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Chapter 8: Problem 31

Without using Fig. 8.3, predict the order of increasing electronegativity in each of the following groups of elements. a. \(\mathrm{C}, \mathrm{N}, \mathrm{O} \quad\) c. $\mathrm{Si}, \mathrm{Ge}, \mathrm{Sn}$ b. \(\mathrm{S}, \mathrm{Se}, \mathrm{Cl} \quad\) d. $\mathrm{TI}, \mathrm{S}, \mathrm{Ge}$

Short Answer

Expert verified
The orders of increasing electronegativity for the given groups of elements are: a. \( C < N < O \) b. \( Se < S < Cl \) c. \( Sn < Ge < Si \) d. \( Tl < Ge < S \)

Step by step solution

01

Determine Groups and Periods

Check the periodic table to find the group and period of each element. a. C, N, O -> second period, Group 14, 15, and 16 b. S, Se, Cl -> third period, Group 16, 16, and 17 c. Si, Ge, Sn -> Group 14, with Si in period 3, Ge in period 4, and Sn in period 5 d. Tl, S, Ge -> Tl is in Group 13, period 6; S is in Group 16, period 3; and Ge is in Group 14, period 4
02

Apply Electronegativity Trends

Use the periodic table trends to determine the order of increasing electronegativity for each group of elements. a. C, N, O -> Electronegativity increases going from left to right across a period (Group 14 to 16), so the order is C < N < O. b. S, Se, Cl -> Electronegativity decreases going down a group (S and Se are in the same group), and it increases from left to right (Group 16 to 17). So, the order is Se < S < Cl. c. Si, Ge, Sn -> Electronegativity decreases going down a group, so the order is Sn < Ge < Si. d. Tl, S, Ge -> Electronegativity decreases going down a group (Tl and Ge in different periods) and increases from left to right (Group 13 to 14). So, the order is Tl < Ge < S.
03

Final Answer

The orders of increasing electronegativity for the given groups of elements are: a. C < N < O b. Se < S < Cl c. Sn < Ge < Si d. Tl < Ge < S

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Most popular questions from this chapter

Use formal charge arguments to explain why CO has a much smaller dipole moment than would be expected on the basis of electronegativity

Consider the following reaction: $$A_{2}+B_{2} \longrightarrow 2 A B \quad \Delta H=-285 \mathrm{kJ}$$ The bond energy for \(\mathrm{A}_{2}\) is one-half the amount of the AB bond energy. The bond energy of \(\mathrm{B}_{2}=432 \mathrm{kJ} / \mathrm{mol}\) . What is the bond energy of \(\mathrm{A}_{2}\) ?

Classify the bonding in each of the following molecules as ionic, polar covalent, or nonpolar covalent. a. \(\mathrm{H}_{2} \quad\) e. \(\mathrm{HF}\) b. \(\mathrm{K}_{3} \mathrm{P} \quad\) f. \(\mathrm{CCl}_{4}\) c. \(\mathrm{Nal} \quad\) g. \(\mathrm{CF}_{4}\) d. \(\mathrm{SO}_{2} \quad\) h. \(\mathrm{K}_{2} \mathrm{S}\)

Without using Fig. 8.3, predict which bond in each of the following groups will be the most polar. a. \(\mathrm{C}-\mathrm{H}, \mathrm{Si}-\mathrm{H}, \mathrm{Sn}-\mathrm{H}\) b. $\mathrm{Al}-\mathrm{Br}, \mathrm{Ga}-\mathrm{Br}, \mathrm{In}-\mathrm{Br}, \mathrm{Tl}-\mathrm{Br}$ c.\(\mathrm{C}-\mathrm{O}\) or \(\mathrm{Si}-\mathrm{O}\) d. \(\mathrm{O}-\mathrm{F}\) or \(\mathrm{O}-\mathrm{Cl}\)

The standard enthalpies of formation for $\mathrm{S}(g), \mathrm{F}(g), \mathrm{SF}_{4}(g),\( and \)\mathrm{SF}_{6}(g)\( are \)+278.8,+79.0,-775,$ and \(-1209 \mathrm{kJ} / \mathrm{mol}\) respectively. a. Use these data to estimate the energy of an \(\mathrm{S}-\) F bond. b. Compare your calculated value to the value given in Table \(8.5 .\) What conclusions can you draw? c. Why are the \(\Delta H_{f}^{\circ}\) values for \(\mathrm{S}(g)\) and \(\mathrm{F}(g)\) not equal to zero, since sulfur and fluorine are elements?
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