Diagonal Relationship Atomic Radii pm Pauling Electronegativity Since

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Diagonal Relationship Atomic Radii (pm) Pauling Electronegativity ▪ Since the electronegativity of atoms generally

Diagonal Relationship Atomic Radii (pm) Pauling Electronegativity ▪ Since the electronegativity of atoms generally increases across a period and decreases down a group, two atoms which are diagonally related could have similar electronegativities. ▪ The atomic radii of atoms generally decreases across a period and increases down a group, and therefore diagonally related atoms may have similar atomic radii. ▪ Therefore, there are situations where comparable electronegativities and atomic radii may lead to similarities in chemical properties, even though the two elements concerned have different valencies. ▪ Example: Whereas Group 1 elements form compounds that are essentially ionic in nature, Li and Mg salts have some degree of covalent character in their bonding. 1

Enthalpies of Atomization ▪ Definition: The enthalpy of atomization of an element (ΔHa 0

Enthalpies of Atomization ▪ Definition: The enthalpy of atomization of an element (ΔHa 0 ) is a measure of the energy required to form gaseous atoms. For solids, the enthalpy of atomization is the enthalpy change associated with the atomization of the solid. For molecular species, it is the enthalpy of dissociation of the molecules. ▪ Trend: It increases with increasing number of valence electrons. 2

Enthalpies of Atomization ▪ Explanation: It first increase and then decrease across Periods 2

Enthalpies of Atomization ▪ Explanation: It first increase and then decrease across Periods 2 and 3, reaching a maximum at C in Period 2 and Si in Period 3. The values decrease between C and N, and Si and P: even though N and P each have 5 valence electrons, 2 of these electrons form a lone pair and only 3 are involved in bonding. A similar effect is seen between N and O, where O has 6 valence electrons of which 4 form lone pairs and only 2 are involved in bonding. 3

Metallic Character ▪ Definition: It can be defined as the ability an atom to

Metallic Character ▪ Definition: It can be defined as the ability an atom to lose electrons to form the electron sea that binds together the cations and accounts for metallic bonding. Metallic Bonding ▪ Trend: The metallic character of the elements decreases across a period and increases down a group. 4

Metallic Character ▪ Explanation: Elements with low ionization energies are likely to be metals

Metallic Character ▪ Explanation: Elements with low ionization energies are likely to be metals and those with high ionization energies are likely to be nonmetals. Thus, as ionization energies decrease down a group the elements become more metallic, and as the ionization energies increase across a row the elements become less metallic. ▪ This trend is most noticeable within Groups 13 to 16, where the elements at the head of the group are nonmetals and those at the foot of the group are metals. Allotropic Variations 5

Inert Pair Effect ▪ Trend: Heavier elements of p-block (groups 13– 17) often have

Inert Pair Effect ▪ Trend: Heavier elements of p-block (groups 13– 17) often have oxidation states that are lower by 2 than the maximum predicted for their group. ▪ Examples: In Group 13 whereas the group oxidation number is +3, the +1 oxidation state increases in stability down the group. The most common oxidation state of thallium is Tl(I). Similarly, the stability of +2 oxidation state increases down Group 14. Element Electron Configuration I 1 (k. J/mol) I 1 + I 2 + I 3 (k. J/mol) Average M–Cl Bond Energy (k. J/mol) B [He] 2 s 22 p 1 801 6828 536 Al [Ne] 3 s 23 p 1 578 5139 494 Ga [Ar] 3 d 104 s 24 p 1 579 5521 481 In [Kr] 4 d 105 s 2 p 1 558 5083 439 Tl [Xe] 4 f 145 d 106 s 2 p 1 589 5439 373 ▪ The sum of the first three ionization energies of Tl is similar with the values for Ga and higher than the values for In and Al. 6

Inert Pair Effect ▪ Explanations: 1) Higher-than-expected IE: The valence ns electrons are not

Inert Pair Effect ▪ Explanations: 1) Higher-than-expected IE: The valence ns electrons are not shielded from the nucleus very effectively by the intervening (n-1)d and (n-2)f electrons having poor shielding effects. As the two ns electrons are both held tighter by the nucleus, the ionization energies for these two electrons are unusually large. Therefore, it is more difficult than expected to remove the ns 2 electrons after the removal of np 1 electrons. Because Tl is less likely than Al to lose its two ns 2 electrons, its most common oxidation state is +1 rather than +3. 2) Lower-than-expected BDE: Weaker bond dissociation energies (BDE) are expected for the heavier elements as a result of the diffuse nature of orbital overlap with increasing orbital size. Consequently, bond strengths tend to decrease down a column. ▪ The net effect of these two factors—increasing ionization energies and decreasing bond strengths—is that in going down a group in the p-block, the additional energy released by forming two additional bonds eventually is not great enough to compensate for the additional energy required to remove the two ns 2 electrons. 7