UNIT 3 ( PAGE 2)

Successive Electron affinity
            Like the second and higher ionisation energies , the second and higher electron affinities are also possible. After addition of one electron, the atom becomes negatively charged and the second electron is added to a negatively charged ion. The addition of second electron is opposed by the coulombic force of repulsion and energy has to be supplied for the addition of the second electron. If an atom has spontaneous tendency, i.e., a positive tendency , to gain electron, then conventionally , its electron gain enthalpy is said to be negative and if the atom is reluctant to gain an electron, i.e., it has a negative tendency to gain an electron and is forced to accept it, its electron gain enthalpy is   positive. Thus in the case of oxygen, the first electron gain enthalpy is negative since 141 kJ is released when one mole of oxygen atoms get converted to O- ions. In other words oxygen atom has positive tendency to accept electron. However, the second electron gain enthalpy  is positive  since 770 kJ of energy has to be supplied to convert  1 mol of O-  ions to O2-  ions.
            Similarly in the case of sulphur, while first electron gain enthalpy is negative since 200 kJ of energy is released when        1 mole of S atoms get converted to S- ions and second electron gain enthalpy is positive since 590 kJ of energy has to be supplied to convert 1 mole of S- ions to S2-  ions.
            Thermodynamically , the energy released is given a negative sign and energy absorbed is given a positive sign. Accordingly, when a species has a positive electron affinity ,   DH , accompanying the addition of an electron to the species, is negative and if it has a negative electron affinity, DH is positive. Thus for the reaction Cl + e ®  Cl-  , while electron affinity is positive , while electron gain enthalpy is negative. In various calculations involving DH, the value of electron affinity of chlorine would be taken as -349 kJmol-1and not +349 kJmol-1.
Factors on which Electron Gain Enthalpy depends
            The important factors upon which electron gain enthalpy depends are briefly discussed.
i)  Atomic size :  As the size of an atom increases , the distance between its nucleus and the incoming electron also increases. Consequently, the incoming electron experiences less attraction towards the nucleus of the atom. Therefore , electron gain enthalpy becomes less negative down the group.

ii)    Nuclear charge :  With increase in the nuclear charge, force of attraction between the nucleus and and incoming electron increases and so is the value of electron gain enthalpy. Thus, the electron gain enthalpy becomes more negative with increase in nuclear charge.
iii) Symmetry of electronic configuration :  The symmetry of electronic configuration has very important role to play. The atoms with symmetrical configuration (having filled and half-filled orbitals in the same sub-shell) do not have any urge to take up extra electrons because their configuration will become unsymmetrical or less stable. In case these are made to accept electrons, energy will be needed and electron gain enthalpy will be positive. For example , noble gas elements have positive electron gain enthalpies.
Variation of Electron gain enthalpy
Along a period
As a general rule, electron gain enthalpy becomes more negative with increase in atomic number across a period. The effective nuclear charge increases as we go from left to right across a period and consequently it will be easier to add an electron to smaller atom since the added electron on an average would be closer to the positively charged nucleus.
The trends in electron gain enthalpy values within a period are irregular indicating that atomic size is not the only criterion for determining electron gain enthalpy. Thus electron gain enthalpy of Be is positive (+66 kJmol-1) is while that of Li is negative (-60 kJmol-1). Similarly , electron gain enthalpy of nitrogen is positive(+31 kJmol-1) while that of oxygen although atomic size of oxygen is less  is negative (-141 kJmol-1) . 
Along a group
We should also expect electron gain enthalpy to become less negative as we go down a group because the size of the atom increases and the added electron would be farther from the nucleus. This is generally the case (Table above). However, electron gain enthalpy of O or F is less than that of succeeding element. This is because when electron is added to O or F, the added electron goes to smaller n = 2 quantum level and suffers significant repulsion from the other electrons present in this level.  For the n = 3 quantum level ( S or Cl) , the added electron occupies a larger region of space and the electron-electron repulsion is much less.
Some typical trends in Electron Gain Enthalp
Halogens have high negative electron gain enthalpy. This is due to their strong tendency to gain an additional electron to change into s2p6 configuration. The electron gain enthalpy become less negative from Cl ® Br ® Ii.e., on moving down the group. However, the electron gain enthalpy of fluorine is unexpected less negative. It is probably due to small size of the atom(atomic radius 72 pm). The addition of extra electron produces high electron charge density in a relatively compact 2p sub-shell resulting in strong electron-electron repulsion. The repulsive forces between electrons imply low negative electron gain enthalpy value. Therefore the incoming electron is not accepted with the same ease as in the case of chlorine which the atomic size is  bigger (atomic radius 99 pm) and the electron crowding is comparatively less. As a result, the negative electron gain enthalpy of fluorine (-328 kJmol-1) is less as compared to chlorine    (-349 kJmol-1)

Electron Gain Enthalpies of Halogens

            In the oxygen family (Group 16) the negative electron gain enthalpy of oxygen is also less than that of sulphur although it is expected to be more.
            Apart from the size of the atom, the electronic configuration also influence their electron gain enthalpy values considerably.
            The members of the noble gas family have highly symmetrical electronic configuration and have completely occupied orbitals. Therefore , there atoms have hardly any need or urge to take up extra electron. In case these are made to accept electrons under special conditions, enthalpy or energy will be needed and their electron gain enthalpies (DegH) will be positive. As we come down the group, the electron gain enthalpy decrease. However, the first element helium has exceptionally small electron gain enthalpy(48 kJmol-1) due to very small atomic size and high tendency to accept extra electrons.
            Electron gain enthalpies of beryllium, magnesium, calcium and nitrogen  are also positive. This is attributed to the extra stability of the fully  completed s-orbitals in Be(2s2),          Mg (3s2) and Ca(4s2) and of exactly half-filled 2p-orbitals in                        N (2s2Px1Py1Pz1). Thus, if an atom has fully filled or exactly      half-filled 2p-orbitals, its electron gain enthalpy value will be positive.
19.        Which of the following will have the most negative electron gain enthalpy and which the least negative ?  P, S, Cl, F .  Explain your answer. 
7. Electronegativity
The term electronegativity has been defined differently by different investigators.
i)        Pauling's Scale : Pauling defined electronegativity as the power of atom in a molecule to attract electrons to itself.
ii)      Mulliken's –Jaffe Scale : Mulliken suggested that the average of First Ionisation Energy (IE1) and Electron Affinity (EA) of an atom should be a measure of electronegativity of an atom. Accordingly, he proposed the following empirical correlation for calculating electronegativities of various elements.

iii)    Allred –Rochow scale : Allred and Rochow defined electronegativity as the force of attraction exerted by the nucleus of an atom on the valency electrons. Making use of the effective nuclear charge at the periphery of the atom, Zeff , they proposed the following empirical relation for calculating the electronegativity :
where c is the electronegativity and r is the covalent radius of the atom in angstrom units.
            The electronegativity of any given element is not constant ; it varies depending on the element to which it is bound. Though it is not a measurable quatity, it does provide a means for predicting the nature of force that holds a pair of atoms together.
Factors affecting Magnitude of Electronegativity
Magnitude of electronegativity depends up on the following    factors :
(i)          Size of the atom : Smaller the size of the atom, the greater is the attraction for bonding electrons. Thus, atoms with small size are more electronegative.
(ii)         Nuclear charge : Electronegativity value increases with increase in nuclear charge. This is because the nucleus with higher nuclear charge attracts the shared pair of electrons more strongly towards itself.
(iii)        Type of the ion : A cation attracts the electron pair strongly towards itself as compared to the atom from which it has been derived. This is due to the smaller size of the cation as compared to the neutral atom from which it has been derived. The cation has higher electronegativity than the parent atom. For example , the electronegativity of Li is 1.0 while that of Li+ is 2.5. 
An anion attracts the electron pair less strongly as compared to its parent atom. This again is due to the fact an anion has larger size as compared to the parent atom from which it is derived. Thus, an anion has less electronegativity than the parent atom. For example, electronegativity of F atom is 4.0 while that of F- is 0.8.
If an atom can form different types of cations, then the cation with higher positive charge has more value of electronegativity than the cation with lower positive charge. This is due to the fact that the cation with higher value of positive charge has smaller size and hence greater attraction for electrons. For example, electronegativity of Fe2+ is 1.83 and that of Fe3+ is 1.96 . Similarly electronegativity of Sn2+ is 1.81 and that of Sn4+ is 1.96.
(iv)        Ionisation energy and electron affinity : The electronegativity of an element is the mean of ionisation energy and electron affinity. Thus elements having higher value of ionisation energy and electron affinity also have high values of electronegativity. For example, the elements of Group 1 (alkali metals) having lowest electron affinity and ionisation energy , have lowest electronegativity values. Similarly, elements of group 17 (halogens) which have high values of ionisaton energy and electron affinity have high electronegativity values.
Applications of Electronegaivities
Electronegativities have a very wide range of applications. Some of the important applications are given below :
(i)      Non-metallic elements have strong tendency to gain electrons. Therefore, electronegativity is directly related to that non-metallic properties of elements. Electronegativity is inversely related to the metallic properties of elements.  Thus increase in electronegativities across a period is accompanied by an increase in non-metallic properties (or decrease in metallic properties) of elements . Similarly , the decrease in electronegativity down a group is accompanied by a decrease in non-metallic properties (or increase in metallic properties) of elements.
(ii)     Calculation of  partial ionic character of a covalent bond. The development of ionic character in a covalent bond between two atoms , say A and B , is due to the difference in the electronegativities of A and B. The greater the difference in the electronegativities, the greater would be the development of ionic character and consequently the higher would be the stability of the resulting bond.
(iii)    Calculation of enthalpies of  formation of compounds.
(iv)    Calculation of bond length. If the two atoms A and B bonded together through a covalent bond differ in their electronegativities, then the covalent bond would acquire some ionic character, i.e., the bond acquires polarity. The greater the polarity, the shorter would be the length of the bond formed between A and B.
(v)     Explanation of bond angles . The lesser the electronegativity of the central atom in a polyatomic molecule, the lesser would be the bond angle.
(vi)    Explanation of diagonal relationship :The electronegativity increases as we go from Li to Be (variation of electronegativity in a period) but it decreases as we move from Be to Mg (variation of electronegativity in a group). As a result of these two opposite changes(one along the period and other down the group), as we move diagonally , these two effects partly cancel each other and there is no marked change in electronegativity.

      This the reason why Li and Mg have similar chemical    properties.
(vii)   Type of bondsThe type of bond formed between two atoms would evidently, depend upon the differences in their electronegativities.  If this difference is zero or very small, the bond formed would be covalent and  and if the difference exceeds 2.5 , the bond formed would be electrovalent. If however , the electronegativity difference is less is less than 2.5 but otherwise quite appreciable, the bond formed will be polar covalent. The following table gives the relationship between the percentage ionic character and electronegativity difference.
Electronegativity difference
% Ionic character

Electronegativity Scale
The one which is the most widely used is Pauling scale. Pauling (1922) assigned arbitrarily a value of 4.0 to fluorine , the element considered to have the greatest ability to attract electrons. Approximate values of electronegativity of a few elements are given in the Table
Elecronegativities of the Main Group Elements


a) Variation along a Period.
Electronegativity increases along a period from left to right . Across the period, the combined effect of increased nuclear charge and decreasing atomic size results in an increasing tendency to attract the binding electrons in compounds.
b) Variation along a Group
In a given group, the electronegativity decreases with increase in atomic number. This is due to increased atomic size
Periodic Trends
All the Periodic Trends are summarised in the following Fig.

The periodic trends of elements in the periodic table
Anomalous properties of Second Period Elements
            The first element  of each of the group 1 (Lithium) and 2 (Beryllium) and groups 13 – 17 (boron to fluorine) differs in many respects from other members of their respective group. For example, lithium unlike other alkali metals and beryllium unlike other alkaline earth metals form compounds with pronounced covalent character ; the other members of these groups predominantly form ionic compounds. In fact the behaviour of lithium and beryllium is more similar with the second element of the following group i.e., magnesium and aluminium respectively. This sort of similarity is commonly referred to as diagonal relationship in the periodic properties.
Reasons for the anomalous behaviour
            The anomalous behaviour is attributed to their small size, large charge / radius ratio and high electronegativity of the elements. In addition, the first member of the group  has only four valence orbitals ( 2s and 2p) available for bonding, whereas the second member of the groups have nine valence orbitals          (3s, 3p, 3d) . As a consequence of this, the maximum covalency of first member of each group is 4 (e.g. boron can only form [BF4]4- , whereas the other members of the groups can expand their valence shell to accommodate more than four pairs of electrons e.g.  aluminium forms [AlF6]3-. Further , the first member of        p-block elements displays greater ability to form Pp-Pp multiple bonds to itself( e.g. C=C , CºC , N=N , NºN) and to other second period elements (e.g. C=O, C=N, CºN, N=O) compared to subsequent members of the same group.
20.  Are the oxidation state and covalency of Al in [AlCl(H2O)5]2+
same ?
Periodic Trends and Chemical Reactivity
            All chemical and physical properties are manifestation of the electronic configuration of elements. The atomic and ionic radii generally decrease in a period from left to right. As a  consequence , ionisation enthalpies generally increase  and electron gain enthalpies  become more negative  across a period. In other words ionisation enthalpy of extreme left element in a period is the least and electron gain enthalpy of the element on the extreme right  is highest negative( note ; noble gases having completely filled shells  have rather positive electron gain enthalpy values). This results into high chemical reactivity  at the two extremes and lowest in the centre. Thus, the maximum chemical reactivity at the extreme left (among alkali metals) is exhibited by the loss of an electron leading to the formation of a cation and at the extreme right (among halogens) shown by gain of an electron forming an anion. This property is related with the reducing and oxidising behaviour of the elements. The metallic character of an element , which is highest at extremely left decreases and the non-metallic character increases while moving from left to right across the period. The chemical reactivity of an element can be best shown by its reactions with oxygen and halogens. Elements on the two extremes of a period easily combine with oxygen to form oxides. The normal oxide  formed by the element on extreme left is most basic (e.g. Na2O), whereas that formed by the element on extreme right is the most acidic (e.g. Cl2O7). Oxides of the elements in the centre are amphoteric (e.g. Al2O3, As2O3) or neutral (e.g. CO, NO, N2O). Amphoteric oxides behave as acidic with bases and as basic with acids, whereas neutral oxides have no acidic or basic properties.
            Among transition metals (3d series) , the change in atomic radii is much smaller as compared to those of representative elements across the period. The change in atomic radii is still smaller among inner-transition metals (4 f series). The ionisation enthalpies are intermediate between those of s- and    p-blocks. As a consequence, they are less electropositive than group 1 and 2 metals.
            In a group, the increase in atomic and ionic radii with increase in atomic number generally results in a gradual decrease in ionisation enthalpies and a regular decrease (with exception in some third period elements) in electron gain enthalpies in the case of main group elements. Thus , the metallic character increases down the group and non-metal and non-metallic character decreases. The trend can be related with their reducing and oxidising property. In the case of transition elements , a reverse trend is observed. This can be explained in terms of atomic size and ionisation enthalpy.

Properties of Halides, Hydroxides, Sulphates and Carbonates of Alkali and Alkaline earth Metals.
Periodic variations are observed in the properties of many compounds within a group of the periodic table. Some of the trends are given below.
1.    Melting and boiling points of halides of alkali metals.
The halides of alkali metals are all high melting, colourless crystalline solids. The high melting and boiling points of these alkali metal halides is due to their ionic nature. For the alkali metal halides, the melting and boiling point follow the following order:
(i) For the same alkali metal ion, the melting and boiling points decrease in the order :  F-  >  Cl-    >   Br-   >  I-.           The melting points of different halides of sodium metal are plotted in the following Fig.

Variation of melting points of Sodium halides.

(i)      For the same halide ion, melting points of lithium halides are less than those of sodium halides and thereafter they decrease as we go down from sodium to caesium. The melting points of alkali metal chlorides are given in the following Fig.

 Variation of Melting points of
alkali metal chlorides.
The low melting point of lithium chloride is probably due to very small size of the atom as a result of which its compounds are mostly covalent. The melting point of LiCl , is therefore, less than that of NaCl.
2.     Solubility and basic character of hydroxides
The solubilities and basic character of hydroxides of alkali and alkaline earth metals show periodic variation. The hydroxides of alkali metals are strongly basic in nature due to their low ionisation energies. The M-O bond can easily cleave to ionise as M+and OH- ions in solution. The hydroxides are , therefore basic in nature.
          MOH  M+  +   OH-  (basic nature)
The basic nature of hydroxides increases down the group due to decrease in ionisation energy. As the ionisation energy of of alkali metals decreases down the group, M-O bond is more and more readily cleaved and consequently, the basic character increases.
                LiOH      NaOH       KOH      RbOH     CsOH
Basic character increases   ®
            The hydroxides of alkaline earth metals are also basic , but they are less basic than corresponding alkali metals of the same periods. The lesser basic strength of alkaline earth hydroxides is due to higher ionisation energies, smaller ionic size and dipositive charge on the ions. As a result, OH- ions is strongly held by M 2+ ion and M-OH bond is not readily cleaved. As we go down the group, the basic character of the hydroxides increases , mainly due to decreasing  ionisation energies of the metal atoms.  For example, Be(OH)2 is amphoteric, Mg(OH)2 is a weak base, Ca(OH)2and Sr(OH)2 are moderately strong bases and Ba(OH)2 is nearly as strong base as the alkali metal hydroxides.

Be(OH) 2           Mg(OH) 2         Ca(OH)   Sr(OH) 2     Ba(OH) 2
Amphoteric      Weak base                                  Strong base
Basic character increases   ®
            The hydroxides  of alkaline earth metals are slightly soluble in water. With the increase in atomic size of alkaline earth metal as we go down the group, the hydroxides becomes more and more soluble. The solubilities of hydroxides of Mg, Ca, Sr and Ba are shown in the following Fig.

Variation of Solubilities(g/L)  of hydroxides of
alkaline earth metals.
3. Solubilities of Carbonates and Bicarbonates
The carbonates and bicarbonates of alkali metals are generally soluble in water. The solubilities of these salts normally increase as we go down the group from lithium to casesium. The solubilities of carbonates and bicarbonates of alkali metals are given below:
Solubility (wt %) of Carbonates and bicarbonates.
Alkali metal
1.3(at 86K)
70(at 293K)
74(at 293)
The solubilities of carbonates of alkaline earth metals decrease on moving down the group. For example, MgCO3, is slightly soluble in water but BaCO3 is almost insoluble.
4. Stabilities of Carbonates
The stabilities of carbonates also show interesting trends. Except Li2CO3 , the carbonates of other alkali metals are quite stable towards heat and do not decompose. Li2CO3 is not stable and easily decompose on heating.
            The instability of Li2CO3 is due to small size of Li+ ion. Li+ ion being smaller in size can stabilise only small O2- ion and not bigger CO32- ion. Other alkali metal ions are larger in size and can easily stabilise larger CO32- ion and therefore M2CO3 lattices are more stable. The carbonates of alkaline earth metals are relatively less stable towards heat and decompose to give CO2. However, their stability increases as we move down the group. The temperature at which these carbonates decompose are :
     BeCO3           MgCO3             CaCO3                   SrCOBaCO3
  < 373 K       813 K         1173K          1563K         1635 K
5.  Solubilities of Sulphates of alkaline earth metals.
The sulphates of alkaline earth metals are less soluble in water. The solubility decreases down the group. For example, MgSO4  is soluble in water. CasO4 is slightly soluble while SrSO4 and BaSO4 are insoluble.
Diagonal Relationship  
On moving diagonally across the periodic table the elements show certain similarities. These are usually weaker than the similarities within a Group but are quite pronounced in the case of few elements given below.


This  can be more clearly understood in the light of electronegativity across a period and on descending a Group.
            Starting from Lithium and moving a step right brings us to Beryllium which is a little more electronegative than Lithium. Moving one step down  from Beryllium brings us to magnesium which is less electronegative than Beryllium. We might therefore expect lithium and Magnesium to have the same values of electronegativity (both being a little less electronegative than Beryllium). Being equally electronegative Lithium and Magnesium resemble each other in a number of respects( Diagonal relationships).
21.   Show by chemical reaction with water that Na2O is a basic oxide and Cl2O7 is an acidic oxide.
22.   Among the elements of third period  Na to Ar pick out the element :
(i)        with the highest ionisation enthalpy
(ii)       with the largest atomic radius
(iii)      that is the most reactive non-metal
(iv)      that is the most reactive metal.
23.   Which of the following pairs of elements would have a more negative electron gain ethalpy ? Explain.
 (i)  N or O      (ii)   F or Cl
24.  Name a species that will be isoelectronic with each of the following atoms or ions :
(i)   Ne           (ii)   Cl-    (iii)   Ca2+     (iv)    Rb
25.   In each of the following pairs , which species has a larger  size ? Explain.
(i)  K or K+   (ii)  Br or Br-  (iii) O2- or F-   (iv)  Li+ or Na+  (v)  P or As    (vi)  Na+ or Mg2+
26.   Arrange the following ions in the order of size
Be2+, Cl-, S2-, Na+, Mg2+, Br-
27. From each set , choose the atom which has the largest ionization enthalpy and explain your answer
(a)   F, O, N     (b)  Mg, P, Ar     (c) B, Al, Ga
28.  Select from each group the species which has the smallest radius stating appropriate reason.
(a)   O, O-, O2-  (b)   K+, Sr2+, Ar     (c)   Si, P, Cl
29.  Arrange the following elements in the increasing order of metallic character.   B, Al, Mg, K
30.  Predict the position of the element in the periodic table satisfying the electronic configuration (n-1) d1ns2 for         n = 4.
31.    Elements A, B, C, D and E have the following electronic
       configurations :
        A :   1s22s22p1              B.   1s22s22p63s23p1
C :   1s22s22p63s23p3    D.    1s22s22p63s23p5
E :   1s22s22p63s23p64s2
Which among these will belong to the same group in the periodic table ?
32. The first (IE1) and the second (IE2) ionization enthalpies
      (kJ/mol) of three elements I, II, III are given below.
   I             II              III
            IE1         403        549          1142
            IE2         2640       1060         2080
        Identify the element which is likely to be
        (a) non-metal   (b) an alkali metal   (iii) an alkaline earth

33.  The element 119 has not been discovered . What would be the IUPAC name and symbol for the element ? On the basis of periodic table, predict the electronic configuration of this element and also its formula of its most stable chloride and oxide.
34.  Among elements B, Al, C and Si
(a)       which has the highest first ionization enthalpy.
(b)       which has the most negative electron gain enthalpy ?
(c)       which has the largest atomic radius ?
(d)       which has the most metallic character ?
35.  Which of the elements Na, Mg , Si and P would have the greatest difference between the first and second ionization enthalpies. Briefly explain your answer.
36. Which of the following species will have the largest and smallest size ?  Mg, Mg2+, Al, Al3+
37. Using the Periodic Table, predict the formulae of compounds which might be formed by the following pairs of elements :  
(a)   silicon and bromine      (b)   aluminium and sulphur
38. Considering the atomic number and position in the periodic table, arrange the following elements in the increasing order of metallic character : Si, Be, Mg , Na, P.
39. Which of the following will have the most negative electron gain enthalpy and which the least negative ? P, S, Cl, F. Explain your answer.
40. In terms of period and group where would you locate the element with Z = 114 ?
41. Write the atomic number of element present in the third period and seventeenth group of the periodic table.
42.  Which element do you think would have been named by :
(i)   Lawrence Berkley Laboratory
(ii)   Seaborg’s group ?
43. The first (DiH1) and the second(DiH2) ionisation enthalpies
      (in kJ mol-1) of few elements are given below :
Elements    DH1        DH2        DegH
I           520       7300       - 60
II          419       3051       - 48
III        1681       3374       - 328
IV         1008       1846       - 295
V          2372       5251       + 48
VI           738      1451       - 40
          Which of the following elements is likely to be :
       (a)   the least reactive element
       (b)  the most reactive metal
       (c)   the most reactive non-metal
       (d)  the least reactive non-metal
(e)   a metal which can form a stable binary halide of the
       formula MX ( X = halogen) ?
44.   Predict the formula of the stable binary compounds that would be formed by the combination of the following pairs of elements.
   (a)  Lithium and oxygen
   (b)  Magnesium and nitrogen
   (c)  aluminium and iodine
   (d)  silicon and oxygen
   (e)  Phosphorus and fluorine
   (f)  Element 71 and fluorine
45.    Assign the position of the element having outer electronic configuration :
       (i)   ns2np4  for n = 3      (ii)   (n -1) d2ns2 for n = 4 and (iii) (n - 2 ) f7(n -1) d1ns2 for n = 6 , in the periodic table.
46.     The increasing order of reactivity among group 1 elements is Li < Na < K < Rb < Cs whereas that among group 17 elements is F > Cl > Br > I.
47.    Use the periodic table to answer the following questions.
       (a)   Identify an element with five electrons in the outer
       (b)  Identify an element that would tend to lose two
       (c)   Identify an element that would tend to gain two
       (d)  Identify the group having metal, non-metal, liquid as
                well as gas at room temperature.
48.   The atomic mass of germanium is 72.6 and its density is 5.47 g cm-3. What is the atomic volume of  germanium ?
49.    The ionisation enthalpy of lithium is 520 kJ mol-1. Calculate the amount of energy required to convert 70 mg of lithium atoms in gaseous state into Li+ ions.
50.    The ionisation potential of hydrogen is 13.6 eV. Calculate the energy required to produce H+ ions from 0.5 g of hydrogen atoms.
1.        Write a note on Mendeleev's Periodic Table.
2.        Give the uses of Mendeleev's Periodic table.
3.        What are drawbacks of Mendeleev's periodic Table.
4.        Give an account of Modern Periodic Table.
5.        Give the features of Log form of the Periodic Table.
6.        Compare Modern periodic Table with that of Mendeleev's Periodic Table.
7.        What are s-Block elements ? Discuss their characteristics.
8.        What are p-block elements. Discuss their characteristics.
9.        What are d-block elements ? Discuss their characteristics.
10.     What are f-block elements ? Discuss their characteristics.
11.      What property did Mendeleev used to classify the elements in  the periodic Table ?
12.     State Modern Periodic Law.
13.     Explain the term 'periodicity' in the properties of elements.
14.     Explain the term regular gradation of properties of elements.
15.     Explain the terms :
(i)   Covalent radius         
(ii)    van der waal's radius 
(iii)   Ionic radius.
16.   How do atomic sizes vary in a group and in a period ? Give reasons for the variation.
17.   Explain the term atomic volume. How does it vary :
i)         in a period
 ii)   in a group
18.   Which of the following pairs would have a larger size ? Explain.
i)        K   or K+             
ii)     Br or Br-  
iii)    O2- or F-
iii)    Li+ or Na+ 
iv)   Na++ or Mg2+               v)    P or As     
19.   Explain the following :
i)        The size of calcium is smaller than its parent atom.
ii)      The size of an anion is larger than its parent atom.
iii)     Mg2+  ion   is smaller than O2- ion, although both have the same electronic structure.
20.   Choose the ion with the largest ionic size.
        Ar   , K+, S2- O2-, F-
21.   Which of the following species are iso-electronic with noble gas argon ?
         Ar   , K+, S2- O2-, Ca2+
22.   Explain the terms :
i)        ionisation energy   
 ii)   Electron affinity.
24.   Discuss the factors that influence the magnitude of ionisation energy.
25.   Discuss the factors that influence the magnitude of electron affinity.
26.   What are the general trends of variation of electron affinity in periodic table ?
27.   The electron affinity of chlorine is more that of fluorine. Give reason.
28.   Discuss the variation of ionisation energy amongst the elements of a period of the periodic table.
29.   Why does first ionisation energy decreases as we go from top to bottom in the group of the periodic     table ?
30.   Among the elements : Li, K, Ca, S and Kr, which one has the lowest first ionisation energy ? Which has the highest first ionisation energy ?
31.   Which one of the following pairs of elements would you expect to have lower first ionisation energy ? Explain.
i)         Cl or F    
ii)     Cl or S  
iii)     K or Ar  
iv)     Kr or Xe
32.   Explain the first ionisation energy of :
i)        Aluminium is lower than that of magnesium.
ii)      Nitrogen is more that of elements on its sides i.e., (Carbon and Oxygen) in the same period ?
33.   Second and third ionisation energies of an element are always greater than its first ionisation energy. Explain why ?
34.   Which of the following pairs of elements would have a higher electron affinity ?
 i) N or O 
ii) F  or Cl. Explain.
35.   Explain why electron affinity of atoms increases from left to right along a period in the periodic table ?
36.   How does electron affinity vary in a group in the periodic table ?
37.   Electron affinity values of inert gases are zero. Comment.
38.   Why electron affinity of halogens are high ?
39.   What is meant by the term 'electronegativity' of an element ?
40.   How does electronegativity vary in a :
(i)   Group         
(ii)  Period ?
41.   Electronegativity values of inert gases are zero. Explain.
42.   How does chemical reactivity vary in a :
(a)     Group     (b) Period.
43.   Among the elements of the second period, pick out the  element :
i)         with the highest first ionisation energy.   
ii)        with the highest atomic radius.
iii)       with largest atomic radius.
iv)       that is most reactive non-metal.
v)        that is the most reactive metal.
44.   Predict the density of Cs from the density of the following elements:
       K :  0.86 g/cm3   Ca  :  1.548 g/cm3 Sc : 2.991 g/cm3
       Rb: 1.532 g/cm3 Sr  : 2.68 g/cm3     Y  : 4.34 g/cm3
       Cs   :  ?               Ba  : 3.51 g/cm3      La : 6.16 g/cm3
45.   Account for the fact that the 4th period has eighteen and not eight elements.
46.   Give the formula of a species that will be isoelectronic with the following atoms or ions:
i)        Ne          ii) Cl-        iii) Ca2+      iv) Rb+
47.   The valence of representative element is either equal to the number of the valence electrons or eight minus this number. What is the basis of this rule
48.   The first (IE1) and the second (IE2) ionisation energies (kJ/mol) of a few elements designated by Roman numerals are shown below:

      Which of the above element is likely to be :
(i)      a reactive metal
(ii)     a reactive non-metal.
(iii)   a noble gas.
(iv)    a metal that forms a stable binary halide of the formula AX2 (X = halogen) ?
49.   Lanthanides and actinides are placed in separate rows at the bottom of the periodic tabble. Explain the reasons for this arrangement.
50.   The elements Z = 107 and Z = 109   has not yet been made. Indicate the groups which you will place the above elements.
51.   What is meant by electropositive character of an element ? How does it vary in a :
(a)     Period            b)  Group ?
52.   The ionisation energy of beryllium is much larger than that of boron, though boron is slightly larger in size than beryllium.   Why ?
53.   Explain the following :
       The first ionisation energy of carbon atom is greater
        than that of boron atom ; whereas the reverse is
        true for the second ionisation energy.
54.   Explain with arguments why the first ionisation energy of oxygen is lower than that of nitrogen
55.   Give an example of diagonal relationship in periodic table. Why is it observed ?


Atoms and Molecules

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