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Chapter 8: Problem 15

True or false? In general, a large atom has a smaller electronegativity. Explain.

Short Answer

Expert verified
True. In general, a large atom has a smaller electronegativity due to the greater distance between the nucleus and valence electrons, which reduces its ability to attract electrons towards itself in a chemical bond.

Step by step solution

01

Definition of Electronegativity

Electronegativity is the tendency of an atom to attract a bonding pair of electrons. It is a measure of how strongly an atom can attract electrons towards itself in a chemical bond. Electronegativity values can be found on the periodic table, with higher values indicating a higher electronegativity.
02

Periodic Trend of Electronegativity

Electronegativity generally increases from left to right across a period and decreases as we move down a group on the periodic table. This is because as we go across a period, the number of protons in the nucleus increases, resulting in a stronger positive charge that attracts electrons more effectively. When we move down a group, atoms have more energy levels which leads to a greater distance between the nucleus and valence electrons, reducing the attractive force.
03

Atom Size and Electronegativity

A larger atom has more energy levels and a greater distance between the nucleus and its valence electrons. This makes it more difficult for the atom to attract electrons, resulting in a lower electronegativity value. Conversely, smaller atoms have fewer energy levels and a shorter distance between the nucleus and valence electrons, making it easier for the atom to attract electrons and resulting in a higher electronegativity value.
04

Conclusion

The statement "In general, a large atom has a smaller electronegativity" is true. As the size of an atom increases, its electronegativity tends to decrease due to the greater distance between the nucleus and valence electrons, leading to a decreased ability to attract electrons towards the atom in a chemical bond.

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Most popular questions from this chapter

Use bond energies (Table \(8.5 ),\) values of electron affinities (Table 7.7\()\) , and the ionization energy of hydrogen \((1312 \mathrm{kJ} / \mathrm{mol})\) to estimate \(\Delta H\) for each of the following reactions. a. \(\mathrm{HF}(g) \rightarrow \mathrm{H}^{+}(g)+\mathrm{F}^{-}(g)\) b. \(\mathrm{HCl}(g) \rightarrow \mathrm{H}^{+}(g)+\mathrm{Cl}^{-}(g)\) c. \(\mathrm{HI}(g) \rightarrow \mathrm{H}^{+}(g)+\mathrm{I}^{-}(g)\) d. $\mathrm{H}_{2} \mathrm{O}(g) \rightarrow \mathrm{H}^{+}(g)+\mathrm{OH}^{-}(g)$

Predict the molecular structure (including bond angles) for each of the following. a. \(\mathrm{PCl}_{3}\) b. \(\mathrm{SCl}_{2}\) c. \(\mathrm{SiF}_{4}\)

Think of forming an ionic compound as three steps (this is a simplification, as with all models): (1) removing an electron from the metal; (2) adding an electron to the nonmetal; and (3) allowing the metal cation and nonmetal anion to come together. a. What is the sign of the energy change for each of these three processes? b. In general, what is the sign of the sum of the first two processes? Use examples to support your answer. c. What must be the sign of the sum of the three processes? d. Given your answer to part c, why do ionic bonds occur? e. Given your above explanations, why is NaCl stable but not $\mathrm{Na}_{2} \mathrm{Cl} ? \mathrm{NaCl}_{2} ?$ What about MgO compared to \(\mathrm{MgO}_{2} ? \mathrm{Mg}_{2} \mathrm{O} ?\)

Use the formal charge arguments to rationalize why \(\mathrm{BF}_{3}\) would not follow the octet rule.

Write electron configurations for a. the cations \(\mathrm{Sr}^{2+}, \mathrm{Cs}^{+}, \mathrm{In}^{+},\) and \(\mathrm{Pb}^{2+} .\) b. the anions \(\mathrm{P}^{3-}, \mathrm{S}^{2-},\) and \(\mathrm{Br}^{-}\)
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