Understanding Periodic Trends in Atomic Radii and Ionization Energy

1. Periodic properties Periodic properties Periodicity If the elements are arranged in order of increasing atomic numbers, then their properties get repeated after regular intervals. The repetition of similar properties of the elements after regular intervals is called periodicity. Atomic radii The atomic radius is defined as the distance between the centre of nucleus and the outermost shell of electron. Nucleus + Shell

2. The atomic radii are sub-divided into three types. Atomic radii Covalent radii van der waals radii metallic radii 1) Covalent radii It is defined as, one half of the distance between the nuclei of two covalently bonded atoms in a molecule. Covalent radius

3. a) Homo-nuclear diatomic molecule: In a homonuclear diatomic molecule the covalent radius of an atom is obtained by dividing inter-nuclear distance by two. For eg. The internuclear distance between the two atoms in F2molecule is 1.28 (Fig. 1) 0.64 Å F F Covalent radius of F = 1.28 /2 = 0.64 Å Internuclear distance= 1.28 Å

4. Periodic trends in covalent radii A) Along period As we move from left to right along period, atomic radii decreases because 1) The nuclear charge increases due to that valence shell electron are attracted more towards the nucleus. 2) The electron are added in the last same shell and hence electron cloud is concentrated closer to the nucleus and therefore atomic radius decreases. B) Along group As we move from top to bottom in the group. atomic radii increases because 1) The nuclear charge increases but new orbital's are created to fill the electron and nucleus is shielded more. 2) Due to shielding effect the effective nuclear charge decreases which causes increases in atomic radii.

5. 2) Van der waals radii Van der waals radius is defined as one half of the distance between the nuclei of two non bonded neighboring atoms of two adjacent molecules in a solid state. 1.80 Å Cl Cl Van der waals radius 3.60 Van der waals radius of Cl =3.60 2 =1.80 Å

6. Metallic radii The metallic radius is defined as one half of the distance between the nuclei of two adjacent atoms of metals in a metallic closed packed crystal. For eg. The internuclear distance between two adjacent sodium atom in crystal lattice of sodium metal is 3.72 Å 3 .72 2 Metallic radius of Na-atom= 3 .72 Metallic radius

7. Ionic radii Ionic radius is defined as the distance between the centre of nucleus of an ion and point up to which the nucleus has influence on its electron charge cloud. Internuclear distance = Radius of cation + Radius of anion rc+- d(c+-A-) = + rA- 2.76 Na+ Cl- For eg. Ionic radius of Na+ion is 0.95 and internuclear distance between Na and Cl ion pair in Na+ Cl- ionic crystal is 2.76 find out ionic radius of Cl ion. d(Na+- Cl-) = rNa - rCl = 0.95 + rCl 2.76 = 2.76 - 0.95 rCl = 1.81

8. A) Along period The ionic radii decreases from left to right in any period in periodic table due to increase in ENC per electron and hence valence shell electron are bound more tightly to the nucleus. B) Along group The ionic radii increases from top to bottom in any group in periodic table because the electron are added in extra valence shell and due to it d and f orbital's do not shield the nucleus very effectively and hence electron are loosely held to the nucleus and ionic radii increases.

9. Ionization energy The ionization energy is defined as the amount energy required to remove the most loosely bounded electron from valence shell of an isolated gaseous atom to convert it into gaseous cation. Mg (g) Mg+ (g) + e- 744 kJmol-1 The energy required to remove the first electron from an isolated gaseous atom is called first ionization energy. Al (g) Al+ (g) + (2,8,3) (2,8,2) e- = + I.E.1 The energy required to remove the second electron from an isolated gaseous atom is called second ionization energy. Al (g) (2,8,2) Al (g) (2,8,1) Al2+ (g) (2,8,1) Al3+ (g) (2,8) + e- = + I.E2 + e- = + I.E3 Thus, the magnitude of successive ionization energy increases in the following order I.E.1<I.E.2< I.E.3

10. Factors affecting ionization energy IE of an element depends upon the following factors. 1) Atomic size The IE increases as the atomic size decreases and vice versa. In larger atom valence shell electron are present at larger distance from the nucleus and experiences less nuclear attraction. Therefore it is easier to nock out an electron from the outer shell of an atom. 2) Nuclear charge Increase in nuclear charge on atom or ion increases the ionization potential. As nuclear charge Increase the attraction of electron from outermost shell to the nucleus increases so that it requires more amount of energy for its removal.

11. 3) Screening effect When large no. of electron are in the inner shell then low will be the ionization energy. The combined effect of this attractive and repulsive force acting on the valence electron is that the valence electron experiences less attraction from the nucleus this is known as screening effect. Thus larger the no. of electron in the inner shell lesser is the attractive force holding. 4) Electronic configuration The atom with half filled or full filled orbital's in valence shell acquires extra stability and requires somewhat more IE than that of normally expected from its position in the periodic table.

12. Trends of ionization energy A) Along period The IE increases from left to right along period in periodic table due to  atomic size decreases  nuclear charge increases  electron are added in the same principal energy level. The attractive force between outermost electron and nucleus increases therefore IE increase. B) Along group In general IE decreases as we move from top to bottom along a group in the periodic table. This is because of  Nuclear charge increase  The gradual increase in the atomic size due to increase in the no. of principal energy shell.  There is increase in the no. of inner electron which shields the valence electron from the nucleus.

13. Electron affinity The electron affinity is defined as the amount of energy released when an extra electron is added to the valence shell of an isolated gaseous atom to convert it into gaseous negative ion. Cl-(g) (2,8,7) (2,8,8) Cl(g) + e- + Energy The addition of first electron involves release of energy according to thermodynamic conversions it is exothermic reaction When first electron is added to the valence shell of an atom to form gaseous anion then it is called first electron affinity energy O(g) O-(g) (2,6) (2,7) O-(g) O- -(g) + (2,7) (2,8) The EA1 of oxygen is 141.13 KJ/mole and EA2 is 844.0 KJ/mole. If 2 electron is added to a mononegative gaseous anion the energy is required not liberated and consequently the process defining EA2 is endothermic process. + e- + Energy + e- Energy

14. Factors affecting electron affinity 1) Atomic size In case of smaller atoms the attraction of the nucleus for the electron to be added is stronger. Thus smaller is the size of atom grater is its electron affinity. The electron affinity of halogen is grater than the rest of the elements in respective group, period. Grater is the nuclear charge of elements (along period) stronger is the attraction of its nucleus for the electron to be added. Thus increase in the magnitude of nuclear charge increases the electron affinity. 3) Screening effect The electron affinity decreases with increase in screening effect or shielding because due to screening effect the added electrons experiences less nuclear attraction. 4) Electronic configuration The elements which have stable electronic (half or full filled) configuration For eg. Be (1s2,2s2), N (1s2,2s2, 2p3) and Mg (1s2,2s2, 2p6,3s2) possesses negligible electron affinity and nobel gases passes almost zero electron

15. Trends of electron affinity A) Along period On moving from left to right along period electron affinity increases. It is due to  Atomic size decreases  Nuclear charge increases  Number of valence electron increases so that addition of electron becomes easier from left to right along period. B) Along group Electron affinity decrease from top to the bottom in a group in a periodic table because of  Atomic size increases due to addition of new main energy shells.  Effective nuclear charge decreases  Electron are less attracted towards the nucleus, hence electron affinity decreases from top to bottom in a group.

16. Electronegativity It is defined as the tendency of an atom to attract shared pair of electrons towards itself in a covalently bonded molecule. The formation of H-Cl molecule takes place as follows. In this the electronegativity of chlorine is more therefore shared pair electron is attracted towards chlorine atom due to this molecule becomes polar i.e. Chlorine have –ve and hydrogen have +ve charge as shown below.

17. Factors affecting Electronegativity 1) Atomic size The atoms with small atomic size has grater tendency to attract the shared electron pair towards itself than the larger atom. When effective nuclear charge increases then the electronegativity increases. Electropositive character of element is opposite to that of electronegative character. In short the elements which are highly electropositive are weakly electronegative.

18. Trends of Electronegativity A) Along period As we move from left to right along period electronegativity increases. It is due to  Atomic size decreases  Nuclear charge increases  As we move from left to right along period electropositive character decreases. It means electronegative character increases in the same direction. B) Along group Electronegativity decrease from top to the bottom in a group in a periodic table because of  Atomic size increases due to addition of new main energy shells.  Effective nuclear charge decreases  Down a group electropositive character increases it results in decrease in electronegativity.

19. Determination of Electronegativity Different chemist in various ways defined the term electronegativity and hence they have their own definitions and method to calculating it. Some important methods scales are given below 1) Pauling electronegativity scale(1932) According to Linus Pauling electronegativity of an atom in a molecule is its ability to attract shared electron pairs towards itself. Consider the formation of A-B molecule by combination of ½ A2and ½ B2 molecule ½ A2 ½ B2 AB 1 Since A2,B2,and AB are covalent molecule so that can be represented as A-A, B-B, and A-B respectively. Then the equation 1 can be written as ½ B-B ½ A-A A-B 2

20. Then pauling putting a equation for determination of electronegativity 0.182[EA-B(EA-Ax EB-B)1/2]1/2 XA-XB= This scale is not accepted universal because the bond dissociation energy cannot be determined easily and correctly

21. 2) Mullikens electronegativity scale It is based on the ionization energy and electron affinity of an atom. He proposed that the electronegativity of an atom can be calculated with the help of following equation (IE)A + (EA)A 2 XA= 0.374 + 0.17 ………………[1] XA= electronegativity of atom A IEA= ionization energy of atom A EAA= electron affinity of atom A The above equation holds good when EA and IE both are measured in electron volt (ev). If it is in KJ/Mole then equation 1 becomes (IE)A + (EA)A 540 ………………[2] XA=

22. Applications of electronegativity 1) Nature of bond If the electronegativity difference between two atom is grater then bond A-B is predominantly ionic and if it is equal or slightly differ from each other then bond A-B is covalent 2) Percentage of ionic character Linus pauling has given the following relation between (XA-XB) and percentage ionic characters. XA-XB 0.6 1.0 1.4 1.7 2.0 2.4 3.0 3.2 % ionic character 9 22 39 51 63 76 91 92 The above data shows that a. If (XA-XB) = 1.7 A-B bond is 50% ionic and 50% covalent in nature b. If (XA-XB) < 1.7 ionic charcter is less than 50% covalent in nature and covalent character is more than 50% c. If (XA-XB) > ionic charcter is more than 50% and covalent charcter is less than 50%

23. 3) Stability of bond in molecule If the electronegativity difference between two atom is more then the bond is more stable molecule HF HCl HBr HI Diff. in electronegativity 1.9 0.9 0.7 0.4 Stability HF > HCl > HBr > HI Heat of formation 297 92 52 6 4) Metallic and non-metallic charcter On the basis of electronegativity of an element the metallic and non-metallic charcter can be determined. The electronegativity are non-metal and electronegativity are metallic in nature. element those have having higher lower value value of of

24. Radius of cation Radius of cation is always smaller than its parent atom because 1) Due to formation of +ve ion, the outer most shell electron is generally removed completely. 2) Due to formation of cation no. of electron are decreases while the nuclear charge remain same. electron Na+ Na-atom For eg. Sodium atom is converted into Na+cation by loosing one electron In Na-atom there are 11 protons and 11 electron but in Na+ cation (10 electron). The ENC per electron is more than the parent atom so it attracts the electron from the valence shell more strongly towards the nucleus. Hence the radius of cation is small than its parent.

25. Radius of anion Radius of anion is always grater than its parent atom because If one or more electron are added in the valence shell of an atom then it possess –ve charge. Due to addition of electron the ENC per electron Decreases & hence electron are held less tightly by nucleus and electron cloud extends and ionic size decreases. electron + + Cl=2,8,7 Cl-=2,8,7 rCl=1.81 rCl=0.99 For eg. Cl (17) having electronic configuration 1s2,2s2,2p6,3s2,3p5 The parent cl atom contains 17-electron but Cl-ion has 17 protons and 18 electron. The ENC per electron decreases and electron are loosely bound toward the nucleus and due to this ionic radius is more than the parent atom.