Introduction Of Chemical Concepts: Caesium And Iodine Xinyi Huang

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Caesium (Cs), which is also written as cesium in American spelling, has the atomic number of 55 and molar mass of 132. 91. The name of caesium was given because of two beautiful ‘sky blue’ colour lines in its spectra and the word Caesium means sky blue in Latin. [1] Because of the physical and chemical properties, caesium is unstable and highly reactive in nature and consequently, the abundance of caesium is significantly less than the abundance of other elements from Group 1. Its main nature occurrence form is hydrated aluminosilicate, Cs4Al4Si9O26∙H2O, but according to Greenwood and Earnshaw [2], ‘the world’s only commercial source is at Bernice Lake, Manitoba’. As caesium is a part of the products coming with the extraction of aluminium, one of the extraction methods of caesium is based on the hydrated aluminosilicate in mineral with sulphuric acid and carbon. The general process of the extraction method is firstly using sulphuric acid to remove silicon from the hydrated aluminosilicate and secondly using carbon in high-temperature reaction environment to get the caesium sulphate, Cs2SO4.

The density of caesium in the solid state is 1879 kg m-3 and the melting point is 301. 59 K, which is about 28. 44℃. [4] As is well-known, the valence shell electrons play an important role in both physical and chemical properties of the elements. The valence shells with ns1 (n=2,3,4,5,6,7) of Group 1 elements contribute to the low melting points and softness of alkali metals [5] as only one electron of each group 1 atom is provided when forming the metallic bond, which results in the metallic bonds in alkali metals are weaker than the ones in other metals. What’s more, caesium and other alkali metals have a good ability of electricity conduction because when forming the metallic bonds, the ns1 orbitals provide unfilled bands, conduction bands (each has only one electron) and fulfilled bands. When the voltage is provided to the metal, the electrons in the fulfilled bands will get enough energy to jump to the unfilled band. The originally unfilled band which now contains one electron has the negative charge and moves to the positive electrode, while the originally fulfilled band which now loses one electron has the positive charge and moves to the negative electrode. When the chemical properties of elements are discussed, not only the valence shell electrons but also other ones such as the radii of atoms should be considered as well. With the increasing atomic number, the radius of the atom from the Group 1 descends, and this leads to the decreasing of the first ionization energy. According to the values from Emsley, the first ionization energy of Cs is 376 kJ mol-1 [6] due to the increasing distance between the valence shell electron and the nucleus down the group, the attraction between the nucleus and the valence shell electron becomes smaller.

The definition of electron affinities is the change of energy when an atom turns in to an anion and Norman explained that [7] ‘it is the reverse of ionization’. Caesium has the lowest electron affinity as the electron affinity decreases down the group 1. When the anion is formed, the ns orbitals are fulfilled which needs high energy and as shown above, the energy needed to form a fulfilled ns orbital decreases with the attraction between the nucleus and valence shell electron getting smaller. Another critical chemical property is electronegativity which is one of the factors should be considered with chemical compounds. The electronegativity within a group has the trend of decreasing down the group and according to the values from Weller, Overton, Rourke and Armstrong [8], the Pauling electronegativity of caesium is 0. 79 that is smaller than the most elements in the periodic table. Before starting the oxides and halides of caesium, its oxidation states should be illustrated as this feature influences the synthesis and analysis of the caesium compounds distinctly. According to table 4. 1 from Norman [9], the common oxidation state of group I elements is +1. The group 1 elements usually give the ns1 electron to the atoms with higher electronegativity to form chemical bonds and end with the fulfilled and completed valence shell. For example, in the oxide Cs2O, the oxidation states of Cs and O are +1 and -2 respectively and, in the halides, CsF, CsCl, CsBr and CsI, the oxidation state of Cs is +1. The next paragraph will briefly show the formation and reaction of typical oxides and halides of caesium. Alkali metals can make various compounds with oxygen and among the alkali metals, caesium has the most various compounds with oxygen. [10]

The two typical oxides of caesium that will be discussed are Cs2O2 and Cs2O. According to Arora [11], when the caesium is burning in the air, both two oxides above can be formed. However, the most reaction between oxygen and caesium is, 2Cs + O2 → 2Cs2O2 (the reaction condition is burning) which product is peroxide caesium. When the Cs2O2 is formed, the two electrons from two 6s orbitals of caesium with two electrons from 2p orbitals of oxygen form two σ bonds respectively, while the two parallel 2p orbitals from each oxygen atom form a π bond. Another more efficient way to prepare Cs2O, which has the bent structure using the VSEPR model, is using the reduction between caesium hydroxide and caesium [11]:Cs + CsOH → Cs2O + H2↑ The caesium chloride, CsCl, is a typical example of caesium halides and it is one of the reactants to extract Cs. The interesting thing is, due to the large radius of a caesium atom, compared to other halides in group 1, caesium chloride and other halides of caesium have the primitive cubic in solid structure while other halides have the body centred cubic. The caesium chloride can be obtained by the neutralization between hydroxide caesium, CsOH, and hydrochloric acid, HCl: CsOH + HCl → CsCl + H2O [12]or the reaction between caesium carbonate, Cs2CO3, and hydrochloric acid, HCl: Cs2CO3 + 2HCl → 2CsCl + H2O + CO2↑ [12]

The method to extract caesium from ‘highly purified’ caesium chloride is using reduction with ‘fractionally redistilled’ Ca:2CsCl + Ca → CaCl2 + 2Cs [13]Although caesium has not been discovered any important biological uses in the human body, its industrial values are wildly applied to various fields such as being the catalyst of several hydrogenations in Organic Chemistry, ‘ion propulsion system’ in Astronomy and even in light-sensitive cells and tubes in Biological area. [14] Through those industrial uses of caesium, the most famous one is the atomic clock, and this clock is mostly based on the physical properties of caesium. In International System of Units, ‘the second is the duration of 9 192 631 1770 periods of the radiation corresponding to the transition between the two hyperfine levels of the ground state of caesium 133 atoms. ’ [15] The basic mechanism of an atomic clock is using the exact frequency of caesium-133, an isotope of caesium. The caesium-133 atom is put into a long funnel with a detector in the atomic clock and every time when the caesium jumps to and jumps back between the two energy states, the detector will record once, and time took between two records is one second. [16] However, even caesium benefits the scientific and industrial fields significantly in the last century, the serious effects on the environment because of its radioactive isotopes cannot be ignored. In 1986, a fire in a nuclear reactor in Europe caused severe 137Cs radioactive pollution. [17] Iodine (I) Word Count: 1201Iodine, one of the least abundant elements across the periodic table, has the atomic number of 53 and the molar mass of 126. 90. In 1811, the discovery of Iodine was made by a chemist called B. Courtois and the name of Iodine was given for its dark purple colour. [18] According to Cox [19], iodine is mostly found in the seaweeds and oceans, and the main nature form of iodine is IO3-. There are several methods to extract iodine, and the most vital and earliest method is to extract iodine from seaweed. The first person who extracted iodine from seaweed is Courtois by using concentrated H2SO4. [20] There are three major steps in this method. Firstly, iodide, I-, is obtained by burning the seaweed and then burnt seaweed is put into water so the iodide moves from seaweed ash into the water. [21] Secondly, hydrogen peroxide is used to convert iodide to iodine by the oxidation in an aqueous solution which is usually sulphuric acid [22]:(a). H2O2 + I- + H+ → HIO + H2O(b). HIO + I- + H+ → I2 + H2OFinally, the pure I2 can be obtained by heating which leads to the evaporation of water as the boiling point of I2 is higher. In the human body, Iodine plays an important role in the endocrine system, especially in the thyroid gland.

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According to Sargis [23], the main function of thyroid gland is transferring iodine coming with food especially seafood into T3 (thyroxine) and T4 (triiodothyronine). Both iodine hormones are vital figures when the thyroid gland is checked and the content of them in the human body can reflect the illness of thyroid gland, for example, hyperthyroidism and hypothyroidism. In some countries such as China, iodised salt is produced and sold to the public to prevent some thyroid illness during the last decades. However, even iodine has an irreplaceable position in the human body, Hora illustrated [24] only 3% productions of iodine are used for human and the main industrial uses of iodine such as X-ray contrast media and liquid crystal display occupies more percentage in the total production of iodine. As shown in the table 17. 1 from Weller, Overton, Rourke and Armstrong [25], iodine has the ionic radius of 220pm, the melting point of 220℃, the first ionization energy of 1008 kJ mol-1, the electron affinity of 295 kJ mol-1 and the Pauling electronegativity of 2. 6. The ionic radius of group 17 element increases down the group.

Due to the atomic number increases, more shells are gained around the atom and this is the reason for the radius trend in group 17. For example, the valence electron configurations of bromine and iodine are [Ar]3d104s24p5 and [Kr]4d105s25p2, which shows iodine contains one more shell than bromine and in fact, the ionic radius of bromine is 34 pm less than the ionic radius of iodine. According to Norman [26], when the solid states are considered, ‘the ratio of the short, covalent bonds to the longer, intermolecular distances decreases as the group is descended’ and this means the solid structure stableness decreases with decreasing atomic number. Another reason for the trend is that the electronegativity of halogens decreases down the group and less electronegativity leads to stronger metallic bonds in the solid structure. Therefore, it is obviously shown that the melting point of halogens increasing in a descended group.

There is a decreasing trend of the first ionization energy down the group 17 and this is because the radii of halogens increase and the attractions between the nucleus and the valence shell electron on np orbitals descend. Another trend is halogens have higher first ionization energies than the first ionization energies of other elements from same periods and the principal reason is the electrons removed of halogens are from almost fulfilled shells. Because both electron affinities and electronegativities of group 17 elements decrease down the group and higher than the figures of elements from the same period, these will be explained as the results of increasing radii down the group and more filled valence orbitals across the same period, while the actual situation is far more complicated. Halogens normally have the oxidation range from -1 to +7 and the most common oxidation states of iodine are +1, +5, +7 and -1. [27] The four oxidation states will be illustrated in detail with the synthesises and reactions of specific iodine compounds in the next two paragraphs. IF, iodine monofluoride, is made by two halogens.

In this compound, the oxidation state of iodine is +1, while the oxidation state of fluorine is -1. As shown in the previous paragraph, fluorine has a higher electronegativity than iodine so when the IF is formed, the fluorine atom attracts the electron more strongly. According to Greenwood and Earnshaw [28], the synthesis of IF is using the reaction between I2 and F2 at low temperature (about 45℃) environment with the catalyst CCl3F:I2 + F2 → 2IFHowever, this compound is not stable under the room temperature and will quickly react and results in IF5, where the oxidation of iodine turns to +5. IF5, iodine pentafluoride, is a strong oxidizer and at the same time, hazardous and corrosive. [29] In Lewis Structure, IF5 has five σ bonds and one pair of lone electrons so the structure of square pyramidal according to the VSEPR model theory.

The hybridisation of iodine pentafluoride is sp3d2 as the 5 σ bonds should be same and have enough space for the lone pair. The most common method to prepare IF5 is burning iodine in F2:I2 + 5F2 → 2IF5Iodine pentafluoride can make vigorous reaction with water and results in the production of hydrogen fluoride, HF [29]:IF5 + 3H2O → 5HF + HIO3When considering the +5 oxidation state of iodine, there is one oxide of iodine cannot be ignored, I2O5. The oxides of iodine are the most stable ones among the oxides of halogens. [30] The synthesis method of iodine pentoxide is using the oxidation reaction between iodine and nitric acid, HNO3 [31]:3I2 + 10HNO3 → 3I2O5 + 5H2O + 10NOIt is worth mentioning that, iodine pentoxide can convert CO into CO2 [30]:5CO + I2O5 → I2 + 5CO2

Besides IF and IF5, iodine can also form IF7, iodine heptafluoride, in which the oxidation state of iodine is +7. Iodine heptafluoride has the molecule structure of pentagonal bipyramidal because the Lewis Structure of it contains seven σ bonds without any lone electron pair and the hybridization of iodine heptafluoride is sp3d3 which contains 7 σ orbitals. The compounds of iodine above all contain the iodine with positive oxidation states, while in hydrogen iodide and sodium iodide, the oxidation state of iodine is -1 because of the higher electronegativity. Hydrogen iodide, HI, is highly corrosive and can be made by reaction of hydrogen sulphide and iodine and sodium iodide, NaI, can be used to treat iodine deficiency. The synthesis method od sodium iodide is using the simple displacement reaction between hydrogen iodide, HI, and sodium carbonate, Na2CO3, or sodium hydroxide, NaOH [32]:with Na2CO3: 2HI + Na2CO3 → 2NaI + H2O + CO2↑with NaOH: HI + NaOH → NaI + H2O

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