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CHEMICAL SCIENCES SUBJECT CODE – 102 INORGANIC CHEMISTRY UNIT - III OSN ACADEMY www.osnacademy.com LUCKNOW 0522-4006074 9935977317 [1] Chemistry Ch. No. Contents Pg. Nos. 1 Main Group Element 3 – 46 2 Co-ordination Compound 1 - 33 [2] Chemistry 1. The correct order of stability of difluorides is : [June 2011] (a) GeF2>SiF2>CF2 (b) CF2>SiF2 > GF2 (c) SiF2> GeF2>CF2 (d) CF2> GF2>SiF2 Ans. (a) Hint: Inter pair effect (As we move down the group there is a reluctance of ns electrons to participate in bond formation) Therefore +2 state stability will be in the order:- Ge>Si>C. Therefore the correct order of stability of difluorides is therefore GeF2>SiF2>CF2 2. Lewis acidity of BCl3, BPh3 and BMe3 with respect to pyridine follows the order [Dec. 2011] (a) BCl3>BPh3>BMe3 (b) BMe3>BPh3>BCl3 (c) BPh3>BMe3>Bcl3 (d) BCl3>BMe3>BPh3 Ans. (b) Lewis acidity of BCl3, BPh3 and BMe3 with respect to pyridine follows the order. –BMe3> BPh3>BCl3 Among BCl3, BMe3 & PPh3 & BCl3 is weak Lewis acidic due 2 pB 3 pB 3 pCl back bonding. Among BMe3, and BPh3, BMe3 is strong Lewis acid with respect to pyridine due to steric factor in BPh3. 3. The material that exhibits the highest electrical conductivity among the following sulfur – nitrogen compounds is : [Dec 2011] (a) S4N4 (b) S7NH (c) S2N2 (d) (SN)x Ans. (d) Sulfur forms a variety of covalent binary nitrides, but the two most interesting ones are tetrasulfur & tetranitride S4N4, and disulfur dinitirde, S2N2, because they are precursors to an unusual polymer called polythiazyl, (SN)x. This polymeric sulfur nitride is unusual because, even though it is composed solely of two nonmetals, if exhibits some properties normally associated only with metals like electrical and thermal conductivity & It becomes a superconductor at 0.26 kelvin. [June 2012] (a) Cl>S>P>Si (b) Cl>P>S>Si (c) P>S>Si>Cl (d) Si>P>S>Cl Ans. (d) Hint: The size of the s, p and d orbitals generally decreasses with decrease in size of an element And as we know in periods, the atomic radius decreases across the period from left to right. Therefore order of the sizes of d orbitals is Si>P>S>Cl 4. The size of the d orbitals in Si, P, S and Cl follows the order. 5. The electronegativity differences is the highest for the pair 6. Which of the following pairs has the highest difference in their first ionization energy ? [June 2013] [Dec. 2012] (a) Li, Cl (b) K,F (c) Na, Cl (d) Li, F Ans. (b) Among metal, order of electronegativity is Li>Na>K & among halide order of electronegativity is F>Cl. Therefore an electronegativity difference is the highest for the K and F. (a) Xe, Cs (c) Ar, K (b) Kr, Rb (d) Ne, Na Ans. (d) Order of size of alkali metal is Na<K<Rb<Cs & first ionizatioin energy is Na>K>Rb>Cs while that of Nobel gases, order of size is Ne<Ar<Kr<Xe & first ionization energy is Ne> Ar>Kr>Xe. Difference in first ionization energy betweend [3] Chemistry alkali metal & nobel gases of same period is decreases top to bottom. Thus the Na, Ne pairs has the highest difference in their first ionization energy. [Dec. 2013] (a) BF3OEt2 (b) SF4 (c) PFe6 (d) Sb2S3 Ans. (c) A hypervalent molecule (expanded octet) is a molecule that contains one or more main group elements with more than eight electrons in their valence shells e.g. PFe6 7. Among the following, an example of a hypervalent species is 8. The strength of p d bonding in E-O (E=Si, P, S and C) follows the order [June 2012] (a) Si-O>P-O>S-O>Cl-O (b) P-O>Si-O>S-O>Cl-O (c) S-O>Cl-O>P-O>Si-O (d) Cl-O>S-O>P-O>Si-O Ans. (d) As the size increases, strength of p d bonding decreases because inter nuclear distance between two atoms increases & strength of bond decreases due to poor overlap as bond is formed by side by side overlap. As increasing order dof size is Si>P>S>Cl & therefore order of strength of p d bonding is Cl-O>S-O>P-O>SiO [June 2014] (a) As>Al>Ca>S (b) S>As>Al>Ca (c) Al>Ca>S>As (d) S>Ca>As>Al Ans. (b) The correct order of decreasing electronegativity is : S(2.58)>As(2.18)> Al(1.61)>Ca(1). In periodic table electronegativity increases from left to right in period & decreases top to bottom in group. 9. The correct order of decreasing electronegativity of the following atoms is : 10. The correct order of the size of S, S2-,S2+ and S4+ species is : [June 2014] (a) S S 2 S 4 S 2 (b) S 2 S 4 S 2 S (c) S 2 S S 2 S 4 (d) S 4 S 2 S S 2 Ans. (c) The correct order of the size of . S 2 S S 2 S 4 General order of size:- Anoinic species>Neutral atom>Cation species. 11. The series with correct order of decreasing ionic size is : [GATE 2006] 2 2 2 2 (a) K Ca S Cl (b) S Cl K Ca 2 2 (c) K Cl Ca S (d) Cl K S 2 Ca 2 Ans. (b) The series with correct order of decreasing ionic size is S Cl K Ca 2 General order of size: Anoinic species > neutal atom>Cation species 2 12. 1 [GATE 2007] The Lewis acid character of BF3, BCl3 follows the order : (a) BF3<BBr3BCl3 (c) BF3<BCl3<BBr3 (b) BCl3<BBr3<BF3 (d) BBr3<BCl3<BF3 Ans. (c) The Lewis acidity BF3 is less than BCl3 because of stronger 2 pB 2 pF back bonding in case of BF3 which is not effective in case of BCl3 due to weaker 2 pB 3 pCl back bonding and this back bonding became worse [4] Chemistry in case of BBr3; due to poorer 2 pB 4 pBr back bonding. Therefore Lewis acidity order is BF3 BCl3 BBr3 Bl3 13. [GATE 2009] Among the following, the isoelectronic and isostructural pair is : (a) CO2 and SO2 (c) NO2 and TeO2 (b) SO3 and SeO3 (d) SiO44 and PO43 Ans. (d) Among all SiO44 & PO43 is the isoelectronc and isostructural pair. Pair Molelcular/ion Total electrons (a) CO2 22 (b) (c) (d) 14. Structure according to VSEPR linear SO2 32 Bent or V shaped SO3 40 trigonal planar SeO3 58 trigonal planar NO2 TeO2 24 bent or V shaped 68 bent or V shaped SiO4 50 tetrahedral PO43 50 tetrahedral Among the following, the group of molecules that undergoes rapid hydrolysis is [GATE 2011] (a) SF6, Al2Cl6, SiMe4 (b) BCl3, SF6. SiCl4 (c) BCl3, SiCl4, PCl5 (d) SF6, Al2Cl6, SiCl4 Ans. (c) Among all SF6, Al2Cl6 doesn’t undergo hydrolysis, In SF6S is in high oxidation state (VI) & as F is bad leaving group) nucleophilic atack of water is not possible while Al2Cl6 is dimeric structure therefore does’t undergo hydrolysis. In case of BCl3 SiCl4 & PCl5 water can attack on central atom & undergoes rapid hydrolysis. 15. The reaction of solid XeF 2 with AsF6 in 1:1 ratio affords (a) Al 2Cl6 , SiMe4 (b) BCl3 , SF6 , SiCl 4 (c) XeF ASF6 [GATE 2011] (d) Xe2 F3 AsF6 Ans. (c) The reaction of solid XeF2 with AsF5 in 1 : 1 ratio affords XeF and AsF6 Hint : BF3, AsF5 and SbF5 are the F- acceptor & it leads to the formation nof a lewis base due to electron withdrawing nitro group while in case of (B) & (D) it leads to steric crowding in adduct. [5] Chemistry CHAPTER-1 MAIN GROUP ELEMENTS HYDROGEN AND THE HYDRIDES: ELECTRONIC STRUCTURE : Hydrogen has the simple atomic structure of all the elements, and consists of a nucleus containing one proton with a charge +1 and one orbital electron. The electronic structure may be written as 1s1. 1. By forming an electron pair (covalent) bond with another atom Non-metals typically form this type of bond with hydrogen, for example H 2 , H 2O, HCl gasor CH 4 and many metals do so too. 2. By losing an electron to form H+. A proton is extremely small (radius approximately 1.5 105 Å for most atoms). Because H+ is so small, it has a very high polarizing power, and therefore distorts the electron cloud on other atoms. Thus protons are always associated with other atoms or molecules. For example, in water or aqueous solutions of HCl and H2SO4 , protons exist as H 3O , H 9O4 H H 2O n ions. Free protons do not exist under ‘normal conditions’, though they are found in low pressure gaseous beams, for example in a mass spectrometer. 3. By gaining an electron to form HCrystalline solids such as LiH contain the H ion and are formed by highly electropositive metals (all of Group 1, and some of Group2) However, H ions are uncommon. Since hydrogen has an electro negativity of 2.1, it may use any of the three methods, but the most common way is forming covalent bonds. PREPARATION OF HYDROGEN Hydrogen is manufactured on a large scale by a variety of methods. 1. Hydrogen is made cheaply, and in large amounts, by passing steam over red hot coke. The product is water gas, which is a mixture of CO and H2. This is an important industrial fuel since it is easy to make and it burns, evolving a lot of heat. C C H 2O 1000 CO H 2 o water gas CO H 2 O2 CO2 H 2O heat It is difficult to obtain pure H 2 from water gas, since CO is difficult to remove. The CO may be liquefied at a low temperature under pressure, thus separating it from H 2 .Alternatively the gas mixture can be mixed with steam, cooled to 4000oC and passed over iron oxide in a shift converter, giving H 2 and CO2 . The CO2 so formed is easily removed either by dissolving in water under pressure, or reacting with thus giving H 2 gas K 2CO3 solution, giving KHCO3 , and H 2O CO H 2 2 H 2 CO2 450o C water gas Fe 2O3 2. Hydrogen is also made in large amounts by the steam reformer process. The hydrogen produced in this way is used in the Haber process to make NH 3 and for hardening oils . Light hydrocarbons such as methane are mixed with steam and passed over a nickel catalyst at 800900oC. These hydrocarbons are present in natural gas, and are also produced at oil refineries when ‘cracking’ hydrocarbons. CH 4 H 2O CO 3H 2 CH 4 2 H 2O CO2 4 H 2 [6] Chemistry The gas emerging from the reformer contains CO, CO2 .H 2 and excess steam. The gas mixture is mixed with more steam, cooled to 400oC and passed into a shift converter. This contains an iron/ copper catalyst and CO is converted into CO2. CO H 2O CO2 H 2 3. Pure hydrogen (99.9%pure) is made by electrolysis of water or solutions of NaOH or KOH. This is the most expensive method. Water does not conduct electricity very well. So it is usual to electrolyse aqueous solutions of NaOH or KOH in a cell with nickel anodes and iron cathodes. The gases produced in the anode and cathode compartments must be kept separate. 2OH H 2O ½O2 2e 2H 2O 2e 2OH H 2 Anode Cathode Overall H 2O 2OH H 2 4. The laboratory preparation is the reaction of dilute acids with metals. Zn H 2 SO4 ZnSO4 H 2 2 Al 2 NaOH 6 H 2O 2 NaAl OH 4 3H 2 PROPERTIES OF MOLECULAR HYDROGEN Hydrogen is the lightest gas known, and because of its low density, it is used instead of helium to fill balloons for me meteorology. It is colourless, odourless and almost insoluble in water, Hydrogen forms diatomic molecules H 2 , and the twos atoms are joined by a very strong covalent bond (bond 1 energy 435.9 kJ mol ) . Hydrogen is not very reactive under normal conditions. The lack of reactivity is due to kinetics rather than thermodynamics, and relates to the strength of the H H bond to produce atoms of hydrogen. This requires 435.9kJ mol 1 .hence there is a high activation energy to such reactions . Consequently many reactions are slow or require high temperatures, or catalysts (often transition metals). Many important reactions of hydrogen involve heterogeneous catalysis, where the catalyst first reacts with H 2 and either breaks or weakens the H H bond, and thus lowers the activation energy, Examples include : 1. The Haber process for the manufacture of NH 3 from N 2 and H 2 using a catalyst of activated Fe at 380 4500 C and 200 atmospheres pressure. 2. The hydrogenation of a variety of unsaturated organic compounds, (including the hardening of oils), using finely divided Ni, Pd or Pt as catalysts. 3. The production of methanol by reducing CO with H 2 over a Cu / Zn catalyst at 300oC. Thus hydrogen will react directly with most elements under the appropriate conditions. Hydrogen burns in air or dioxygen, forming water, and liberates a large amount of energy. This is used in the oxy-hydrogen flame for welding and cutting metals . Temperatures of almost 3000oC can be attained. Care should be taken with these gases since mixtures of H 2 and O2 close to a 2 : 1 ratio are often explosive. 2H 2 O2 2H 2O H 485kJ mol 1 Hydrogen reacts with the halogens. The reaction with fluorine is violent, even at low temperatures. The reaction with chlorine is slow in the dark, but the reaction is catalysed by light (photocatalysis), and becomes faster in daylight, and explosive in sunlight. Direct combination of the elements is used to produce HCl. H 2 F2 2 HF H 2 Cl2 2 HCl A number of metals react with H2, forming hydrides. The reaction are not violent, and usually require a high temperature. These are described in a later section [7] Chemistry Large quantities of H2 are used in the industrial production of ammonia by the Haber process. The reaction is reversible, and the formation of NH3 is favoured by high pressure, the presence of a catalyst (Fe), and a low temperature . In practice a high temperature of 380450oand a pressure of 200 atmosphere are used to get a reasonable conversion in a reasonable time. N 2 3H 2 g 298k 33.4Kj mol 1 Large amounts of H2 are used for hydrogenation reactions, in which hydrogen is added to a double bond in an organic compound. An important example is the hardening of fats and ioils. Unsaturated fatty acids acids which have higher melting points. By removing double bonds in the carbon chain in this way, edible oils which are liquid at room temperature may be converted into fats which are solid at room temperature. The reason for doing this is that solid fats are more useful than oils, for example in the manufacture of margarine. CH 3.CH 2 n .CH CH .COOH H 2 CH 3.CH 2 n .CH 2 .COOH Hydrogen is also used to reduce nitrobenzene to aniline (dyestuffs industry), and in the catalytic reduction of benzene (the first step in the production of nylon-66). It also reacts with CO to form methyl alcohol. CO 2H 2 catalyst CH 3OH The hydrogen molecule is veryd stable, and has little tendency to dissociate at normal temperatures, since the dissociation reaction is highly endothermic. H 2 2H H 435.9kJ mol 1 However, at high temperatures, inm an electric arc, or under ultraviolet light, H2 does dissociate. The atomic hydrogen produced exists for less than half a second, after which it recombines to give molecular hydrogen Atomic hydrogen is a strong reducing agent, and is commonly prepared in solution by moons of a zinc-copper couple or a mercury-aluminium couple. ORTHO AND PARA HYDROGEN The hydrogen molecule H2 exists in two different forms known as ortho. And para hydrogen. The nucleus of an atom has nuclear spin, in a similar way to electrons having a spin. In the H 2 molecule, the two nuclei may be spinning in either the same direction, or in opposite directions. This gives rise to spin isomerism, that is two different forms of H2 may exist. These are called ortho and para hydrogen. Spin isomerism is also found in other symmetrical molecules whose nuclei have spin momenta, e.g. D2,N2, F2, Cl2, There are considerable differences between the physical properties(e.g. boiling points, specific heats and thermal conductivities) of the ortho and para forms, because of differences in their internal energy. There are also differences in the band spectra of trhe ortho and para forms of H2. The para form has the lower energy, and at absolute zero the gas contain 100%of the para form. As the temperature is raised, some of the para form changes into the ortho form. At high temperatures the gas contains about 75% ortho hydrogen. (a) (b) Ortho and para hydrogen,(a) ortho, parallel spins;(b) para, opposite Para hydrogen is usually prepared by passing a mixture of the two forms of hydrogen through a tube packed with charcoal cooled to liquid air room temperature in a glass vessel, because the ortho-para conversion is slow in the absence of catalysts. Suitable catalysts include activated charcoal, atomic [8] Chemistry hydrogen, metals such as Fe, Ni, Pt and W and paramagnetic substances or ions (which contain unpaired electrons) such as O2, NO, NO2, Co2+and Cr2O3. Electronic Structure Group 1 elements all have one valency electron in their outer orbital-an s electron which occupies a spherical orbital. Ignoring the filled inner shells the electronic structures may be written: 2s1, 3s1, 4s1, 5s1, 6s1 and 7s1. The single valence electron is a long distance from the nucleus, is only weakly held and is readily removed. In contrast the remaining electrons are closer to the nucleus, more tightly held, and are removed only with great difficulty. Because of similarities in the electronic structures of these elements, many similarities in chemical behaviour would be expected. Electronic Structure Element Atomic Symbol Electronic Structure Number Lithium 3 Li 1s2 2s1 or [He] 2s1 2 2 6 1 Sodium 11 Na 1s 2s 2p 3s or [Ne] 3s1 Potassium 19 K 1s2 2s2 2p6 3s2 3p6 4s1 or [Ar] 4s1 2 2 6 2 6 10 2 6 1 Rubidium 37 Rb 1s 2s 2p 3s 3p 3d 4s 4p 5s or [Kr] 5s1 2 2 6 2 6 10 2 6 10 2 6 1 Caesium 35 Cs 1s 2s 2p 3s 3p 3d 4s 4p 4d 5s 5p 6s or [Xe] 6s1 Francium 87 Fr or [Rn] 7s1 Size of Atoms and Ions Group 1 atoms are the largest in their horizontal periods in the periodic table. When the outer electron is removed to give a positive ion, the size decreases considerably. There are two reasons for this. 1. The outermost shell of electrons has been completely removed. 2. Having removed an electron, the positive charge on the nucleus is now greater than the charge on the remaining electrons, so that each of the remaining electrons is attracted more strongly towards the nucleus. This reduces the size further. Positive ions are always smaller than the parent atom. Even so, the ions are very large and they increase in size from Li+ to Fr+ as extra shells of electrons are added. The Li+ is much smaller than the other ions. For this reason, Li only mixes with Na above o 380 C and it is immiscible with the metals K, Rb and Cs, even when molten; nor will Li form substitutional alloys with them. In contrast the other metals Na, K, Rb and Cs are miscible with each other in all proportions. Density The atoms are large, so Group 1 elements have remarkably low densities. Lithium metal is only about half as dense as water, whilst sodium and potassium are slightly less dense than water. It is unusual for metals to have low densities, and in contrast most of the transition metals have densities greater than 5g cm-3 for example iron 7.9 g cm-3, mercury 13.6 g cm-3, and osmium and iridium (the two most dense elements) 22.57 and 22.61 g cm-3 respectively. Table: Size and density Ionic radius M+ Density Metallic radius( A ) (gcm-3) six-coordinate( A ) Li 1.52 0.76 0.54 Na 1.86 1.02 0.97 K 2.27 1.38 0.86 Rb 2.48 1.52 1.53 Cs 2.65 1.67 1.90 Ionization Energy The first ionization energies for the atoms in this group are appreciably lower than those for any other group in the periodic table. [9] Chemistry The second ionization energy-that is the energy to remove a second electron from the atoms-is extremely high. The second ionization energy is always larger than the first, often by a factor of two, because it involves removing an electron from a smaller positive ion, rather than from a larger neutral atom. The difference between first and second ionization energies is much larger in this case since in addition it corresponds to removing an electron from a closed shell.A second electron is never removed under normal conditions, as the energy required is greater than that needed to ionize the noble gases. The elements commonly form M* ions. Ionization energies Li Na K Rb Cs First ionization energy (kJ mol-1) 520.1 495.7 418.6 402.9 375.6 Second ionization energy (kJ mol-1) 7296 4563 3069 2650 2420 Electronegativity and Bond Type The electronegativity values for the elements in this group are very small-in fact the smallest values of any element. Thus when these elements react with other elements to form compounds, a large electronegativity difference between the two atoms is probable, and ionic bonds are formed. Na electronegativity 0.9 Cl elect ronegativity 3.0 Electronegativity difference 2.1 Electronegativity values Pauling's electronegativity Li 1.0 Na 0.9 K 0.8 Rb 0.8 Cs 0.7 An electronegativity difference of approximately 1.7-1.8corresponds to 50% ionic character. TRhe value 2.1 exceeds this, so the bonding in NaCl is predominantly ionic. Similar arguments apply to other compounds: for example, the electronegativity difference in Lif is 3.0 and in KBr is 2.0 , and both compounds are ionic. The chemistry of the alkali metals is largely that of their ions. Born-Haber Cycle:Enthalpy changes in the Formation of Ionic Compounds When elements react to form compounds, G (the free energy of formation) is negative. For a reaction to proceed spontaneously, the free energy of the products must be lower than that of the reactants. Usually the energy changes are measured as enthalpy values H, and G is related to H by the equation. G = H - TS In many cases enthalpy values are used instead of free energy values, and the two are almost the same if the term TS is small, At room temperature T is almost 300K, so G and H are only similar when the change in entropy S is very small. Entropy changes are large when there is a change in physical state, e.g. solid to liquid, or liquid to gas, but otherwise entropy changes are usually small. There is no direct way of obtaining electron affinity values, and these have been calculated from this type of energy cycle. Hess’s law states that the energy change occurring during a reaction depends only on the energy of the initial reactants and the energy of the final products and not on the reaction mechanism, or the route taken. Thus, by Hess’s law, the energy change for the reaction of solid sodium and [10] Chemistry chlorine gas to from a sodium chloride crystal by the direct route (measured as the enthalpy of formation) must be the same as the sum of all the energy changes going round the cycle by the long route, i.e. by producing first gaseous almost of the all the energy changes going round the cycle by the long route, i.e. by producing first gaseous atoms of the elements, then gaseous ions, and finally packing these to give the crystalline solid. This may be expressed elements, then gaseous ions, and finally packing these to give the crystalline solid. This may be expressed as : Hf = +Hs + 1+ ½ Hd + E + U All the halides MCI have negative enthalpies of formation, which indicates that thermodynamically (that is in terms of energy) it is feasible to form the compounds MCI from the elements. 1. The most negative enthalpies of formation occur with t1e fluorides. For any given metal, the values decrease in the sequence fluoride > chloride > bromide > iodide. Thus the fluorides are the most stable, and the iodides the least stable. 2. The enthalpies of formation for the chlorides, bromides and iodides become more negative on descending the group. This trend is observed with most salts, but the opposite trend is found in the fluorides. Ionic compounds may also be formed in solution, when a similar cycle of energy changes must be considered, but the hydration energies of the positive and negative ions must be substituted for the lattice energy. Structures of Metals, Hardness and Cohesive Energy At normal temperatures all the Group 1 metals adopt a body-centred cubic type of lattice with a coordination number of 8. However, at very low temperatures lithium forms hexagonal close-packed structure with a coordination number of 12. The metals are all very soft, and can be cut quite easily with a knife. Lithium is harder than the others, but is softer than lead. The cohesive energy is the force holding the atoms or ions together in the solid. (This is the same in magnitude, but the opposite in sign, to the enthalpy of atomization, which is the energy required to break the solid up into gaseous atoms.) The cohesive energies of Group l-metals are about half of those for Group 2, and one third of those for Group 13 elements. The magnitude of the cohesive energy determines the hardness, and it depends on the number of electrons that can participate in bonding and on the strength of the bonds formed. The softness, low cohesive energy and weak bonding in Group 1 elements are consequences of these metals having only one valency electron which can participate in bonding (compared with two or more electrons in most other metals), and of the large size and diffuse nature of the outer bonding electron. The atoms become larger on descending the group from lithium to caesium, so the bonds are weaker, the cohesive energy decreases and the softness of the metals increases. Table: Cohesive energy Li Na Cohesive energy (enthalpy of atomization) (kJ mol-1) 161 108 [11] Chemistry K Rb Cs 90 82 78 Melting and Boiling Point The generally low values for cohesive energy are reflected in the very low values of melting and boiling points in the group. The cohesive energy decreases down the group, and the melting points decrease corresponding. The melting points range from lithium 181oC to caesium 28.5oC. These are extremely low values for metals, and contrast with the melting points of the transition metals, most of which are above 1000oC. The melting point of lithium is nearly twice as high (in oC) as that for sodium. Melting and Boiling points Melting point Li Na K Rb Cs Boiling point (oC) 1347 881 766 688 705 181 98 63 39 28.5 Flame Coloration: The colour actually arises from electronic transitions in short-lived species which are formed momentarily in the flame. The flame is rich in electrons, and in the case of soldium the ions are temporaliy reduced to atoms. Na Na e The sodium D-line (which is actually a doublet at 589.0 nm and 589.6nm)arises from the electronic transition 3s1 3 p1 in sodium atoms formed in the flame. Colour Li Na K Rb Cs Wavelength(nm) 670 589 766 780 455 Crimson Yellow Lilac Red-violet Blue Flame colours and wavelengths Lattice Energies: Salts of alkali metals consist of cations and anions only and are, therefore, called ionic solids. The lattice energy of an ionic solid is defined as the energy relased when cations and anions in gaseous state are brought from infinity to their respective crystal sites to form one mole of the ionic crystal. Li g Cl g Li Cl s ; H Lattice energy of Li Cl Cs g Cl g Cs Cl s ; H Lattice energy of Cs Cl Reaction with water Group 1 metals all react with water, liberating hydrogen and forming the hydroxides. The reaction becomes increasingly violent on descending the group. Thus lithium reacts gently, sodium melts on the surface of the water and the molten metal skates about vigorously and may catch fire and potassium melts and always catches fire. 2Li + 2H2O → 2LiOH + H2 2Na + 2H2O → 2NaOH + H2 [12] Chemistry 2K + 2H2O → 2KOH + H2 The standard electrode potentials Eo are Li+ | Li = -3.05 volts, Na+ | Na = -2.71, K+ | K = -2.93, Rb | Rb = 2.92, Cs+ | Cs = -2.92. Lithium has the most negative standard electrode potential of any element in the periodic table, largely because of its high hydration energy. Standard electrode potentials Eo and Gibbs free energy G are related by the equation. G nFE o Where n is the number of electrons removed from the metal to produce the ion, and F is the Faraday constant. The reaction Li+ + e → Li has the largest negative Eo value, and hence the largest positive G value. Thus the reaction does not occur. However, the reverse reaction Li+ → Li+ + e has a large negative value of G, so lithium liberates more energy than the other metals when it reacts with water. In view of this it is at first sight rather surprising that lithium reacts gently with water, whereas potassium, which liberates less energy, reacts violently and catches fire. The explanation lies in the kinetics (that is the rate at which the reaction proceeds) rather than in the thermodynamics (that is the total amount of energy liberated). Potassium has a low melting point, and the heat of reaction is sufficient to make it melt, or even vaporize. The molten metal spreads out, and exposes a larger surface to the water, so it reacts even faster, gets even hotter and catches fire. + Reaction with dinitrogen Lithium is the only element in the group that reacts with dinitrogen to form a nitride. Lithium nitride, Li3N, is ionic (3Li+ and N3-), and is ruby red. Two reactions of the nitride are of interest. First, on heating to a high temperature it decomposes to the elements, and second, it reacts with water, giving ammonia. 2Li3N 6Li + N2 Li3N + 3H2O → 3LiOH + NH3 OXIDES, HYDROXIDES, PEROXIDES AND SUPEROXIDES Reaction with air The metals all burn in air to form oxides, though the product varies depending on the metal. Lithium forms the monoxide Li2O (and some peroxide Li2O2), sodium forms the peroxide Na2O2 (and some monoxide Na2O), and the others form superoxides of the type MO2. All five metals can be induced to form the normal oxide, peroxide or superoxide by dissolving the metal in liquid ammonia and bubbling in the, appropriate amount of dioxygen. Normal oxides - monoxides The monoxides are ionic, for example 2Li+ and O2-Li2O and Na2O are pure white solids as expected, but surprisingly K2O is pale yellow, Rb2O is bright yellow and Cs2O is orange. Metallix oxides are usually basic. The typical oxides M2O are strongly basic oxides, and they react with water, forming strong bases. Li2O + H2O → 2LiOH Na2O + H2O → 2NaOH K2O + H2O → 2KOH The crystal structures of Li2O, Na2O, K2O and Rb2O are anti-fluorite structures. The antifluorite structure is like that for fluorite CaF2, except that the positions of the positive and negative ions are interchanged. Thus Li+ fill the sites occupied by F- and O2- fill sites occupied by Ca2+. Cs2O has an anti-CdCl2 layer structure. [13] Chemistry Hydroxides Sodium hydroxide NaOH is often called caustic soda, and potassium hydroxide is called caustic potash, because of their corrosive properties (for example on glass or on skin). These caustic alkalis are the strongest bases known in aqueous solution. The hydroxides of Na. K, Rb and Cs are very soluble in water, but LiOH is much less soluble. Solubility of Group 1 hydroxides Element Solubility (g/100g H O) Li 13.0 Na 108.3 K 112.8 Rb 197.6 Cs 385.6 At 25oC a saturated solution of NaOH is about 27 molar, whilst saturated LiOH is only about 5 molar. The bases react with acids to form salts and water, and are used for many neutralizations. The bases also react with CO2 even traces in the air, forming the carbonate. LiOH is used to absorb carbon dioxide in closed environments such as space capsules (where its light weight is an advantage in reducing the launching weight). 2NaOH + CO2 → Na2CO3 + H2O They also react with the amphoteric oxides, Al2O3 forming aluminates, SiO2 (or glass) forming silicates, SnO2 forming stannates, PbO2 forming plumbates and ZnO forming zincates. The bases liberate ammonia from both ammonium salts and coordination complexes where ammonia is attached to a transition metal ion (ammine complexes). NaOH + NH4Cl → Nh3 + NaCl + H2O 6NaOH + 2[Co(NH3)6] Cl3 → 12Nh3 + Co2O3 + 3NaCl + 3H2O Hexamine cobalt(III) chloride NaOH reacts with H2S to form sulphides S 2 , and hydrogen sulphides SH-, and it is used to remove mercaptans from petroleum products. NaOH + H2S → NaSH → Na2S The hydroxides react with alcohols, forming alkoxides. NaOH + EtOH → NaOEt + H2O Sodium ethoxide KOH resembles NaOH in all its reactions, but as KOH is much more expensive it is seldom used. However, KOH is much more soluble in alcohol, thus producing Oc2H5- ions by the equilibrium. C2H5OH + OH- OC2H5- + H2O Peroxides and superoxides The peroxides all contain the [−O−O−]2- ion. They are diamagnetic (all the electrons are paired), and are oxidizing agents. They may be regarded as salts of the dibasic acid H2O2 and they react with water and acid giving hydrogen peroxide H2O2. Na2O2 2H 2O 2 NaOH H 2O2 Na2O2 is pale yellow in colour: It is used industrially for bleaching wood pulp, paper and fabrics such as cotton and linen. It is a powerful oxidant, and many of its reactions are dangerously violent, particularly with materials that are reducing agents such as aluminium powder, charcoal, sulphur and many organic liquids. Because it reacts with CO2 in the air it has been used to purify the air in submarines and confined spaces, as it both removes CO 2 and produces O2. Potassium superoxide KO2 is even better for this purpose. Some typical reactions are. Na2O2 + Al → Al2O3 Na2O2 + Cr3+ → CrO24Na2O2 + Co → Na2CO3 2Na2O2 + 2Co2 → Na2CO3 + O2 [14] Chemistry The industrial process for forming sodium peroxide is a two-stage reaction in the presence of excess air. 1 2 Na O2 Na2O 2 1 Na2O O2 Na2O2 2 The superoxides contain the ion [O2]-, which has a unpaid electron and hence they are paramagnetic and are all coloured (LiO2 and NaO2 yellow, KO2 orange, RbO2 brown and CsO2 orange). The normal oxides contain O2- ion, the peroxides contain O22- ion and superoxides contain O-2 ion. The peroxides and superoxides become more stable with increase in atomic number of the alkali metal. The formation and stability of these oxides can be explained on the basis of their lattice energy Li+ ion being a small ion has a strong positive field around it and can stabilize only a small anion, O 2-, whereas Na+ ion being a large cation can stabilize a large anion and so on. Thus, smaller cation can stabilize smaller anion and larger cation can stabilize larger anion. The higher oxides, viz., peroxides and superoxides are important oxidizing agents. They react with dilute acids forming hydrogen peroxide and oxygen, respectively. Na2O2 + H2SO4 → Na2SO4 + H2O2 4KO2 + 2H2SO4 → 2K2SO4 + 2H2O + 3O2 Superoxides are even stronger oxidizing agents than peroxides, and give both H2O2 andO2 with either water or acids. 1 KO2 + 2H2O → KOH + H2O2 + O2 2 KO2 is used in space capsules, submarines, and breathing masks, because it both produces dioxygen and removes carbon dioxide. Both functions are important in life support systems. 4KO2 + 2CO2 → 2K2CO3 + 3O2 4KO2 + 4CO2 + 2H2O 4KHCO3 + 3O2 Sodium superoxide cannot be prepares by burning the metal in dioxygen at atmospheric pressure, but it is made commercially and in good yields by reacting sodium peroxide with dioxygen at a high temperature and pressure (450oC and 300 atmospheres) in a stainless steel bomb. Na2O2 + O2 → 2 NaO2 Sodium Bicarbonate, NaHCO3 NaHCO3 can be used on its own to make cakes or bread 'rise' since it decomposes between 50oC and 100oC giving bubbles of CO2. 2NaHCO3 Na2CO3 + H2O + CO2 Baking powder is more commonly used, and contains NaHCO3, Ca(H2PO4)2 and starch. The Ca(H2PO4)2 is acidic and when moistened it reacts with NaHCO, giving CO2. The starch is a filler. An improved combination baking powder contains about 40% starch, 30% NaHCO3, 20% NaAl(So4)2 and 10% Ca(H2PO4)2. The NaAl(So4)2 slows the reaction down so the CO2 is given off more slowly. Oxosalts-Carbonates, Bicarbonates, Nitrates and Nitrites Group 1 metals are highly electropositive and thus form very strong bases, and have quite stable oxosalts. [15] Chemistry The carbonates are remarkably stable. And will melt before they eventually decompose into oxides at temperatures above 1000oC. Li2CO3 is considerably less stable and decomposes more readily. Because Group 1 metals are so strongly basic, they also form solid bicarbonates (also called hydrogen carbonates). No other metals form solid bicarbonates, though NH4HCO3 also exists as a solid. Bicarbonates evolve carbon dioxide and turn into carbonates on gentle warming. This is one test for bicarbonates in qualitative analysis. The crystal structures of NaHCO3 and KHCO3 both show hydrogen bonding, but are different. In NaHCO3 the HCO-3 a ions are linked into an infinite chain, whilst in KHCO3 a dimeric anion is formed. Lithium is exceptional in that it does not form a solid bicarbonate, though LiHCO3 can exist in solution. All the carbonates and bicarbonates are soluble in water. Over 50000 tonnes of Li2CO3 are produced annually. Most of it is added as an impurity to Al2O3 to lower its melting point in the extraction of aluminium by electrolysis. Some is used to toughen glass (sodium in the glass is replaced by lithium). Na2CO3 is used as washing soda to soften water in hard water areas, and NaHCO3 is used as baking powder. The nitrates can all be prepared by the action of HNO3 on the corresponding carbonate or hydroxide, and they are all very soluble in water. LiNO3 is used for fireworks and red-coloured distress flares. Large deposits of NaNO3 are found in Chile, and are used a nitrogenous fertilizer. Solid LiNO3 and NaNO3 are deliquescent, and because of this KNO3 is used in preference to NaNO3 is gunpowder (gunpowder is a mixture of KNO3, sulphur and charcoal). KNO3 is usually obtained from synthetic nitric acid and K2CO3, but at one time it was made from NaNO3. 2HNO3 + K2CO3 → 2 KNO3 + CO2 + H2O NaNO3 + KCl → KNO3 + NaCl Group 1 nitrates are fairly low melting solids, and are amongst the most stable nitrates known. However, on strong heating they decompose into nitrites, and at higher temperatures to the oxide. LiNO3 decomposes more readily than the others, forming the oxide. 2NaNO3 2NaNO2 + O2 4NaNO3 2Na2O + 5O2 + 2N2 Halides and Polyhalids Since Li+ is the smallest ion in the group, it would be expected to form hydrated salts more readily than the other metals. LiCl, LiBr and Lil form trihydrates LiX.3H2O, but the other alkali metal halides form anhydrous crystals. All the halides adopt a NaCl type of structure with a coordination number of 6 except for CsCl, CsBr and CsI. The latter have a CsCl type of structure with a coordination number of 8. Rather more compounds adopt the NaCl type of structure than would be expected from the radius ratios of the ions r+ / r-, and the reason for this structure being adopted is that it gives the highest lattice energy. The alkali metal halides react with the halogens and interhalogen compounds forming ionic polyhalide compounds. Kl + l2 → K [l3] KBr + |C| → K [BrlCl] KF + BrF3 → K [BrF4] Hydrides Group 1 metals all react with hydrogen, forming ionic or salt-like hydrides M+H-. However, the ease with which they do so decreases from lithium to caesium. These hydrides contain the H- ion. It can be proved that H- ions exist because on electrolysis hydrogen is liberated at the anode. The hydrides react with water, liberating hydrogen, and lithium hydride is used as a source of hydrogen for military purposes and for filling meteorological balloons. LiH + H2O → LiOH + H2 [16] Chemistry Lithium also forms a complex hydride Li[AlH4], called lithium aluminium hydride, which is a useful reducing agent. It is made from lithium hydride in dry ether solution. 4LiH + AlCl3 → Li[AlH4] + 3LiCl Lithium aluminium hydride is ionic, and the [AlH4]- ion is tetrahedral. Li[AlH4] is a powerful reducing agent and is widely used in organic chemistry, as it reduces carbonyl compounds to alcohols. It reacts violently with water, so it is necessary to use absolutely dry organic solvents, for example ether which has been dried over sodium. Li[AlH4] will also reduce a number of inorganic compounds. BCL3 + Li[AlH4] → B2H6 diborane PCL3 + Li[AlH4] → PH3 phosphine SiCl4 + Li[AlH4] → SiH4 silane Sodium tetrahydridoborate (sodium borohydride) Na[BH4] is another hydride complex. It is ironic, comprising tetrahedral [BH4]- ions. It is best obtained by heating sodium hydride with trimethyl borate. 4NaH + B(OCH3)3 Na[BH4] + 3NaOCH3 Solubilty and Hydration All the simple salts dissolve in water, producing ions, and consequently the solutions conduct electricity. Since Li+ ions are small, it might be expected that solutions of lithium salts would conduct electricity better than solutions of the same concentration of sodium, potassium, rubidium or caesium salts. The small ions should migrate more easily towards the cathode, and thus conduct more than the larger ions. However, ionic mobility or conductivity measurements in aqueous solution give results in the opposite order Cs+ > Rb+ > K+ > Na+ > Li+. The reason for this apparent anomaly is that the ions are hydrated in solution. Since Li+ is very small, it is heavily hydrated. This makes the radius of the hydrated ion large, and hence it moves only slowly. In contrast, Cs + is the least hydrated, and the radius of the hydrated Cs+ ion is smaller than the radius of hydrated Li+, and hence hydrated Cs+ moves faster, and conducts electricity more readily. The size of the hydrated ions is an important factor affecting the passage of these ions through cell walls, It also explains their behaviour on cation exchange columns, where hydrate Li + ions are attached less strongly and hence eluted first. The decreases in hydration from Li+ to Cs+ is also shown in the crystalline salts, for nearly all lithium salts are hydrated, commonly as trihydrates. In these hydrated Li salts Li+ is coordinated to 6H2O, and the octahedra share faces, forming chains. Many sodium salts are hydrated, e.g. Na 2CO3. 10H2O, Na2CO. 7H2O and Na2CO3.H2O. Few potassium salts and no rubidium or caesium salts are hydrted. The simple salts are all soluble in water, and so in qualitative analysis these metals need to be precipitated as less common salts. Thus Na+ is precipitated by adding zinc (or copper) uranyl acetate solution and precipitating NaZn(UO2)(Ac)9. H2O sodium zinc uranyl acetate. K+ is precipitated by adding a solution of sodium cobaltinitrite and precipitating potassium cobaltinitrite K 3[Co(NO2)6] or by adding perchloric acid and precipitating potassium perchlorate KCIO4. Group 1 metals can be estimated gravimetrically, sodium as the uranylacetate and potassium, rubidium and caesium as tetraphenylborates. However, modern instrumental methods such as flame photometry and atomic absorption spectrometry are much quicker and easier to use and are now used in preference to gravimetric analysis. K Na3 CONO2 6 Na K3 CONO2 6 K NaClO4 Na KClO4 K NaBC6 H 5 4 Na K BC6 H 5 4 potassium cobaltinitrite potassium perchlorate potassium tetraphenylborate quantitative precipitate If a salt is insoluble its lattice energy is greater than the hydration energy. K[B(C 6H5)4] is insoluble because the hydration energy is very small as a result of the large size of its ions. [17] Chemistry The solubility of most of the salts of Group 1 elements in water decreases on descending the group. For a substance to dissolve the energy evolved when the ions are hydrated (hydration energy) must be larger than the energy required to break the crystal lattice (lattice energy). Conversely, if the solid is insoluble, the hydration energy is less than the lattice energy. The reason why the solubility of most Group 1 metals decreases on descending the group is that the lattice energy only changes slightly, but the free energy of hydration changes rather more. For example, the difference in lattice energy between NaCl and KCl is 67 kJ mol -1, and yet the difference in G (hydration) for Na+ and K+ is 76 kJ mol-1. Thus 1 is less soluble than NaCl. The Group 1 fluorides and carbonates are exceptional in that their solubilities increase rapidly on descending the group. The reason for this is that their lattice energies change more than the hydration energies on descending the group. The lattice energy depends on electrostatic attraction between ions, and is proportional to the distance between the ions, that is proportional, to I/(r +/r-). It follows that the lattice energy will vary most when r- is small, that is with F-, and will vary least when r+ is large (with I-). The weight if solute dissolving does not provide a very useful comparison of the solubilities, because the molecular weights differ. The easiest way to compare the number of ions is to compare the solubilities as molar quantities. Reactions with NH3 (gaseous): When ammonia gas is passed over alkali metals, the corresponding amides are formed. 2M 2 NH 3 2MNH 2 H 2 The alkali amides are also formed when the solutions of alkali metals in liquid NH3 are evaporated. Solubility in Mercury: The alkali metals dissolve readily in mercury forming what are called amalgams. The process is highly exothermic. Solubility in Liquid Ammonia: All the alkali metals dissolve in anhydrous liquid ammonia giving deep blue solutions. The colour depends with increase in concentration of the solution. The concentrated solutions are bronze in colour and posses metallic luster. The solutions are good conductors of electricity and are paramagnetic in nature. The paramagnetic character decreases with increase in concentration. The dissolution of the metal is accompanied by its ionization. M → M+ + eWhen the alkali metals are dissolved in liquid ammonia, the ammonia molecules solvate the cation as well as the electron, as represented below. M + NH3 (in excess) → M+ (NH3)x + e- (NH3)y The values of x and y depend upon the extent of salvation of the alkali metal ion and the electron, respectively. The exact mechanism for the salvation of electrons is not clear. It is argued, however, that the electron released by the alkali metal causes polarization of the charge of the electrons of the surrounding ammonia molecules causes a 'hole' in which the released electron is trapped. It is the ammoniated electron which is responsible for the blue colour of the solution. The electrical conductivity is due to the ammoniated cation as well as the ammoniated electron. The dilute solutions are paramagnetic because they contain free electrons. The Decrease in paramagnetic character with increase in concentration suggests that in concentrated solutions the ammoniated electrons associate to form electron pairs. 2e- (NH3)y → [e- (NH3)y]2 If blue colour is allowed to stand, the colour slowly fades until it disappears owing to formation of metalamide. At concentration above 3M, solutions are copper-bronze coloured and have a metallic luster because metal clusters are formed. In the presence of impurities or catalysts such as Fe, the alkalie metals react with NH3 to form metal amide and hydrogen. M + NH3 → MNH2 + H2 If all impurities or catalysts are absent, the group 1 metals dissolve in liquid ammonia. The alkali metal solutions in liquid ammonia are stable and becompose only slowly liberating hydrogen. The presence of transition metals (such as Fe, Pt, Zn, etc.) even in traces, catalyses the reaction. [18] Chemistry Due to the presence of the free or ammoniated electrons, these solutions behave as strong reducing agents as is evident from the following reactions. These solutions reduce aerial oxygen to superoxide and peroxide ions without causing cleavage of O–O bond. These solutions reduce metal ions to their unusually low oxidation states. K2[Ni(CN)4] + 2K Mn2(CO)10 + 2K Fe(CO)5 + 2Na K4[Ni(CN)4] 2K[Mn(CO)5] Na2[Fe(CO)4] +CO (Ni is in 0 oxidation state) (Mn is in -1 oxidation state) (Fe is in –2 oxidation state) Ionic Conductance: The alkali metal ions in aqueous solution have high electrical conductance. The electrical conductance of Li+ ion, which is the smallest of alkali metal ions, is expected to be the highest. Since ionic size increases on moving from Li+ to Cs. the ionic conductance is expected to decrease in the order. Li+ > Na+ > K+ > Rb+ > Cs+ However, we find by experiment that ionic conductance of Cs+ is the highest and that of Li+ is the lowest. Thus, ionic conductance actually decreases in the order. Cs+ > Rb+ > K+ > Na+ > Li+ This is evidently the reverse of what we expect. This anomaly arises from the fact that alkali metal ions are hydrated to different extents. The Li+ ion is hydrated to the maximum extent due to its high charge/size ratio compared to other alkali metal ions. The extent of hydration decreases on moving from Li+ to Cs+. Thus, Cs+ ion is the least hydrated. As a result of the variation in the extent of hydration, Li+ ion due to heavy hydration is the biggest while Cs+ ion due to much smaller hydration is the smallest. Thus, the ionic size in aqueous solutions, instead of increasing, actually decreases on moving from Li+ to Cs+ as a result of which the ionic mobility increases as we move from Li+ to Cs+. Hence, the ionic conductance, instead of decreasing, actually increases on moving from Li+ to Cs+. The ionic radii, radii of hydrated ions and ionic conductances of different alkali metal ions at infinite dilution are given in following table. Role of Na+ and K+ ions in Biological Systems: Sodium cation is a major cation of extracellular fluids of animals including human beings which is known to activate certain enzymes in the animal body. Sodium ions are relatively harmless except when present in excess amounts in which case they may cause hypertension. The saline water containing excessive amounts of NaCl is harmful to plant and aquatic life because of the toxicity of Na+ ions. The potassium ion is essential to all organisms with the possible exception of blue green algae. It is a major cation of intracellular fluid of animal cells. It is moderately toxic to mammals but only when injected intravenously. The sodium and potassium ions, taken together, constitute what is known as Na+ - K+ pump or simply as 'sodium pump'. The functioning of this pump is actually a biological process which goes on occurring in each and every cell of all animals. It transfers Na+ ions from the intracellular fluids to the extra cellular fluids with the help of 'carrier proteins'. Simultaneously it transfers K+ ions from the extra cellular fluids to the intracellular fluids. Since each operation of the pump (i.e., each cycle of this biological process) pumps out larger number of sodium ions from the cell than the number of K + ions that it pumps into the cell, the interior of the cell acquires an excess negative charge and the exterior of the cell acquires an excess positive charge. This results in the development of electrical potential gradient across the cell membrane which is responsible for the transmission of nerve signals in animals. The Na+ – K+ pump also maintains the volume of the cell. Without the existence of the Na+ – K+ pump in the cell, the latter would have swelled in volume and ultimately bursted. [19] Chemistry