certain changes in the learning teaching methodology, evaluation and assessment are expected. The newly introduced Teachers’ Guide provides learning outcomes, a guideline for teachers to mould the learning events, assessment and evaluation. When implementing the previous curricula, the use of internationally recognized standard textbooks published in English was imperative for the Advanced Level science subjects. Due to the contradictions of facts related to the subject matter between different textbooks and inclusion of the content beyond the limits of the local curriculum, the usage of those books was not convenient for both teachers and students. This book comes to you as an attempt to overcome that issue. As this book is available in Sinhala, Tamil, and English, the book offers students an opportunity to refer the relevant subject content in their mother tongue as well as in English within the limits of the local curriculum. It also provides both students and teachers a source of reliable information expected by the curriculum instead of various information gathered from the other sources. This book authored by subject experts from the universities and experienced subject teachers is presented to you followed by the approval of the Academic Affairs Board and the Council of the National Institute of Education. Thus, it can be recommended as a material of high standard. Dr. A. D. A. De Silva Director Department of Science IV  Guidance Dr. (Mrs.) T. A. R. J. Gunasekara Director General National Institute of Education Supervision Dr. A. D. A. De Silva Director, Department of Science National Institute of Education Mr. R. S. J. P. Uduporuwa Former Director, Department of Science National Institute of Education Subject Leader Mrs. M. S. Wickramasinghe Assistant Lecturer, Department of Science National Institute of Education Internal Editorial Panel Mr. L. K. Waduge Senior Lecturer, Department of Science Mrs. G. G. P. S. Perera Assistant Lecturer, Department of Science Mr. V. Rajudevan Assistant Lecturer, Department of Science Writing Panel Dr. M.N. Kaumal - Senior Lecturer, Department of Chemistry, University of Colombo (Unit 6) External Editorial Panel Prof. S. P. Deraniyagala - Senior Professor, Department of Chemistry, University of Sri Jayewardenepura Prof. M. D. P. De Costa - Senior Professor, Department of Chemistry, University of Colombo Prof. H. M. D. N. Priyantha - Senior Professor, Department of Chemistry, University of Peradeniya Prof. Sudantha Liyanage - Dean, Faculty of Applied Sciences, University of Sri Jayewardenepura Mr. K. D. Bandula Kumara - Deputy Commissioner, Education Publication Department, Ministry of Education Mrs. Deepika Nethsinghe - SLTS-1 (Rtd), Ladies College, Colombo 07 Mrs. Muditha Athukorala - SLTS-1, Prajapathi Balika Vidyalaya, Horana Miss. C. A. N. Perera - SLTS-1, Princess of Wales’, Moratuwa Mrs. V. K. W. D. Salika Madavi - SLTS-1, Muslim Ladies College, Colombo 04 V Mrs. H.M.D.D. D. Manike - SLTS-1, Viharamhadevi Balika Vidyalaya, Kiribathgoda Mr. S. Thillainathan - SLTS-1 (Rtd), Hindu Ladies College, Colombo 06 Miss. S. Veluppillai - SLTS-1 (Rtd), Hindu Ladies College, Colombo 06 Mrs. M. Thirunavukarasu - SLTS-1 (Rtd), Hindu Ladies College, Colombo 06 Mrs. S. Rajadurai - SLTS-1 (Rtd), St. Peters' College, Colombo 04 Language Editing Dr. Chandra Amarasekara Consultant, National Institute of Education Mr. M. A. P. Munasinghe Chief Project Officer (Rtd.), National Institute of Education Cover Page Mrs. R. R. K. Pathirana Technical Assitant, National Institute of Education Supporting Staff Mrs.Padma Weerawardana Mr. Mangala Welipitiya Mr. Ranjith Dayawansa VI  Content Message from the Director General ……………………………………………………………..…… iii Message from the Director ………………………………………………………………………........ iv Subject Committee………………………………………………………………………………..…... v 4.0 Chemistry of s, p and d block elements……………………………………..…….. 121-173 s Block Elements 4.1 Group 1 elements………………………………………………………………….122 4.1.1 Group trends 4.1.2 Reactions of Group 1 elements 4.1.3 Thermal stability of salts 4.1.4 Solubility of Group 1 salts 4.1.5 Flame test 4.2 Group 2 elements………………………………………………………………….126 4.2.1 Group trends 4.2.2 Reactions of alkaline earth Group 2 elements 4.2.3 Thermal stability of salts 4.2.4 Solubility of Group 2 salts 4.2.5 Flame test p Block Elements 4.3 Group 13 elements…………………………………………………………………130 4.3.1 Group trends 4.3.2 Aluminium 4.4 Group 14 elements…………………………………………………………………132 4.4.1 Group trends 4.4.2 Diamond and graphite 4.4.3 Carbon monoxide and carbon dioxide 4.4.4 Oxoacid of carbon 4.5 Group 15 elements…………………………………………………………………135 4.5.1 Group trends 4.5.2 Chemistry of nitrogen 4.5.3 Oxoacids of nitrogen 4.5.4 Ammonia and ammonium salts 4.6 Group 16 elements…………………………………………………………………140 4.6.1 Group trends 4.6.2 Hydrides of Group 16 4.6.3 Oxygen 4.6.4 Sulphur 4.6.5 Oxygen containing compounds 4.6.6 Hydrogen peroxide 4.6.7 Sulphur containing compounds 4.6.8 Oxoacids of sulphur 4.7 Group 17 elements…………………………………………………………………147 4.7.1 Group trends 4.7.2 Simple compounds of Group 17 4.7.3 Reactions of chlorine VII 4.8 Group 18 elements…………………………………………………………………152 4.8.1 Group trends 4.8.2 Simple compounds of group 18 elements 4.9 Periodic trends shown by s and p block elements……………………………….153 4.9.1 The valence electron configuration 4.9.2 Metallic character 4.9.3 Reactions of third period oxides with water, acids and bases 4.9.4 Acid, base and amphoteric nature of hydroxides and hydrides 4.9.5 Nature of the halides across the third period d Block Elements 4.10 Transition elements………………………………………………………………158 4.10.1 Occurrence 4.10.2 Properties of fourth period d block elements 4.10.3 Oxides of d block oxides 4.10.4 Chemistry of some selected d block oxides 4.10.5 Coordination compounds of transition metal ions 4.10.6 Nomenclature of simple complex ions and compounds 4.10.7 Factors affecting the colour of the complexes 4.10.8 Importance of d block elements 4.10.9 Identification tests for selected cations of d block elements VIII G.C.E. (A/L) CHEMISTRY: UNIT 6 Chemistry of s, p and d block elements 4. Chemistry of s, p and d Block Elements Content s Block Elements 4.7 Group 17 elements 4.1 Group 1 elements 4.7.1 Group trends 4.1.1 Group trends 4.7.2 Simple compounds of Group 17 4.1.2 Reactions of Group 1 elements 4.7.3 Reactions of chlorine 4.1.3 Thermal stability of salts 4.1.4 Solubility of Group 1 salts 4.8 Group 18 elements 4.1.5 Flame test 4.8.1 Group trends 4.2 Group 2 elements 4.8.2 Simple compounds of group 18 4.2.1 Group trends elements 4.2.2 Reactions of alkaline earth Group 2 elements 4.9 Periodic trends shown by s and p 4.2.3 Thermal stability of salts block elements 4.2.4 Solubility of Group 2 salts 4.9.1 The valence electron configuration 4.2.5 Flame test 4.9.2 Metallic character p Block Elements 4.9.3 Reactions of third period oxides with 4.3 Group 13 elements water, acids and bases 4.3.1 Group trends 4.9.4 Acid, base and amphoteric nature of 4.3.2 Aluminium hydroxides and hydrides 4.9.5 Nature of the halides across the third 4.4 Group 14 elements period 4.4.1 Group trends 4.4.2 Diamond and graphite 4.4.3 Carbon monoxide and carbon dioxide d Block Elements 4.4.4 Oxoacid of carbon 4.10 Transition elements 4.10.1 Occurrence 4.5 Group 15 elements 4.10.2 Properties of fourth period d block 4.5.1 Group trends elements 4.5.2 Chemistry of nitrogen 4.10.3 Oxidation states of d block oxides 4.5.3 Oxoacids of nitrogen 4.10.4 Chemistry of some selected d block 4.5.4 Ammonia and ammonium salts oxides 4.6 Group 16 elements 4.10.5 Transition metal ions of coordination 4.6.1 Group trends complexes 4.6.2 Hydrides of Group 16 4.10.6 Nomenclature of simple complex ions 4.6.3 Oxygen and compounds 4.6.4 Sulphur 4.10.7 Factors affecting the colour of the 4.6.5 Oxygen containing compounds complexes 4.6.6 Hydrogen peroxide 4.6.7 Sulphur containing compounds 4.10.8 Importance of d block elements 4.6.8 Oxoacids of sulphur 4.10.9 Identification tests for selected cations 121  G.C.E. (A/L) CHEMISTRY: UNIT 6 Chemistry of s, p and d block elements Introduction This section describes the physical and chemical properties of elements in s, p and d blocks. This section will help to identify trends and patterns among elements in the periodic table. s Block Elements 4.1 Group 1 elements All Group 1 elements are metals except hydrogen which is a nonmetal. Unlike most other metals, they have low densities. All Group 1 elements have the valence shell electron configuration of ns1 therefore, they are highly reactive. Sodium can be found naturally as various salts such as NaCl (rock salt) and Na2B4O7·10H2O (borax). Some examples of naturally occurring potassium salts are KCl (sylvite) and KCl·MgCl2·6H2O (carnallite). 4.1.1 Group trends All alkali metals are lustrous. They are high electrical and thermal conductors. These metals are soft and become even softer when progress down the group. The melting point of Group 1 metals decreases down the group. The values given in Table 4.1 below can be used to understand the trends among these elements. Group 1 metals always show oxidation number of +1 when they form compounds. Most compounds are stable ionic solids. Table 4.1 Properties of Group 1 elements Li Na K Rb Cs Ground state electronic [He]2s1 [Ne]3s1 [Ar]4s1 [Kr]5s1 [Xe]6s1 configuration Metallic radius/ pm 152 186 231 244 262 Melting point/ °C 180 98 64 39 29 Radius of M+/ pm 60 95 133 148 169 1st ionization energy/ kJ mol-1 520 495 418 403 375 2nd ionization energy/ kJ mol-1 7298 4562 3052 2633 2234 Increase in the atomic radius from Li to Cs makes the ionization energy of these elements to decrease down the group, and this can be used to explain the chemical properties of Group 1 elements. Reactivity of the Group 1 elements increases down the group. ** When writing equations in inorganic chemistry, it is not always essential to indicate the physical state of reactants or products. However, always balance equations must be written to consider as a complete answer. 122  G.C.E. (A/L) CHEMISTRY: UNIT 6 Chemistry of s, p and d block elements 4.1.2 Reactions of Group 1 elements With oxygen (O2) 4M + O2 2M2O With excess oxygen (O2) Na form peroxides 2Na + O2 Na2O2 With excess oxygen (O2) K, Rb and Cs form M + O2 MO2 superoxides With nitrogen (N2) only Li forms stable 6Li + N2 2Li3N nitride With hydrogen (H2) 2M + H2 2MH With water (H2O) 2M + 2H2O 2MOH + H2 With acids (H+) 2M + 2H+ 2M+ + H2 Reaction with water Group 1 metals show an increase in reactivity with water down the group. The reactivity trend with water is as follows. Li Na K Rb Cs Vigourously Gently Vigorously Explosively Explosively with ignition Lithium reacts non-vigorously with water or with water vapour available in the air to produce lithium hydroxide and hydrogen gas. However, both sodium and potassium react vigorously with water to produce metal hydroxide and hydrogen gas. These reactions are highly exothermic except with Li. Reactions with oxygen/ air Lithium can react both with oxygen and nitrogen. When heated, lithium burns to produce lithium oxide (Li2O), a white powder. With nitrogen gas, lithium gives lithium nitride (Li3N). However, both sodium and potassium do not react with nitrogen gas. When sodium is burnt in air, sodium peroxide is mainly produced with some sodium oxide. In contrast, when potassium is burnt in air, potassium superoxide is formed as the main product with some potassium oxide and peroxide. Oxidation numbers of oxygen in sodium or potassium peroxide are -1 and in potassium superoxide, oxidation numbers are -1 and 0. Group 1 metal oxides react with water to produce metal hydroxides as shown below. Na2O(s) + H2O(l) 2NaOH(aq) 123  G.C.E. (A/L) CHEMISTRY: UNIT 6 Chemistry of s, p and d block elements When heated, lithium forms lithium nitride with nitrogen. Only lithium forms a stable alkali-metal nitride. With water, lithium nitride produces ammonia and lithium hydroxide. Li3N(s) + 3H2O(l) 3LiOH(aq) + NH3(g) Group 1 hydroxides react with carbon dioxide to produce relevant carbonates. These carbonates can further react with carbon dioxide to produce metal hydrogen carbonates. 2NaOH(aq) + CO2(g) Na2CO3(aq) + H2O(l) Na2CO3(aq) + CO2(g) + H2O(l) 2NaHCO3(s) Sodium hydrogen carbonate is less soluble than sodium carbonates in water. Reactions with hydrogen gas Group 1 elements react with hydrogen to produce solid, ionic metal hydrides. In these hydrides, hydrogen has the oxidation number of –1. These metal hydrides react vigorously with water to produce hydrogen gas. 2Na(s) + H2(g) 2NaH(s) NaH(s) + H2O(l) NaOH(aq) + H2(g) Reactions with acids Lithium, sodium and potassium react vigorously with dilute acids to produce hydrogen gas and relevant metal salts. These reactions are highly exothermic and explosive. A few selected reactions are shown below. 2Li(s) + dil. 2HNO3(aq) 2LiNO3(aq) + H2(g) 2Na(s) + dil. H2SO4(aq) Na2SO4(aq) + H2(g) 4.1.3 Thermal stability of salts Decomposition of nitrates Group 1 nitrates are used as fertilizers and explosives. These nitrates decompose upon heating. LiNO3 decomposes to produce lithium oxide, nitrogen dioxide and oxygen. However, the other Group 1 nitrates on heating produce relevant metal nitrite and oxygen. 4LiNO3(s) 2Li2O(s) + 4NO2(g) + O2(g) 2KNO3(s) 2KNO2(s) + O2(g) 124  G.C.E. (A/L) CHEMISTRY: UNIT 6 Chemistry of s, p and d block elements Decomposition of carbonates Carbonates are stable and they will melt before they decompose into oxides. However, Li2CO3 is less stable and decomposes readily. Li2CO3(s) Li2O(s) + CO2(g) Decomposition of bicarbonates Decomposition of bicarbonates of Group 1 is shown below. 2NaHCO3(s) Na2CO3(s) + H2O(g) + CO2(g) Thermal stability increases down the group. 4.1.4 Solubility of Group 1 salts All Group 1 salts are soluble in water except some lithium salts such as LiF, Li2CO3 and Li3PO4. All these salts are white solids unless the salt anion is a coloured ion. Solubility of Group 1 halides increase down the group is shown in Table 4.2. Table 4.2 The solubility of halides of sodium Salt Solubility/ mol L-1 NaF 0.99 NaCl 6.2 NaBr 9.2 NaI 12.3 Variation in the solubility can be understood using the energy cycle for the solvation of ionic solids. The solubility can be explained using Gibbs free energy. For almost all ionic solids of Group 1, are soluble in water due to the negative Gibbs free energy in the solvation process. 125  G.C.E. (A/L) CHEMISTRY: UNIT 6 Chemistry of s, p and d block elements Enthalpy and entropy cycles for the solvation process are shown below. Figure 4.1 Enthalpy and entropy cycles for the solvation process Using these two energy cycles, enthalpy and entropy change of solvation can be calculated and these calculated values are given in Table 4.3. Free energy is calculated using the equation, ΔGϴ = ΔHϴ - T ΔSϴ Table 4.3 Free energy change of salts during solvation Salt Enthalpy change/ Entropy change × T Free energy change/ kJ mol-1 (K × kJ mol-1 K-1) kJ mol-1 NaF +1 -2 +3 NaCl +4 +13 -9 NaBr -1 +18 -19 NaI -9 +23 -32 Calculated Gibbs free energies match with the solubility trend for the sodium halides. The free energy change gets more negative from sodium fluoride to sodium iodide. 4.1.5 Flame test The flame test can be used to identify alkali metals and their compounds. Flame colours of Group 1 metals and compounds are given below. Lithium – Crimson red Sodium – Yellow Rubidium – Red-violet Caesium – Blue - violet Potassium – Lilac 4.2 Group 2 elements Group 2 elements are known as alkaline earth metals. They are less reactive than Group 1 metals due to its valence shell ns2 electron configuration. Both calcium and magnesium can be found naturally in dolomite (CaCO3∙MgCO3). Magnesite (MgCO3), kieserite (MgSO4∙H2O) and carnallite (KMgCl3∙6H2O) are 126  G.C.E. (A/L) CHEMISTRY: UNIT 6 Chemistry of s, p and d block elements examples of minerals with magnesium. Fluoroapatite [3(Ca3(PO4)2)∙CaF2] and gypsum (CaSO4∙2H2O) are commercially important calcium contacting minerals. 4.2.1 Group trends Beryllium and magnesium are greyish metals and other Group 2 metals are soft and silvery in colour. Group 2 metal oxides produce basic oxides except for BeO which shows amphoteric properties. Beryllium behaves similar to Al and this can be understood using the diagonal relationship between Al and Be in the periodic table. Elements of Group 2 have higher densities and stronger metallic bonds compared to the Group 1 metals. This is due to the availability of a greater number of electrons to form a stronger metallic bond and their smaller size in atomic radii. The first ionization energies of Group 2 elements are higher than that of Group 1 elements due to their electron configuration of ns2. Elements become more reactive and produce +2 oxidation state easily down the group. The properties of Group 2 elements are given in Table 4.4. Table 4.4 Properties of Group 2 elements Be Mg Ca Sr Ba Ground state electronic [He]2s2 [Ne]3s2 [Ar]4s2 [Kr]5s2 [Xe]6s2 configuration Metallic radius/ pm 112 160 197 215 224 Melting point/ °C 1560 923 1115 1040 973 Radius of M2+/ pm 30 65 99 113 135 1st ionization energy/ 899 736 589 594 502 kJ mol-1 2nd ionization energy/ 1757 1451 1145 1064 965 kJ mol-1 3rd ionization energy/ 14850 7733 4912 4138 3619 kJ mol-1 127  G.C.E. (A/L) CHEMISTRY: UNIT 6 Chemistry of s, p and d block elements 4.2.2 Reactions of alkaline earth Group 2 elements With oxygen (O2) 2M + O2 2MO With excess oxygen (O2) Ba forms its Ba + O2 BaO2 peroxide With nitrogen (N2), at high 3M + N2 M3N2 temperatures With water (H2O(l)), at room M + 2H2O M(OH)2 + H2 temperature (e.g.: Ca, Sr and Ba) With hot water (H2O(l)) Mg + 2H2O Mg(OH)2 + H2 (e.g.: Mg reacts slow) With steam (H2O(g)) Mg + H2O MgO + H2 With acids (H+) M + 2H+ M2+ + H2 With hydrogen (H2), at high M + H2 MH2 temperatures with Ca, Sr, Ba at high pressure with Mg With concentrated acids Mg + 2H2SO4 MgSO4 + SO2 + 2H2O Mg + 4HNO3 Mg(NO3)2+2NO2+2H2O Reaction with water Beryllium does not react with water, but it reacts with steam. The reaction of magnesium with water at room temperature is negligible. However, magnesium reacts slowly with hot water. Calcium, strontium and barium react readily with cold water. Reaction with water produces metal hydroxide and hydrogen gas. Ca(s) + 2H2O(l) Ca(OH)2(aq) + H2(g) Reactions with hydrogen All Group 2 elements, except Be, react with hydrogen to produce metal hydrides which are ionic solids. In these hydrides, hydrogen has an oxidation number of – 1. These metal hydrides (not violent as Group 1) react vigorously with water to produce hydrogen gas. Ca(s) + H2(g) CaH2(s) CaH2(s) + 2H2O(l) Ca(OH)2(aq) + 2H2(g) Reaction with nitrogen All Group 2 elements burns in nitrogen to form M3N2, nitrides. These nitrides react with water to produce ammonia in the same way as lithium does. 3Mg (s) + N2(g) Mg3N2(s) Mg3N2(s) + 6H2O(l) 3Mg(OH)2(aq) + 2NH3(g) 128  G.C.E. (A/L) CHEMISTRY: UNIT 6 Chemistry of s, p and d block elements 4.2.3 Thermal stability of salts Decomposition of nitrates Upon heating, Group 2 nitrates behave much similar to lithium nitrate. Group 2 nitrates decompose to produce metal oxide, nitrogen dioxide and oxygen. All Group 2 nitrates are soluble in water. 2Mg(NO3)2(s) 2MgO(s) + 4NO2(g) + O2(g) Decomposition of carbonates Thermal stability of these carbonates increases down the group. Thermal stability of these carbonates increases with the size of the cation. The polarizing power of the cation decreases down the group due to the decrease of charge density of the cation. Carbonate anion attached to Mg2+ cation is highly polarized than that of carbonate attached to Ba2+. Highly polarized carbonate anion can undergo decomposition easily and this explains the lower decomposition temperature of MgCO3 than that of BaCO3.The general decomposition of metal carbonates is shown below. MCO3 MO + CO2 Decomposition temperature increases from 540 °C for MgCO3 to 1360 °C for BaCO3. Decomposition of bicarbonates Group 2 hydrogen carbonates are only stable in aqueous solutions and solid Group 2 hydrogen carbonates are not stable at room temperature. Ca(HCO3)2(aq) CaCO3(s) + CO2(g) + H2O(l) 4.2.4 Solubility of Group 2 salts Solubility of Group 2 changes depending on the compound. Some compounds such as nitrate, nitrite, halides, hydroxides, sulphides, bicarbonates all are soluble in water. The solubility varies down the group for certain compounds such as hydroxides, sulphate, sulphite, carbonate, phosphate and oxalate showing the patterns given in Table 4.5. Salts of Group 2 metals with uninegative anions, such as chloride and nitrates are generally soluble. However, salts formed with anions containing more than one negative charge, such as carbonates and phosphates, are insoluble. All carbonates are insoluble except BeCO3. Hydrogen carbonates are more soluble than carbonates. The solubility of Group 2 sulphates changes from soluble to insoluble when comparing solubility from MgSO4 to BaSO4. On the other hand, hydroxides change solubility from insoluble to soluble when moving down the group. For example, Mg(OH)2 is sparingly soluble whereas Ba(OH)2 is soluble and produces a strongly basic solution. 129  G.C.E. (A/L) CHEMISTRY: UNIT 6 Chemistry of s, p and d block elements Table 4.5 Solubility of Group 1 and 2 compounds Na+ K+ Mg2+ Ca2+ Sr2+ Ba2+ Cl- aq aq aq aq aq aq Br- aq aq aq aq aq aq - I aq aq aq aq aq aq - OH aq aq IS SS SS aq CO32- aq aq IS IS IS IS HCO3- aq aq aq aq aq aq NO2- aq aq aq aq aq aq - NO3 aq aq aq aq aq aq S2- aq aq aq aq aq aq SO32- aq aq SS IS IS IS 2- SO4 aq aq aq SS IS IS PO43- aq aq IS IS IS IS CrO42- aq aq aq aq IS IS C2O42- aq aq SS IS IS SS aq – soluble, IS – insoluble, SS – sparingly soluble 4.2.5 Flame test Alkaline earth metals and compounds produce characteristic colours with the flame, and the flame test can be used to identify these elements using the flame colors shown below. Calcium – Orange-red Strontium – Crimson red Barium – Yellowish-green 130  G.C.E. (A/L) CHEMISTRY: UNIT 6 Chemistry of s, p and d block elements p Block Elements 4.3 Group 13 elements 4.3.1 Group trends Boron is a metalloid, and most of the boron compounds are covalent. However, aluminium is a metal with amphoteric properties. Gallium, indium and thallium are metals. The first member, B, of Group 13 is different from the other members due to its smaller atomic radius. Boron shows a strong diagonal relationship with Si in Group 14. All elements in Group 13 produce +3 oxidation state. The properties of Group 13 elements are given in Table 4.6. Table 4.6 Properties of Group 13 elements **B Al **Ga **In **Tl Ground state electronic [He]2s22p1 [Ne]3s23p1 [Ar]3d104s24p1 [Kr]4d105s25p1 [Xe]4f145d106s26p1 configuration Metallic radius/ pm - 143 153 167 171 Covalent radius/ pm 88 130 122 150 155 Melting point/ °C 2300 660 30 157 304 3+ Radius of M / pm 27 53 62 80 89 st 1 ionization energy/ 799 577 577 556 590 kJ mol-1 2nd ionization energy/ 2427 1817 1979 1821 1971 kJ mol-1 3rd ionization energy/ 3660 2745 2963 2704 2878 kJ mol-1 **Not a part of current G. C. E. (A/L) syllabus 4.3.2 Aluminium Aluminium is the third most abundant element in the earth crust. The exposed surface of aluminium produces a layer of Al2O3. This layer makes aluminium resistant to further reactions with oxygen. Due to this impermeable layer, Al can be considered as a non- reactive element with air. Reactions of aluminium Aluminium reacts readily with O2 and halogens. Also, it reacts with N2. With oxygen (O2) : 4Al + O2 2Al2O3 With halogen (X2): 2Al + 3X2 2AlX3 131  G.C.E. (A/L) CHEMISTRY: UNIT 6 Chemistry of s, p and d block elements With nitrogen (N2): 2Al + N2 2AlN Aluminum is less reactive than Groups 1 and 2 elements. Similar to beryllium, aluminium reacts with both acids and bases. The equations for the reactions of Al with acids and bases are given below. 2Al(s) + 6HCl(aq) 2AlCl3(aq) + 3H2(g) 2Al(s) + 2OH ̅ (aq) + 6H2O(l) 2[Al(OH)4] ̅(aq) + 3H2(g) Aluminium ion in aqueous solution is expected to be present as hexaaquaaluminium ion. However, hydrolysis of Al3+ produces [Al(OH2)5(OH)]2+(pentaaquahydroxidoaluminium ion) and then produces [Al(OH2)4(OH)2]+ (tetraaquadihydroxidoaluminium ion) as shown below. [Al(OH2)6]3+(aq) + H2O(l) [Al(OH2)5(OH)]2+(aq) + H3O+(aq) [Al(OH2)5(OH)]2+(aq) + H2O(l) [Al(OH2)4(OH)2]+(aq) + H3O+(aq) Addition of OH- ions to aluminum ions first produces a gelatinous precipitate of aluminum hydroxide. With excess OH- ions, the precipitated aluminum hydroxide is converted to tetrahydroxidoaluminate complex ion. Al3+(aq) + 3OH ̅ (aq) Al(OH)3(s) (white gelatinous ppt) Al(OH)3(s) + OH ̅ (aq) [Al(OH)4] ̅ (aq) or AlO2 ̅ (aq) + 2H2O(l) Group 13 elements can have six electrons in their valence shell by forming three covalent bonds due to their ns2np1 electron configuration. As a result, many of the Group 13 covalent compounds have an incomplete octet, so can act as Lewis acids to accept a pair of electrons from a donor. These compounds with incomplete octet are called electron deficient compounds. Both B and Al compounds with incomplete octet form dimers in the gaseous phase to satisfy the octet rule (Figure 4.2). Cl Cl Cl Al Al Cl Cl Cl Figure 4.2 Structure of gaseous Al2Cl6 132  G.C.E. (A/L) CHEMISTRY: UNIT 6 Chemistry of s, p and d block elements 4.4 Group 14 elements 4.4.1 Group trends Due to the formation of covalent bond network structure, the first three elements of group 14 have high melting points. Carbon is a nonmetal, whereas silicon and germanium are metalloids. Last two elements in the group, tin and lead are metals. Carbon can be found in nature mainly in coal, crude oil, calcite (CaCO3), CO2 in air, magnesite (MgCO3) and dolomite (CaCO3·MgCO3). Graphite, diamond and fullerenes are the allotropic forms of carbon. Fullerenes are recently found, and most well-known fullerene is C60, buckminsterfullerene (or bucky-ball). Carbon is the basis of life and the most important element in organic chemistry. Silicon and germanium are mainly used in the semiconductor industries. In addition, silicon is heavily used in inorganic polymer industry. The properties of Group 14 elements are given in Table 4.7. Table 4.7 Properties of Group 14 elements C **Si **Ge **Sn **Pb Ground state electronic [He]2s22p2 [Ne]3s23p2 [Ar]3d104s24p2 [Kr]3d105s25p2 [Xe]4f145d106s26p2 configuration Metallic radius/ - - - 158 175 pm Covalent radius/ 77 118 122 140 154 pm Melting point/ °C 3730 1410 937 232 327 Radius of M /4+ - - 53 69 78 pm **Not a part of current G. C. E. (A/L) Chemistry syllabus 4.4.2 Diamond and graphite Diamond and graphite are composed of homoatomic (same atoms) lattice structures. Diamond (sp3 hybridized carbon, tetrahedral) has a cubic crystalline structure. Graphite (sp2 hybridized carbon, trigonal planar) has stacked two-dimensional carbon layers. Carbon-carbon bonds in graphite are shorter than that of diamond (diamond 154 pm and graphite 141 pm) due to the hybridization of carbon atoms. These two crystalline lattice structures are hard however diamond structure is the strongest lattice. Graphite is an electrical and a thermal conductor due to delocalizing π electrons (Figure 4.3). Interactions between layers of carbon in graphite are weak and this makes graphite a good lubricant. 133  G.C.E. (A/L) CHEMISTRY: UNIT 6 Chemistry of s, p and d block elements Figure 4.3 Delocalizing π bonds of graphite Fullerenes are another series of carbon allotropes. In fullerenes, carbon atoms are connected in a spherical manner. Structures of graphite, diamond and fullerene (C60) are shown in Figure 4.4. (a) (b) (c) Figure 4.4 Structures of (a) graphite, (b) diamond and (c) fullerene (C60) 4.4.3 Carbon monoxide and carbon dioxide Carbon monoxide is a colourless, odourless, highly poisonous gas. Bond enthalpy of carbon monoxide is more than that of the C=O double bond. In carbon monoxide, CO bond length is shorter than that of a typical C=O double bond. This suggests that the bonding between C and O in carbon monoxide is not a typical C=O double bond. It has a triple bond nature between the two atoms of C and O. The Lewis structure of CO is shown in Figure 4.5. Figure 4.5 The Lewis structure of CO 134  G.C.E. (A/L) CHEMISTRY: UNIT 6 Chemistry of s, p and d block elements Carbon monoxide is mostly used as a reducing agent in the production of iron. Also, CO plays an important role in many catalytic reactions as a ligand due to the lone pair of electrons on the C atom. Carbon dioxide (Figure 4.6) solidifies due to London forces at low temperatures and/ or under high pressures. Solid CO2 (dry ice) sublimes to produce gaseous carbon dioxide under normal atmosperic conditions. It is commonly used as a freezing agent in the food industry and to produce artificial rain. Figure 4.6 The Lewis structure of CO2 4.4.4 Oxoacid of carbon Oxoacid of carbon is referred to as carbonic acid (H2CO3) which is a weak acid. The bond structure of H2CO3 is given in Figure 4.7. Carbonic acid can be prepared by dissolving CO2 in water under pressure. CO2(aq) + H2O(l) H2CO3(aq) H2CO3(aq) + H2O(l) HCO3-(aq) + H3O+(aq) HCO3-(aq) + H2O(l) CO32-(aq) + H3O+(aq) Figure 4.7 The bond structure of H2CO3 Hydrogen atom which is directly connected to oxygen atom can be released as a proton to the solution by exhibiting the acidic property of carbonic acid. Carbon dioxide reacts with bases to produce carbonates showing its acidic property. In the presence of excess CO2 thus formed carbonates of Group 1 and 2 produce hydrogen carbonates. CO2(g) + 2NaOH(aq) Na2CO3(aq) + H2O(l) Na2CO3(aq) + excess CO2(g) + H2O(l) 2NaHCO3(aq) 135  G.C.E. (A/L) CHEMISTRY: UNIT 6 Chemistry of s, p and d block elements 4.5 Group 15 elements 4.5.1 Group trends The first element, nitrogen of Group 15 shows different properties from the other elements in this group (Table 4.8). Metallic character of the Group 15 elements increases down the group. Nitrogen and phosphorous are nonmetals and show oxidation numbers -3 to +5. Nitrogen can achieve +5 oxidation state with oxygen and fluorine. Dinitrogen, N2 is greatly stable (inert) under normal conditions due to strong triple bond (942 kJ mol-1). Except nitrogen, all the other elements exist as solids. The higher electronegativity, the smaller atomic radius and the absence of d orbitals make nitrogen different from the other elements in the group. Table 4.8 Properties of Group 15 elements N **P **As **Sb **Bi Ground state electronic [He]2s22p3 [Ne]3s23p3 [Ar]3d104s24p3 [Kr]3d105s25p3 [Xe]4f145d106s26p3 configuration Metallic radius/ pm - - - - 182 Covalent radius/ pm 75 110 122 143 152 Melting point/ °C -210 44 (white) 613 630 271 590 (red) Pauling 3.0 2.2 2.2 2.0 2.0 electronegativity **Not relevant to the current G. C. E. (A/L) Chemistry syllabus 4.5.2 Chemistry of nitrogen Nitrogen (boiling point is 195.8 °C) is slightly soluble in water under atmospheric pressure, but the solubility greatly increases with pressure. Nitrogen does not form allotropes. Dinitrogen shows only a few reactions and one of them is given below. 3Mg(s) + N2(g) Mg3N2(s) Since nitrogen is an inert gas its chemical reactions occur under strong conditions. For an instance nitrogen gas reacts with oxygen in the presence of external energy from an elctrical spark. This reaction naturally occurs in lightening. N2(g) + 2O2(g) 2NO2(g) Nitrogen shows oxidation states from –3 to +5. Compounds with these oxidation states are shown in Table 4.9. 136  G.C.E. (A/L) CHEMISTRY: UNIT 6 Chemistry of s, p and d block elements Table 4.9 Oxidation states of nitrogen Oxidation Compound Formula Bond structure state -3 Ammonia NH3 -2 Hydrazine N2H4 -1 Hydroxylamine NH2OH 0 Dinitrogen N2 +1 Dinitrogen N2O monoxide +2 Nitrogen NO monoxide +3 Dinitrogen N2O3 trioxide +4 Nitrogen NO2 dioxide +4 Dinitrogen N2O4 tetroxide +5 Nitric acid HNO3 +5 Dinitrogen N2O5 pentoxide 137  G.C.E. (A/L) CHEMISTRY: UNIT 6 Chemistry of s, p and d block elements 4.5.3 Oxoacids of nitrogen Nitrous acid which is unstable under normal atmospheric conditions is a weak acid. The bond structure of nitrous acid is given in Figure 4.8. Figure 4.8 The bond structure of nitrous acid Nitrous acid can undergo disproportionation to produce nitric acid and nitrogen monoxide which is a colourless gas. 3HNO2(aq) HNO3(aq) + 2NO(g) + H2O(l) Futher reaction of nitrogen monoxide with oxygen forms nitrogen dioxide which is redish brown in colour. 2NO(g) + O2(g) 2NO2(g) Nitric acid (Figure 4.9) is an oily and hazardous liquid. This acid is a strong oxidizing agent and can undergo vigorous chemical reactions. Figure 4.9 The bond structure of nitric acid Due to the light-induced decomposition, nitric acid produces oxygen and nitrogen dioxide. hν 4HNO3(aq) 4NO2(g) + O2(g) + 2H2O(l) Due to this reason concentrated nitric acid is stored in brown colour glass bottles in laboratories. Oxidizing and reducing reactions of nitric acid Dilute nitric acid reacts with metals to produce metal nitrate and hydrogen gas. In these reactions nitric acid acts as an oxidizing agent with respect to hydrogen. When magnesium and copper reacts with concentrated nitric acid it acts as an oxidizing agent with respect to nitrogen. 138  G.C.E. (A/L) CHEMISTRY: UNIT 6 Chemistry of s, p and d block elements Mg(s) + dil. 2HNO3(aq) Mg(NO3)2(aq) + H2(g) Mg(s) + conc. 4HNO3(l) Mg(NO3)2(aq) + 2NO2(g) + 2H2O(l) 3Cu(s) + dil. 8HNO3(aq) 3Cu(NO3)2(aq) + 2NO(g) + 4H2O(l) Cu(s) + conc. 4HNO3(l) Cu(NO3)2(aq) + 2NO2(g) + 2H2O(l) The reactions of conc. HNO3 acting as an oxidizing agent with non metals such as carbon and sulphur are given below. C(s) + conc. 4HNO3(l) CO2(g) + 4NO2(g) + 2H2O(l) S(s) + conc. 6HNO3(l) H2SO4(l) + 6NO2(g) + 2H2O(l) 4.5.4 Ammonia and ammonium salts Ammonia is a colourless gas with a strong characteristic smell. Ammonia is a basic gas which is readily soluble in water. NH3(g) + H2O(l) NH4OH(aq) Ammonium hydroxide is a weak base and partially dissociates to produce ammonium ions and hydroxide ions. NH4OH(aq) NH4+(aq) + OH ̅ (aq) Like any other base it reacts with dilute acids to produce aqueous salts. 2NH4OH(aq) + dil. H2SO4(aq) (NH4)2SO4(aq) + 2H2O(l) Hydrolysis of the ammonium ion in aqueous solution produces the conjugate base, ammonia. NH4+(aq) + H2O(l) NH3(aq) + H3O+(l) All amonium salts reacts with alkali to liberate amonia. NH4Cl(aq) + NaOH(aq) NaCl(aq) + NH3 (g) + H2O(l) Reactions of ammonia Ammonia acts as a reducing agent with chlorine, and the products vary with the amount of ammonia and chlorine used. In the presence of excess ammonia, chlorine produces nitrogen gas as one of the products. However, with excess chlorine, nitrogen trichloride is produced as one of the products, which is used for water disinfection. excess ammonia, 2NH3(g) + 3Cl2(g) N2(g) + 6HCl(g) 139  G.C.E. (A/L) CHEMISTRY: UNIT 6 Chemistry of s, p and d block elements The formed HCl further reacts with unreacted ammonia to form NH4Cl. excess chlorine, 3Cl2(g) + NH3(g) 3HCl(g) + NCl3(l) Nitrogen trichloride is a covalent chloride. It reacts with water to produce ammonia and hypochlorous acid. Due to the ability to produce hypochlorous acid, nitrogen trichloride is used as a water disinfecting agent. NCl3(l) + 3H2O(l) NH3(g) + 3HOCl(aq) Gaseous ammonia reacts with hydrogen chloride to produce a white smoke of solid ammonium chloride. This can be used as a confirmation test for ammonia. NH3(g) + HCl(g) NH4Cl(s) Ammonia acts as a weak reducing agent with CuO and Cl2. 3CuO(s) + 2NH3(g) N2(g) + 3Cu(s) + 3H2O(g) 2NH3(g) + 3Cl2(g) N2(g) + 6HCl(g) Ammonia can act as an oxidizing agent as well as an acid with metals under dry condition. 2Na(s) + 2NH3(l) 2NaNH2(l) + H2(g) 3Mg(s) + 2NH3(l) Mg3N2(l) + 3H2(g) Thermal decomposition of ammonium salts Some ammonium salts decompose upon heating to ammonium gas and to the acidic gas. (NH4)2CO3(s) 2NH3(g) + CO2(g) + H2O(g) NH4Cl(s) NH3(g) + HCl(g) (NH4)2SO4(s) NH3(g) + H2SO4(g)* *Prodcts of this reaction can vary with conditions. However, anions in some ammonium salts can oxidize the ammonium ion to produce many products upon heating. NH4NO2(s) N2(g) + 2H2O(g) NH4NO3(s) N2O(g) + 2H2O(g) (NH4)2Cr2O7(s) N2(g) + Cr2O3(s) + 4H2O(g) 140  G.C.E. (A/L) CHEMISTRY: UNIT 6 Chemistry of s, p and d block elements Identification of ammonium salts All ammonium salts produce ammonium gas with NaOH on warming. This gas produces white fumes of ammonium chloride when a glass rod moistened with concentrated hydrochloric acid. NH4+(aq) + OH¯(aq) NH3(g) + H2O(l) NH3(g) + HCl(g) NH4Cl(s) white fumes Reactions of nitrate Reaction of nitrate with iron(II)/ conc. sulphuric acid can be used to identify nitrate ion. This test is known as brown ring test. The brown coloured [Fe(NO)]2+ ring formed in the test tube, confirms the presence of nitrate. 2NO3ˉ (aq) + 4H2SO4(l) + 6Fe2+(aq) 6Fe3+(aq) + 2NO(g) + 4SO42ˉ (aq) + 4 H2O(l) Fe2+(aq)+ NO(g) [Fe(NO)]2+(aq) brown colour Nitrate reacts with Al/ NaOH to produce ammonia. 3NO3ˉ (aq) + 8Al(s) + 5OHˉ(aq) + 18H2O(l) 3NH3(g) + 8[Al(OH)4]ˉ(aq) 4.6 Group 16 elements 4.6.1 Group trends First element, oxygen of Group 16 shows different properties to the other elements in the group. Metallic nature increases going down the group. However, none of the Group 16 elements behaves as true metals. Both oxygen and sulphur are non-metals and other elements in the group show metallic and nonmetallic properties. Only oxygen exists as a gas, and other elements in the group are solids. Except for oxygen, other elements in the group can form even-numbered oxidation states from +6 to -2. Stability of +6 and -2 oxidation states decreases down the group whereas the stability of the +4 oxidation state increases. 141  G.C.E. (A/L) CHEMISTRY: UNIT 6 Chemistry of s, p and d block elements Table 4.10 Properties of Group 16 elements O S **Se **Te **Po 2 2 10 10 Ground state [He]2s [Ne]3s [Ar]3d [Kr]4d [Xe]4f 5d106s2 14 electronic 2p4 3p4 4s24p4 5s25p4 6p4 configuration Ionic radius X2-/ pm 140 184 198 221 - Covalent radius/ pm 73 103 117 137 140 Melting point/ °C -218 113(α) 217 450 254 Pauling 3.4 2.6 2.6 2.1 2.0 electronegativity 1st electron gain -141 -200 -195 -190 -183 enthalpy/ kJ mol-1 X(g) + e X-(g) 2nd electron gain 844 532 - - - enthalpy/ kJ mol-1 X-(g) + e X2-(g) ** Not a part of the current G. C. E. (A/L) Chemistry syllabus 4.6.2 Hydrides of Group 16 Group 16 elements form simple hydrides with hydrogen. All of them are covalent hydrides. The variation of selected properties down the group of hydrides are shown in Table 4.11. Table 4.11 Selected properties of Group 16 hydrides H2 O H2 S H2Se H2Te Melting point/ °C 0.0 -85.6 -65.7 -51 Boiling point / °C 100.0 -60.3 -41.3 -4 Bond length/ pm 96 134 146 169 Bond angle/ ° 104.5 92.1 91 90 Due to the extensive hydrogen bonding, H2O shows abnormally high boiling and melting points than the other hydrides of the group. Water is the only non-poisonous hydride among all the other hydrides of the group. The observed varation in bond length of covalent hydrides is due to the increase of size of the central atom. Therefore, bond length increases down the group. The covalent bond angle decreases as you come down in the group due to the less repulsion of the bonding electrons as a result of electronegativity of the central atom 142  G.C.E. (A/L) CHEMISTRY: UNIT 6 Chemistry of s, p and d block elements decreases down the group. In H2S, H2Se and H2Te the bond angles become close to 90°. This may also suggest that almost pure p orbitals on selenium and tellurium especially are used for binding with hydrogen. 4.6.3 Oxygen Oxygen has two allotropes, dioxygen (O2) and trioxygen (ozone, O3). Dioxygen is a colourless and an odourless gas which is slightly soluble in water. Ozone has a pungent odour. Ozone has a bond angle of 111.5°. Structure of these two allotropes are shown below. (a) (b) Figure 4.10 Structure of oxygen and ozone Catalytic decomposition of potassium chlorate and hydrogen peroxide can be used to produce oxygen. 2KClO3(s) 2KCl(s) + 3O2(g), heating in the presence of MnO2 or Pt 2H2O2(aq) 2H2O(l) + O2(g), heating in the presence of MnO2 Metals react with dioxygen to produce metal oxides. Ozone is a powerful oxidizing agent stronger than dioxygen. Ozone is used to disinfect water in many developed countries to kill pathogens. Unlike chlorine, ozone does not produce any harmful byproducts in the disinfection process. 4.6.4 Sulphur Sulphur can be classified as it is explained below. Sulphur Crystalline Amorphous Rhombic Monoclinic Plastic Colloidal Milk of (α-Sulphur) (β –Sulphur) Sulphur Figure 4.11 Classification of sulphur Unlike oxygen, sulphur forms single bonds with itself rather than double bonds. The most commonly occurring allotrope is rhombic sulphur which is referred to as α-sulphur (α- 143  G.C.E. (A/L) CHEMISTRY: UNIT 6 Chemistry of s, p and d block elements S8). It has a crown shape with eight-membered ring that has a cyclic zigzag arrangement as shown below. When heated above 93 °C, α-S8 changes its packing arrangement to the other commonly found form of monoclinic sulphur, β-sulphur (β-S8). These two forms are allotropes of each other. (a) (b) (c) Figure 4.12 (a) crown form of S8 (b) Rhombic sulphur (c) Monoclinic sulphur Crystalline form of rhombic and monoclinic sulphur consist of S8 rings in the shape of crown. These can be packed together in two different ways to form rhombic crystals and to form needle shaped monoclinic crystals as shown above. Below 95 °C the rhombic form is the most stable allotropic form of sulphur. Amorphous sulphur is an elastic form of sulphur which is obtained by pouring melted sulphur into water. Sudden cooling of molten sulphur with open chains converts liquid sulphur to amorphous sulphur with open chains. With time, amorphous sulphur converts to crystalline sulphur. The amorphous form of sulphur is malleable but it is unstable. 4.6.5 Oxygen containing compounds Water and hydrogen peroxide Structures of H2O and gaseous H2O2 are shown in the figures below. (a) (b) Figure 4.13 Structures of (a) H2O and (b) H2O2 Water is the most widely used solvent. Water ionizes as follows. This is reffered to as self-ionization of water. 144  G.C.E. (A/L) CHEMISTRY: UNIT 6 Chemistry of s, p and d block elements 2H2O(l) H3O+(aq) + OH-(aq) An amphiprotic molecule can either donate or accept a proton. Therefore, it can act as an acid or a base. Water is an amphiprotic compound since it has the ability to accept and release a proton. The amphoteric nature of water is shown below H2O(l) + HCl(aq) H3O+(aq) + Cl-(aq) H2O(l) + NH3(aq) NH4+(aq) + OH-(aq) 4.6.6 Hydrogen peroxide Hydrogen peroxide (H2O2) is a nonplanar molecule. The H2O2 molecule contains two OH groups which do not lie in the same plane and have a bent molecular shape with the bond angle in the gaseous phase for H-O-O as 94.8°. The structure shown in Figure 4.13 is the one that reduces with a minimum repulsion between the lone pairs found on the ‘O’ atoms. The two H-O groups have a dihedral angle of 111.5° between each other as indicated above in Figure 4.13. Due to the extensive hydrogen bonding, H2O2 is a viscous liquid H2O2 can act as an oxidizing as well as a reducing agent. It oxidizes to oxygen and reduces to water. Reducing half-reaction; H2O2(aq) + 2H+(aq) + 2e 2H2O(l) Oxidizing half-reaction; H2O2(aq) 2H+(aq) + O2(g) + 2e Disproportionation; 2H2O2(aq) O2(g) + 2H2O(l) Reactions of H2O2 H2O2 as an oxidizing agent; H2O2(aq) + 2H+(aq) + 2I-(aq) I2(aq) + 2H2O(l) H2O2(aq) + 2H+(aq) + 2Fe2+(aq) 2Fe3+(aq) + 2H2O(l) H2O2 as a reducing agent; 2MnO4-(aq) + 5H2O2(aq) + 6H+(aq) 2Mn2+(aq) + 5O2(g) + 8H2O(l) Cr2O72-(aq) + 3H2O2(aq) + 8H+(aq) 2Cr3+(aq) + 7H2O(l) + 3O2(g) 145  G.C.E. (A/L) CHEMISTRY: UNIT 6 Chemistry of s, p and d block elements 4.6.7 Sulphur containing compounds Hydrogen sulphide Hydrogen sulphide, H2S is a colourless, toxic and acidic gas with the odour of “rotten eggs”. H2S can be produced by reacting metal sulphides with strong acids. It dissolves in water to produce weak acidic solutions. Reactions of hydrogen sulphide H2S as an acid with strong bases; NaOH(aq) + excess H2S(g) NaHS(s) + H2O(l) 2NaOH(aq) + limited H2S(g) Na2S(s) + 2H2O(l) H2S reacts with metals as an acid as well as an oxidizing agent; 2Na(s) + excess 2H2S(g) 2NaHS(s) + H2(g) 2Na(s) + limited H2S(g) Na2S(s) + H2(g) Mg(s) + H2S(g) MgS(s) + H2(g) H2S as a reducing agent; 2KMnO4(aq)+3H2SO4(aq)+5H2S(g) K2SO4(aq)+5S(s)+2MnSO4(aq)+8H2O(l) K2Cr2O7(aq)+4H2SO4(aq)+3H2S(g) K2SO4(aq)+Cr2(SO4)3(aq)+3S(s)+7H2O(l) 2H2S(aq) + SO2(g) 3S(s) + 2H2O(l) Sulphur dioxide Sulphur dioxide is a colourless gas and soluble in water. Sulphur dioxide can act as an oxidizing and a reducing agent. Reactions of sulphur dioxide As an oxidizing agent; 2Mg(s) + SO2(g) 2MgO(s) + S(s) 3Mg(s) + SO2(g) 2MgO(s) + MgS(s) As a reducing agent; 5SO2(g) + 2KMnO4(aq) + 2H2O(l) K2SO4(aq) + 2MnSO4(aq) + 2H2SO4(aq) 3SO2(g) + K2Cr2O7(aq) + H2SO4(aq) K2SO4(aq) + Cr2(SO4)3 (aq)+ H2O(l) 146  G.C.E. (A/L) CHEMISTRY: UNIT 6 Chemistry of s, p and d block elements SO2(g) + 2FeCl3(aq) + 2H2O(l) H2SO4(aq) + 2FeCl2 (aq) + 2HCl(aq) 4.6.8 Oxoacids of sulphur Common oxidation numbers of sulphur are -2, 0, +2, +4 and +6. Sulphuric acid Sulphuric acid is a strong diprotic acid. Sulphur trioxide reacts with water to produce sulphuric acid. SO3(g) + H2O(l) H2SO4(aq) H2SO4(aq) + H2O(l) HSO4-(aq) + H3O+(aq) HSO4-(aq) + H2O(l) SO42-(aq) + H3O+(aq) Concentrated sulphuric acid can act as a dehydrating agent. conc. H2SO4 C6H12O6(s) 6C(s) + 6H2O(g) conc. H2SO4 C2H5OH(l) C2H4(g) + H2O(l) Concentrated hot sulphuric acid can act as an oxidizing agent. With metals, 2H2SO4(l) + Mg(s) SO2(g) + MgSO4(aq) + 2H2O(l) 2H2SO4(l) + Cu(s) SO2(g) + CuSO4(aq) + 2H2O(l) With nonmetals, S(s) + 2H2SO4(l) 3SO2(g) + 2H2O(l) C(s) + 2H2SO4(l) CO2(g) + 2SO2(g) + 2H2O(l) Dilute H2SO4 act as an acid. H2SO4(aq) + 2NaOH(aq) Na2SO4(aq) + 2H2O(l) H2SO4(aq) + Mg(s) MgSO4(aq) + H2(g) Dilute sulphuric acid is a strong acid which can protonate to give two H+ ions to water as shown below. H2SO4(aq) + 2H2O(l) SO42-(aq) + 2H3O+(aq) 147  G.C.E. (A/L) CHEMISTRY: UNIT 6 Chemistry of s, p and d block elements Figure 4.14 Structure of sulphuric acid Sulphurous acid Due to the air oxidation of sulphurous acid, it always contains a small amount of sulphuric acid. The reaction of gaseous sulphur dioxide and water produces sulphurous acid. The sulphurous acid reacts with dissolved oxygen in water to produce sulphric acid. Structure of the sulphurous acid is shown below. This acid is a weaker acid than sulphuric acid. Figure 4.15 Structure of sufurous acid Thiosulphuric acid Only the salts of thiosulphuric acid are stable and thiosulphate ion can oxidize as well as reduce to give sulphur and sulphur dioxide as its products. Thiosulphuric is a weak acid. In aqueous solutions, thiosulphuric acid can decompose to produce a mixture of sulphur containing products. H2S2O3(aq) S(s) + SO2(g) + H2O(l) Thiosulphate ion can act as a reducing agent. 2S2O32-(aq) + I2(aq) S4O62-(aq) + 2I-(aq) Structures of thiosulphuric acid and thiosulphate ion are shown below. The oxidation state of the central sulphur atom is +4 where as the terminal sulphur is zero in both structures. Figure 4.16 Thiosulphuric acid and thiosulphate ion 148  G.C.E. (A/L) CHEMISTRY: UNIT 6 Chemistry of s, p and d block elements 4.7 Group 17 elements 4.7.1 Group trends Halogens are reactive and can only be found naturally as compounds. Fluorine is the most electronegative element and exhibits -1 and 0 oxidation states. The halogens other than fluorine form stable compounds corresponding to nearly all values of the oxidation numbers from -1 to +7. However, compounds of bromine with the oxidation state of +7 are unstable. Due to the smaller atomic radius, fluorine can stabilize higher oxidation states of other elements. Oxidizing ability of halogens decreases down the group. Fluorine is a powerful oxidizing agent. The reactivity of halegons decreases down the group. This can be explained by using the displacement reactions of halegons. Cl2(aq) + 2Br-(aq) 2Cl-(aq) + Br2(aq) Br2(aq) + 2I-(aq) 2Br-(aq) + I2(aq) Fluorine and chlorine are gases with pale yellow and pale green colours respectively at room temperature. Bromine is a red-brown fumming liquid and iodine is a violet-black solid with lustrous effect. The bond energy of F2 (155 kJ mol-1) is less than that of Cl2 (240 kJ mol-1) due to repulsion between the non-bonded electron pairs of fluorine atoms. This is a reason for the high reactivity of fluorine gas. Down the Group 17 bond energies show a gradual decrease (Cl2 = 240 kJ mol-1, Br2 = 190 kJ mol-1 and I2 = 149 kJ mol-1). 149  G.C.E. (A/L) CHEMISTRY: UNIT 6 Chemistry of s, p and d block elements Table 4.12 Properties of Group 17 elements F Cl Br I **At Ground state electronic [He]2s22p5 [Ne]3s23p5 [Ar]3d104s24p5 [Kr]4d105s25p5 [Xe]4f145d106s26p5 configuration van der Waals 135 180 195 215 - radius/ pm Ionic radius X-/ pm 133 181 196 220 - Covalent radius/pm 71 99 114 133 - Melting point/ °C -220 -101 -7.2 114 - Boiling point/ °C -188 -34.7 55.8 184 - Pauling 4.0 3.2 3.0 2.7 - electronegativity Electron gain -328 -349 -325 -295 - enthalpy/ kJ mol-1 X(g) + e X-(g) **Not relevant to the current G. C. E. (A/L) Chemistry Syllabus 4.7.2 Simple compounds of Group 17 Hydrogen halides Hydrogen halides are acidic in water. HF has the ability to produce extensive hydrogen bonding, however, HF is a gas (boiling point 20 °C) at room temperature and under atmospheric pressure. Acidic nature of hydrogen halides in aqueous solutions For HF; HF(g) + H2O(l) H3O+(aq) + F-(aq) For other hydrohen halides (HCl, HBr and HI); HX(g) + H2O(l) H3O+(aq) + X-(aq) HF is a weak acid whereas the other hydrogen halides are strong acids in the aqueous medium. HF has the high bond energy (strongest covalent bond), which makes it difficult to dissociate in water to produce H+ ions readily. The acidic strength of hydrogen halides increases down the Group 17. This can be explained using the same fact mentioned above. Some selected properties of Group 17 hydrogen halides are shown in Table 4.13. 150  G.C.E. (A/L) CHEMISTRY: UNIT 6 Chemistry of s, p and d block elements Table 4.13 Selected properties of group 17 hydrogen halides HF HCl HBr HI Melting point/ °C -84 -114 -89 -51 Boiling point / °C 20 -85 -67 -35 Bond length/ pm 92 127 141 161 Bond dissociation energy/ kJ mol-1 570 432 366 298 Silver halides Silver halides can be used to identify the halides (chloride, bromide, and iodide) using the colour of the precipitate. Few selected properties are shown below. Table 4.14 Silver halides of Group 17 elements Silver halide Colour Solubility in ammonia AgCl White Dissolves in dil. aqueous ammonia AgBr Pale yellow Dissolves in conc. aqueous ammonia AgI Yellow Insoluble in both dil. and conc. aqueous ammonia Oxides and oxoacids of chlorine Chlorine forms several oxides and oxoanions with variable oxidation states. Some oxoanions are strong oxidizing agents. Selected oxides of chlorine are shown in Table 4.15. Table 4.15 Selected oxides and oxoanions of chlorine Oxidation Formula of oxide Formula of Structure of state oxoanion oxoanion +1 Cl2O ClO- +3 ClO2- +5 ClO3- +6 ClO3 and Cl2O6 +7 Cl2O7 ClO4- 151  G.C.E. (A/L) CHEMISTRY: UNIT 6 Chemistry of s, p and d block elements Chlorine forms four types of oxoacids. The acidic strength increases with the increasing oxidation number of the chlorine atom. The stuctures and the oxidation states of oxoacids are given in the Table 4.16. Table 4.16 Structures of oxoacids of chlorine HClO HClO2 HClO3 HClO4 Oxidation state +1 +3 +5 +7 Structure Oxidizing power of oxoacids of chlorine changed as follows. HClO > HClO2 > HClO3 > HClO4 The oxidation state of chlorine in HClO, HClO2, HClO3, HClO4 respectively are +1, +3, +5 and +7. The higher the oxidation state the stronger the acid will be. Therefore the variation of acidic strength is HClO < HClO2 < HClO3 < HClO4. Halides Most covalent halides react vigorously with water. But CCl4 does not hydrolyze. Most fluorides and some other halides are inert. Chlorides of group 14 and 15 elements react with less water as follows. SiCl4(l) + 2H2O(l) 4HCl(aq) + SiO2(s) PCl5(l) + H2O(l) POCl3(aq) + 2HCl(aq) Chlorides of group 14 and 15 elements react excess water as follows. SiCl4 (l) + 3H2O(l) 4HCl(aq) + H2SiO3(aq) NCl3(l) + 3H2O(l) NH3(aq) + 3HOCl(aq) PCl3(l) + 3H2O(l) H3PO3(aq) + 3HCl(aq) PCl5(l) + 4H2O(l) H3PO4(aq) + 5HCl(aq) AsCl3(s) + 3H2O(l) H3AsO3(aq) + 3HCl(aq) SbCl3(aq) + H2O(l) SbOCl(s) + 2HCl(aq) 152  G.C.E. (A/L) CHEMISTRY: UNIT 6 Chemistry of s, p and d block elements BiCl3(aq) + H2O(l) BiOCl(s) + 2HCl(aq) 4.7.3 Reactions of chlorine Chlorine is less reactive than fluorine. Chlorine gas is a strong oxidizing agent. Some reactions of chlorine act as a strong oxidizing agent are given below. 2Cu(s) + Cl2(g) 2CuCl(s) 2CuCl(s) + Cl2(g) 2CuCl2(s) Fe(s) + Cl2(g) FeCl2(s) 2FeCl2(s) + Cl2(g) 2FeCl3(s) excess ammonia, 8NH3(g) + 3Cl2(g) N2(g) + 6HCl(g) excess chlorine, 3Cl2(g) + NH3(g) 3HCl(g) + NCl3(l) Disproportionation reactions of chlorine Chlorine is simultaneously reduced and oxidized when it reacts with water and bases. Reaction of chlorine with water; Cl2(g) + H2O(l) HOCl(aq) + HCl(aq) In this reaction, zero oxidation state of chlorine (Cl2) oxidize to +1 (HOCl) and reduce to -1 (Cl¯). Reaction with sodium hydroxide; With cold dilute sodium hydroxide Cl2(g) + cold and dil. 2NaOH(aq) NaCl(aq) + NaOCl(aq) + H2O(l) With hot concentrated/ hot dilute sodium hydroxide above 80 °C 3Cl2(g) + conc. 6NaOH(aq) 5NaCl(aq) + NaClO3(aq) + 3H2O(l) Reactions of oxoanions ClO ̅ is stable at low temperatures and disproportionates at high temperature to produce Cl ̅ and ClO3 ̅. However, both BrO ̅ and IO ̅ are not stable even at low temperatures and undergo disproportionation. 153  G.C.E. (A/L) CHEMISTRY: UNIT 6 Chemistry of s, p and d block elements Disproportionation reactions of hypochlorite Disproportionation of hypochlorite to produce chlorate and chloride can be written as; 3ClO ̅ ClO3 ̅ + 2Cl ̅ Under acidic conditions, HOCl is more stable than ClO ̅ , which makes disproportionation predominant under basic conditions. 4.8 Group 18 elements 4.8.1 Group trends All group 18 elements are unreactive monoatomic gasses. Only Xe forms a significant range of compounds. All group 18 elements have positive electron gain enthalpy because an incoming electron needs to occupy an orbital belonging to a new shell. Table 4.17 Properties of Group 18 elements He Ne Ar Kr Xe Ground state electronic 1s2 [He]2s22p6 [Ne]3s23p6 [Ar]3d104s24p6 [Xe]4d105s25p6 configuration Atomic radius/ pm 99 160 192 197 240 1st ionization energy/ 2373 2080 1520 1350 1170 kJ mol-1 Electron gain 48.2 115.8 96.5 96.5 77.2 enthalpy/ kJ mol-1 4.8.2 Simple compounds of group 18 elements Compounds of xenon have oxidation numbers of +2, +4, +6 and +8. Xenon reacts directly with fluorine. Some Xe compounds are shown in Table 4.18. 154  G.C.E. (A/L) CHEMISTRY: UNIT 6 Chemistry of s, p and d block elements Table 4.18 Some selected compounds of Xe Oxidation state Compounds Structure +2 XeF2 +4 XeF4 +6 XeF6 +6 XeO3 +8 XeO4 4.9 Periodic trends shown by s and p block elements 4.9.1 The valence electron configuration The valance electron configuration of an element can be predicted from their position in the periodic table. Group number 1 2 13 14 15 16 17 18 1 2 2 1 2 2 2 3 2 4 2 5 Valance shell ns ns ns np ns np ns np ns np ns np ns2np6 electron configuration 155  G.C.E. (A/L) CHEMISTRY: UNIT 6 Chemistry of s, p and d block elements 4.9.2 Metallic character Metals have lower ionization energies compared to the other elements. Hence, metals can easily release electrons to produce cations. The atomic radii increase and ionization energy decreases when going down a group. Therefore, the metallic nature increases down the group. Also, across a period, atomic radii decrease and ionization energy increases. Therefore, the metallic nature decreases. The third period shows a gradual increase in melting point and then a decrease across the period. Most abundant elemental form, type of bonding between similar atoms and the melting point of the third period elements are shown below. Table 4.19 Most abundant elemental form, type of bonding between similar atoms and the melting point of the third period elements Na Mg Al Si P4 S8 Cl2 Ar Melting point/ °C 98 649 660 1420 44 119 -101 -189 Bonding type M M M NC C C C - Metallic – M, Network covalent – NC, Covalent - C Acid, base and amphoteric nature of oxides Across the third period variation of type of bonding in oxides in which the elements are at their highest oxidation number are given below. Table 4.20 Comparison of the third period oxides Na2O(s) MgO(s) Al2O3(s) SiO2(s) P4O10(s) SO3(g) Cl2O7(l) Oxidation +1 +2 +3 +4 +5 +6 +7 number Bonding I I I NC C C C type Nature Strongly B Am Very Weakly A Strong B weakly A A A Ionic – I, Network covalent – NC, Covalent - C Basic – B, Amphoteric – Am, Acidic - A Oxides with the highest oxidation number are considered to compare the chemical nature. The nature from strong basic on the left to strong acidic to the right can be seen. Amphoteric nature can be seen in the middle of the series. 4.9.3 Reactions of third period oxides with water, acids and bases Oxides of sodium and magnesium react with water to produce hydroxides. Na2O(s) + H2O(l) NaOH(aq) MgO(s) + 2H2O(l) Mg(OH)2 156  G.C.E. (A/L) CHEMISTRY: UNIT 6 Chemistry of s, p and d block elements As these two oxides are basic, they react with acids to produce salt and water. Na2O(s) + 2HCl(aq) 2NaCl(aq) + H2O(l) MgO(s) + 2HCl(aq) MgCl2(aq) + H2O(l) Aluminum oxide is amphoteric and it reacts with acids as well as with bases to produce salts. Al2O3(s) + 6HCl(aq) 2AlCl3(aq) + 3H2O(l) Al2O3(s) + 2NaOH(aq) + 3H2O(l) 2Na[Al(OH)4](aq) SiO2 is weakly acidic and reacts with strong bases. Also, SiO2 shows no reaction with water. SiO2(s) + 2NaOH(aq) Na2SiO3(aq) P4O10, SO3, and Cl2O7 are acidic and produce acids when dissolved in water. Those reactions are shown below. P4O10(s) + 6H2O(l) 4H3PO4(aq) SO3(g) + H2O(l) H2SO4(aq) Cl2O7 (l) + H2O(l) 2HClO4(aq) These oxides also react with bases to produce salts and water. P4O10(s) + 12NaOH(aq) 4Na3PO4(aq) + 6H2O(l) SO3(g) + 2NaOH(aq) Na2SO4(aq) + H2O(l) Cl2O7 (l) + 2NaOH 2NaClO4(aq) + H2O(l) 4.9.4 Acid, base and amphoteric nature of hydroxides and hydrides Hydroxides of the third period show a trend similar to oxides of the same period. The following table shows a comparison of the third period hydroxides. 157  G.C.E. (A/L) CHEMISTRY: UNIT 6 Chemistry of s, p and d block elements Table 4.21 Comparison of the third period hydroxides NaOH Mg(OH)2 Al(OH)3 Si(OH)4 P(OH)5 S(OH)6 Cl(OH)7 Stable form H2SiO3 H3PO4 H2SO4 HClO4 Oxidation +1 +2 +3 +4 +5 +6 +7 number Bonding I I C C C C C type Nature Strongly B Am Very Weakly Strongly Very B weakly A A A strongly A Ionic – I, Network covalent – NC, Covalent - C Basic – B, Amphoteric – Am, Acidic - A Nature of hydrides of third period varies from strong bases to strong acids across the period. Amphoteric nature can be seen in the middle of the series. Table 4.22 Comparison of the third period hydrides NaH(s) MgH2(s) (AlH3)x(s) SiH4(g) PH3(g) H2S(g) HCl(g) Oxidation +1 +2 +3 -4 -3 -2 -1 number Nature of Strongly Weakly Am Very N Weakly A Very the aqueous B B weakly strongly A solution A Bonding I I NC C C C C type Ionic – I, Network covalent – NC, Covalent - C Basic – B, Amphoteric – Am, Acidic – A, Neutral - N Hydrides of sodium and magnesium react with water to produce basic solutions. NaH(s) + H2O(l) NaOH(aq) + H2(g) MgH2(s) + 2H2O(l) Mg(OH)2(s) + 2H2(g) AlH3(s) + 3H2O(l) Al(OH)3(s) + 3H2(g) PH3 is weakly soluble in water and produces a neutral solution. H2S and HCl are acidic and aqueous solutions are also acidic. H2S(g) + H2O(l) HS ̶ (aq) + H3O+ (aq) HCl(g) + H2O(l) Cl ̶ (aq) + H3O+(aq) 158  G.C.E. (A/L) CHEMISTRY: UNIT 6 Chemistry of s, p and d block elements 4.9.5 Nature of the halides across the third period As the electronegativity of elements increases across the period from left to right, the ability of hydrolyzation of chlorides increases accordingly. Corresponding reactions are given below. Chlorides of s block elements in the third period are ionic and the p block elements are covalent. Table 4.23 Comparison of the third period chlorides NaCl(s) MgCl2(s) AlCl3(s) SiCl4(l) PCl5(g) SCl2(g) Oxidation number +1 +2 +3 +4 +5 +2 Bonding type I I C C C C Nature of the aqueous N Very A A A A solution weakly A Ionic – I, Covalent - C Basic – B, Amphoteric – Am, Acidic – A, Neutral - N Reactions with water of third period covalent chlorides are, AlCl3(s) + H2O(l) [Al(H2O)5OH]2+(aq) + H3O+(aq) SiCl4(l) + 2H2O(l) SiO2(s) + 4HCl(aq) PCl5(g) + 4H2O(l) H3PO4(aq) + 5HCl(aq) 2SCl2(g) + 3H2O(l) H2SO3(aq) + S(s) + 4HCl(aq) Group 15 can be used to understand the variation of properties down the group. Down a group the ionization energy decreases, and the metallic nature increases. Use the information given for the Group 15 and correlate the variation in ionization energies with the increase of metallic properties down the group. Both N and P are nonmetals and produce acidic oxides. However, As and Sb oxides are amphoteric and bismuth oxide is basic. Reactions with water of group 15 halides are given in the respective section under the halides of group 17. 159  G.C.E. (A/L) CHEMISTRY: UNIT 6 Chemistry of s, p and d block elements d Block Elements Elements in Groups 3 to 12 are collectively classified as d block elements. In d block elements the last electron gets filled into a d orbital. These elements can be categorised into two categories namely transition and non-transition. 4.10 Transition elements d block elements contain incompletely filed d subshell at elemental state or with the ability to form at least one stable ion with incompletely filled d subshell are called transition elements. Therefore, d block elements producing ions only with d10 configurations are considered as non-transition elements. e.g.: Electronic configurations of Zn : [Ar]3d104s2 Electronic configurations of Zn2+: [Ar]3d104s0 Electronic configuration of Sc : [Ar]3d14s2 Electronic configuration of Sc3+ : [Ar]3d04s0 Both Zn and Sc are d block elements (last electron is filled to a 3d orbital). However, Zn is considered as a non-transition element due to the absence of a partially filed d subshell at the elemental stage and Zn2+ ion. Sc can be considered as a transition element since Sc contains partially filed d subshell at the elemental stage. Table 4.24 Comparison of the properties of d block elements in fourth period Group 3 4 5 6 7 8 9 10 11 12 Element Sc Ti V Cr Mn Fe Co Ni Cu Zn Pauling 1.3 1.5 1.6 1.6 1.5 1.9 1.9 1.9 1.9 1.6 electronegativity Atomic 162 147 134 128 127 126 125 125 128 137 radius/pm Covalent 144 132 122 118 117 117 116 115 117 125 radius/pm Ionic radius (M2+)/ - 100 93 87 81 75 79 83 87 88 pm Transition metal ions have less variation in atomic radii across a period than that of the main group elements. Across the period of the transition metals shown in Table 4.23, the atomic radii decrease slightly and then increase. Across the period, to each d electron added nuclear charge is also increased by one. The decrease of the atomic radii at the middle of the period (from Sc to Ni) occurs due to the predominance of attraction power of nuclear charge increase than the repulsion among the electrons. However, at the end of the period (Cu and Zn), radii of the atoms increase due to greater repulsion among electrons as electrons are paired in d orbitals. 160  G.C.E. (A/L) CHEMISTRY: UNIT 6 Chemistry of s, p and d block elements 4.10.1 Occurrence Elements on the left of the 3d series (fourth period d block elements) exit commonly in the nature as metal oxides and cations combined with anions. Few examples are shown below. Table 4.25 Occurrence of some fourth period d block elements Element Example Ti FeTiO3 (Ilmenite) and TiO2 (Rutile) Fe Fe2O3 (Haematite), Fe3O4 (Magnetite) and FeCO3 (Siderite) Cu CuFeS2 (Copper Pyrite) 4.10.2 Properties of fourth period d block elements Oxidation states and ionization energies Except Sc and Zn in the fourth period d block elements, others can form stable cations with multiple oxidation states. The multiplicity of the oxidation state is due to the varying number of d electrons participate in bonding. Both Zn (+2) and Sc (+3) only produce ions with a single oxidation state, and these ions do not contain partially filled d orbitals. Electron configuration and the oxidation states of d block elements are shown in Table 4.26. Sc forms only Sc3+ ions. Except in Sc, +2 oxidation number can be seen in all the other elements since electrons in 4s orbital get removed due to ionization before electrons in 3d orbitals. Reason for this is that the 4s orbital with two electrons in the outermost shell experiences a lesser effective nuclear charge than that of electrons in the 3d orbital. As a result of the 3d104s1 configuration, Cu can form +1 oxidation number commonly. However, Cr+ is extremely rare and unstable even though Cr has 3d54s1 configuration. The highest possible oxidation number that a d block element can show is the sum of 4s and 3d electrons. Transition metals are also capable of producing variable oxidation states similar to p block elements and show the ability to interconvert among their oxidation states. Therefore, they can act as oxidizing as well as reducing agents. First five elements achieve the maximum possible oxidation state by losing all 4s and 3d electrons. With the filling of more 3d electrons, towards the right end of the period, the 3d orbitals become greater in energy as the nuclear charge of the atom increases. This makes d electrons are harder to remove. The most common oxidation state for these elements is +2 due to the loss of 4s electrons. Reactivity d block elements do not react with the water while s block elements react with water vigorously. The 4s electrons of the d block elements are tightly bound to the nucleus due 161  G.C.E. (A/L) CHEMISTRY: UNIT 6 Chemistry of s, p and d block elements to the higher nuclear charge than that of the s block elements. First ionization energy of d block elements lies between the values of those of s and p block elements. Table 4.26 Electronic configuration and oxidation states of d block elements Element Ground state configuration Oxidation states 3d 4s 1 2 Sc [Ar]3d 4s +3 Ti [Ar]3d24s2 (+2), +3, +4 V [Ar]3d34s2 (+2), (+3), +4, +5 Cr [Ar]3d54s1 +2, +3, (+4), (+5), +6 Mn [Ar]3d54s2 +2, +3, +4, (+5), (+6), +7 Fe [Ar]3d64s2 +2, +3, (+4), (+5), (+6) Co [Ar]3d74s2 +2, +3, (+4) Ni [Ar]3d84s2 +2, (+3), (+4) Cu [Ar]3d104s1 +1, +2, (+3), (+4) Zn [Ar]3d104s2 +2 *Less common states are shown in brackets. Ionization energies of fourth period d block elements are higher than that of the s block elements in the same period. The first ionization energies of d block elements are increase slightly across the period when move from left to the right of the period. Variation of the first ionization energy across the d block is less than that of s and p block elements. Increase in the nuclear charge across the fourth period d block elements expect to be increase the first ionization energies due to the greater attraction towards the 4s electrons. However, in all d block elements, extra electrons are inserted in to the 3d orbital moving from left to right across the period, and these d electrons shield the 4s electrons from the inward attraction of the nucleus. Because of these two counter effects, the ionization energy of d block elements increases slightly across the period. Successive ionization energies of the fourth period d block elements are shown in the table given below. 162  G.C.E. (A/L) CHEMISTRY: UNIT 6 Chemistry of s, p and d block elements Table 4.27 Successive ionization energies of fourth period d block metals, K and Ca.** Element 1st ionization 2nd ionization 3rd ionization energy/ kJ mol-1 energy/ kJ mol-1 energy/ kJ mol-1 K 418 3052 Ca 589 1145 4912 Sc 631 1235 2389 Ti 658 1310 2652 V 650 1414 2828 Cr 653 1496 2987 Mn 717 1509 3248 Fe 759 1561 2957 Co 758 1646 3232 Ni 737 1753 3393 Cu 746 1958 3554 Zn 906 1733 3833 ** For K, only first and second ionization energies are given to understand the energy increase due to removal of an electron from an inner orbital. First ionization energies of d block elements are higher than those of s block elements in the same period. This explains the less reactivity of d block elements than the s block elements. All d block elements are metals because 4s electrons in d block elements can be released easily to form cations. Metallic character of the d block elements increases down the group. All d block elements in the fourth period are solids with high melting and boiling points. Melting and boiling points of d block elements are extremely high as compared to those of s and p block elements. d block elements are moderately reactive. Except metal ions with 3d0 and 3d10 configurations, d block metal compounds produce characteristic colours. This means transition metal ion complexes can produce coloured compounds. Most d block metal ions form complex compounds. Electronegativity Table below provides the electronegativity of d block elements and can be used to understand the variation of electronegativity of d block elements in the fourth period. Electronegativity increases with the atomic number. However, Mn and Zn are deviated from the trend due to their stable electron configuration. Due to the higher nuclear charge, d block elements have higher electronegativity than that of the s block elements. 163  G.C.E. (A/L) CHEMISTRY: UNIT 6 Chemistry of s, p and d block elements Element Sc Ti V Cr Mn Fe Co Ni Cu Zn Electronegativity 1.3 1.5 1.6 1.6 1.5 1.8 1.8 1.8 1.9 1.6 When an atom exhibits variable oxidation states, the higher oxidation state has higher electronegativity. Catalytic properties Most transition metals and compounds can act as catalysts due to the presence of partially filed and empty d orbitals. This makes d orbitals to accept or donate electrons. This property makes them effective components of catalysts. Pd for hydrogenation, Pt/Rh for oxidation of ammonia to nitrogen oxide, and V2O5 for oxidation of SO2 to SO3 and TiCl3/Al(C2H5)6 for the polymerization of ethene are some examples for the use of d block element and its compound as a catalyst. Some popular organic reactions such as alkylation and acylation are done in the presence of transition metal ion as the catalyst. Colours of transition metal ions Aqueous solutions of many transition metal ions can absorb radiation in the visible region of the electromagnetic spectrum to produce various colours. This ability is due to the presence of partially filled d subshells. In contrast, metal ions of s block are colourless because these ions have completely filled subshells. The following Table shows some of the colours of transition metal ions and oxoanions in aqueous solutions. For example, [Co(H2O)6]2+ is pink, [Mn(H2O)6]2+ pale pink. In contrast, aqueous solutions of Sc3+ and Zn2+ are colourless due to the unavailability of partially filled d orbitals. Also, ions with d0 or d10 configuration are coloureless when in an aqueous solution. Colours of MnO4- and CrO42- are not due to the electron transition of electrons among the d orbitals. Colours of some elected oxoanions are given in Table 4.28. Table 4.28 Colours of d block metal ions and oxoanions in aqueous solutions. The number of 3d and 4s electrons are shown in brackets next to the metal ion. Ion Colour Ion Colour Sc3+ (d0 s0) Colourless Fe3+(d5 s0) Brown yellow Ti4+(d0 s0) Colourless Fe2+(d6 s0) Pale green Cr3+(d3 s0) Violet Co2+(d7 s0) Pink Mn2+(d5 s0) Pale pink Ni2+(d8 s0) Green Cu2+(d9 s0) Blue Cu+(d10 s0) Colourless Zn2+(d10 s0) Colourless Oxoanion Colour Oxoanion Colour MnO4- Purple CrO42- Yellow MnO42- Green Cr2O72- Orange 164  G.C.E. (A/L) CHEMISTRY: UNIT 6 Chemistry of s, p and d block elements 4.10.3 Oxides of d block elements First four elements form oxides by removing all valence electrons. Unlike main group elements, transition elements produce different oxidation states. Some d block elements can form oxides in which metal atom presence with two different oxidation numbers. Both Mn3O4 and Fe3O4 are examples for binary oxides (which are formed with two oxidation numbers). Mn3O4 is a mixture of Mn(II) and Mn(III). Also, Fe3O4 is a mixture of Fe(II) and Fe(III). 4.10.4 Chemistry of some selected d block oxides Chromium and manganese oxides Properties of an oxide depend on the oxidation number. The bonding type depends on the oxidation number. The change in the bonding type explains the basis in the acid-base behaviour of metal oxides. For the compounds with high oxidation numbers have covalent bonding characteristics are acidic and the compounds with low oxidation numbers have ionic bonding characteristics are basic. Table 4.29 Acid-base nature of chromium oxides Oxide Acid-base nature Oxidation number CrO Weakly basic +2 low oxidation state Cr2O3 Amphoteric +3 moderate oxidation state CrO2 Weakly acidic +4 CrO3 Acidic +6 high-oxidation state Generally, if the metal is in a lower oxidation state, the oxide is basic. Also, if the metal is in a moderate oxidation state, the oxide is amphoteric and metal oxides with higher oxidation state are acidic. This explains why the compounds in Tables 4.29 and 4.30 with lower oxidation states are more metallic while compounds with higher oxidation states are more non-metallic in properties. Table 4.30 Acid-base nature of manganese oxides Oxide Acid-base nature Oxidation number MnO Basic +2 Low oxidation state Mn2O3 Weakly basic +3 MnO2 Amphoteric +4 moderate oxidation state MnO3 Weakly acidic +6 Mn2O7 Acidic +7 high-oxidation state 165  G.C.E. (A/L) CHEMISTRY: UNIT 6 Chemistry of s, p and d block element