Classification of Elements and Periodicity in Properties PDF

CLASSIFICATION OF ELEMENTS AND PERIODICITY IN PROPER TIES 87 Table 3.7 Electron Gain Enthalpies* / (kJ mol–1 ) of Some Main Group Elements Group 1 ∆ eg H Group 16 ∆ eg H Group 17 ∆ eg...

CLASSIFICATION OF ELEMENTS AND PERIODICITY IN PROPER TIES 87 Table 3.7 Electron Gain Enthalpies* / (kJ mol–1 ) of Some Main Group Elements Group 1 ∆ eg H Group 16 ∆ eg H Group 17 ∆ eg H Group 0 ∆ eg H H – 73 He + 48 Li – 60 O – 141 F – 328 Ne + 116 Na – 53 S – 200 Cl – 349 Ar + 96 K – 48 Se – 195 Br – 325 Kr + 96 d Rb – 47 Te – 190 I – 295 Xe + 77 he Cs – 46 Po – 174 At – 270 Rn + 68 large positive electron gain enthalpies because Problem 3.7 the electron has to enter the next higher pu T principal quantum level leading to a very Which of the following will have the most is unstable electronic configuration. It may be negative electron gain enthalpy and which the least negative? re ER noted that electron gain enthalpies have large negative values toward the upper right of the P, S, Cl, F. bl periodic table preceding the noble gases. Explain your answer. The variation in electron gain enthalpies of elements is less systematic than for ionization Solution enthalpies. As a general rule, electron gain Electron gain enthalpy generally becomes be C enthalpy becomes more negative with increase more negative across a period as we move in the atomic number across a period. The from left to right. Within a group, electron effective nuclear charge increases from left to gain enthalpy becomes less negative down o N right across a period and consequently it will a group. However, adding an electron to be easier to add an electron to a smaller atom the 2p-orbital leads to greater repulsion since the added electron on an average would than adding an electron to the larger be closer to the positively charged nucleus. We 3p-orbital. Hence the element with most should also expect electron gain enthalpy to © negative electron gain enthalpy is chlorine; become less negative as we go down a group the one with the least negative electron because the size of the atom increases and the gain enthalpy is phosphorus. added electron would be farther from the nucleus. This is generally the case (Table 3.7). (e) Electronegativity However, electron gain enthalpy of O or F is less negative than that of the succeeding A qualitative measure of the ability of an atom element. This is because when an electron is in a chemical compound to attract shared added to O or F, the added electron goes to the electrons to itself is called electronegativity. smaller n = 2 quantum level and suffers Unlike ionization enthalpy and electron gain tt significant repulsion from the other electrons enthalpy, it is not a measureable quantity. present in this level. For the n = 3 quantum However, a number of numerical scales of level (S or Cl), the added electron occupies a electronegativity of elements viz., Pauling scale, larger region of space and the electron-electron Mulliken-Jaffe scale, Allred-Rochow scale have no repulsion is much less. been developed. The one which is the most * In many books, the negative of the enthalpy change for the process depicted in equation 3.3 is defined as the ELECTRON AFFINITY (Ae ) of the atom under consideration. If energy is released when an electron is added to an atom, the electron affinity is taken as positive, contrary to thermodynamic convention. If energy has to be supplied to add an electron to an atom, then the electron affinity of the atom is assigned a negative sign. However, electron affinity is defined as absolute zero and, therefore at any other temperature (T) heat capacities of the reactants and the products have to be taken into account in ∆egH = –Ae – 5/2 RT. 88 CHEMISTRY widely used is the Pauling scale. Linus Pauling, electrons and the nucleus increases as the an American scientist, in 1922 assigned atomic radius decreases in a period. The arbitrarily a value of 4.0 to fluorine, the element electronegativity also increases. On the same considered to have the greatest ability to attract account electronegativity values decrease with electrons. Approximate values for the the increase in atomic radii down a group. The electronegativity of a few elements are given in trend is similar to that of ionization enthalpy. Table 3.8(a) Knowing the relationship between The electronegativity of any given element electronegativity and atomic radius, can you d is not constant; it varies depending on the now visualise the relationship between element to which it is bound. Though it is not electronegativity and non-metallic properties? he a measurable quantity, it does provide a means of predicting the nature of force that holds a pair of atoms together – a relationship that you will pu T explore later. is Electronegativity generally re ER increases across a period from left to right (say from lithium to bl fluorine) and decrease down a group (say from fluorine to astatine) in the periodic table. How can these trends be explained? Can the be C electronegativity be related to atomic radii, which tend to decrease across each period from o N left to right, but increase down each group ? The attraction between the outer (or valence) Fig. 3.7 The periodic trends of elements in the periodic table © Table 3.8(a) Electronegativity Values (on Pauling scale) Across the Periods Atom (Period II) Li Be B C N O F Electronegativity 1.0 1.5 2.0 2.5 3.0 3.5 4.0 Atom (Period III) Na Mg Al Si P S Cl Electronegativity 0.9 1.2 1.5 1.8 2.1 2.5 3.0 Table 3.8(b) Electronegativity Values (on Pauling scale) Down a Family tt Atom Electronegativity Atom Electronegativity (Group I) Value (Group 17) Value no Li 1.0 F 4.0 Na 0.9 Cl 3.0 K 0.8 Br 2.8 Rb 0.8 I 2.5 Cs 0.7 At 2.2 CLASSIFICATION OF ELEMENTS AND PERIODICITY IN PROPER TIES 89 Non-metallic elements have strong tendency electronic configuration 2s 22p5, shares one to gain electrons. Therefore, electronegativity electron with oxygen in the OF2 molecule. Being is directly related to that non-metallic highest electronegative element, fluorine is properties of elements. It can be further given oxidation state –1. Since there are two extended to say that the electronegativity is fluorine atoms in this molecule, oxygen with 2 inversely related to the metallic properties of outer electronic configuration 2s 2p4 shares elements. Thus, the increase in two electrons with fluorine atoms and thereby electronegativities across a period is exhibits oxidation state +2. In Na2O, oxygen d accompanied by an increase in non-metallic being more electronegative accepts two properties (or decrease in metallic properties) electrons, one from each of the two sodium of elements. Similarly, the decrease in atoms and, thus, shows oxidation state –2. On he electronegativity down a group is accompanied the other hand sodium with electronic by a decrease in non-metallic properties (or configuration 3s 1 loses one electron to oxygen increase in metallic properties) of elements. and is given oxidation state +1. Thus, the pu T All these periodic trends are summarised oxidation state of an element in a particular is in figure 3.7. compound can be defined as the charge acquired by its atom on the basis of re ER 3.7.2 Periodic Trends in Chemical electronegative consideration from other atoms bl Properties in the molecule. Most of the trends in chemical properties of elements, such as diagonal relationships, inert Problem 3.8 pair effect, effects of lanthanoid contraction etc. Using the Periodic Table, predict the be C will be dealt with along the discussion of each formulas of compounds which might be group in later units. In this section we shall formed by the following pairs of elements; study the periodicity of the valence state shown (a) silicon and bromine (b) aluminium and o N by elements and the anomalous properties of sulphur. the second period elements (from lithium to fluorine). Solution (a) Periodicity of Valence or Oxidation (a) Silicon is group 14 element with a valence of 4; bromine belongs to the © States halogen family with a valence of 1. The valence is the most characteristic property Hence the formula of the compound of the elements and can be understood in terms formed would be SiBr4. of their electronic configurations. The valence of representative elements is usually (though (b) Aluminium belongs to group 13 with not necessarily) equal to the number of a valence of 3; sulphur belongs to electrons in the outermost orbitals and / or group 16 elements with a valence of equal to eight minus the number of outermost 2. Hence, the formula of the compound electrons as shown below. formed would be Al2S3. Nowadays the term oxidation state is tt Some periodic trends observed in the frequently used for valence. Consider the two valence of elements (hydrides and oxides) are oxygen containing compounds: OF2 and Na2O. shown in Table 3.9. Other such periodic trends The order of electronegativity of the three which occur in the chemical behaviour of the no elements involved in these compounds is F > elements are discussed elsewhere in this book. O > Na. Each of the atoms of fluorine, with outer Group 1 2 13 14 15 16 17 18 Number of valence 1 2 3 4 5 6 7 8 electron Valence 1 2 3 4 3,5 2,6 1,7 0,8 90 CHEMISTRY Table 3.9 Periodic Trends in Valence of Elements as shown by the Formulas of Their Compounds Group 1 2 13 14 15 16 17 Formula LiH B2 H6 CH4 NH3 H2O HF of hydride NaH CaH2 AlH3 SiH4 PH3 H2 S HCl KH GeH4 AsH3 H2 Se HBr d SnH4 SbH3 H2 Te HI Formula Li2O MgO B2 O3 CO2 N2O 3, N2 O5 – he of oxide Na2 O CaO Al2O3 SiO2 P4 O6, P4O 10 SO3 Cl2 O 7 K2O SrO Ga2 O3 GeO2 As2O 3, As2 O5 SeO3 – pu T BaO In2O3 SnO2 Sb2O 3, Sb2O 5 TeO3 – is PbO2 Bi2 O3 – – re ER bl There are many elements which exhibit variable following group i.e., magnesium and valence. This is particularly characteristic of aluminium, respectively. This sort of similarity transition elements and actinoids, which we is commonly referred to as diagonal shall study later. relationship in the periodic properties. be C (b) Anomalous Properties of Second Period What are the reasons for the different Elements chemical behaviour of the first member of a group of elements in the s- and p-blocks The first element of each of the groups 1 o N compared to that of the subsequent members (lithium) and 2 (beryllium) and groups 13-17 in the same group? The anomalous behaviour (boron to fluorine) differs in many respects from is attributed to their small size, large charge/ the other members of their respective group. radius ratio and high electronegativity of the For example, lithium unlike other alkali metals, elements. In addition, the first member of © and beryllium unlike other alkaline earth group has only four valence orbitals (2s and metals, form compounds with pronounced 2p) available for bonding, whereas the second covalent character; the other members of these member of the groups have nine valence groups predominantly form ionic compounds. orbitals (3s, 3p, 3d). As a consequence of this, In fact the behaviour of lithium and beryllium the maximum covalency of the first member of is more similar with the second element of the each group is 4 (e.g., boron can only form − Property Element [BF4 ] , whereas the other members of the groups can expand their Metallic radius M/ pm Li Be B valence shell to accommodate more tt 152 111 88 than four pairs of electrons e.g., 3− Na Mg Al aluminium forms [ AlF6 ] ). Furthermore, the first member of no 186 160 143 p-block elements displays greater Li Be ability to form pπ – p π multiple bonds + to itself (e.g., C = C, C ≡ C, N = N, Ionic radius M / pm 76 31 N ≡ Ν) and to other second period Na Mg elements (e.g., C = O, C = N, C ≡ N, N = O) compared to subsequent 102 72 members of the same group. CLASSIFICATION OF ELEMENTS AND PERIODICITY IN PROPER TIES 91 Problem 3.9 non-metallic character increases while moving from left to right across the period. The Are the oxidation state and covalency of chemical reactivity of an element can be best 2+ Al in [AlCl(H2 O)5] same ? shown by its reactions with oxygen and Solution halogens. Here, we shall consider the reaction of the elements with oxygen only. Elements on No. The oxidation state of Al is +3 and the two extremes of a period easily combine with covalency is 6. oxygen to form oxides. The normal oxide d formed by the element on extreme left is the 3.7.3 Periodic Trends and Chemical most basic (e.g., Na2O), whereas that formed Reactivity by the element on extreme right is the most he acidic (e.g., Cl2O 7). Oxides of elements in the We have observed the periodic trends in certain centre are amphoteric (e.g., Al2O 3, As2O3) or fundamental properties such as atomic and neutral (e.g., CO, NO, N2O). Amphoteric oxides ionic radii, ionization enthalpy, electron gain pu T behave as acidic with bases and as basic with enthalpy and valence. We know by now that is acids, whereas neutral oxides have no acidic the periodicity is related to electronic or basic properties. re ER configuration. That is, all chemical and physical properties are a manifestation of the bl Problem 3.10 electronic configuration of elements. We shall now try to explore relationships between these Show by a chemical reaction with water fundamental properties of elements with their that Na2O is a basic oxide and Cl2O7 is an chemical reactivity. acidic oxide. be C The atomic and ionic radii, as we know, Solution generally decrease in a period from left to right. Na2O with water forms a strong base As a consequence, the ionization enthalpies o N whereas Cl2O7 forms strong acid. generally increase (with some exceptions as outlined in section 3.7.1(a)) and electron gain Na2O + H2O → 2NaOH enthalpies become more negative across a period. In other words, the ionization enthalpy Cl2O 7 + H2O → 2HClO4 © of the extreme left element in a period is the least and the electron gain enthalpy of the Their basic or acidic nature can be element on the extreme right is the highest qualitatively tested with litmus paper. negative (note : noble gases having completely filled shells have rather positive electron gain Among transition metals (3d series), the change enthalpy values). This results into high in atomic radii is much smaller as compared chemical reactivity at the two extremes and the to those of representative elements across the lowest in the centre. Thus, the maximum period. The change in atomic radii is still chemical reactivity at the extreme left (among smaller among inner-transition metals (4f series). The ionization enthalpies are tt alkali metals) is exhibited by the loss of an electron leading to the formation of a cation intermediate between those of s- and p-blocks. and at the extreme right (among halogens) As a consequence, they are less electropositive shown by the gain of an electron forming an than group 1 and 2 metals. no anion. This property can be related with the In a group, the increase in atomic and ionic reducing and oxidizing behaviour of the radii with increase in atomic number generally elements which you will learn later. However, results in a gradual decrease in ionization here it can be directly related to the metallic enthalpies and a regular decrease (with and non-metallic character of elements. Thus, exception in some third period elements as the metallic character of an element, which is shown in section 3.7.1(d)) in electron gain highest at the extremely left decreases and the enthalpies in the case of main group elements.