Question 1

Which noble gas has the highest ionization energy?

Accepted Answer

Helium (He)

Question 2

Given the following successive ionization energies (in kJ/mol) for a period 3 element: IE2 = 2250, IE3 = 3360, IE4 = 4560, IE5 = 7010, IE6 = 8500, IE7 = 27,100. Which element is this, and what does the large jump in IE7 indicate?

Accepted Answer

Phosphorus; the jump indicates removal of a core electron.

Question 3

Why do Li, Na, and K have similar chemical properties?

Accepted Answer

They have the same number of valence electrons.

Question 4

What is the ground state electron configuration for $$\mathrm{Cd$$^{2+}$$}$$?

Accepted Answer

$$[\mathrm{Kr}]\ 4d$$^{10}$$

Question 5

Why is the first ionization energy of sulfur smaller than that of phosphorus?

Accepted Answer

Sulfur has paired electrons in the 3p orbital, increasing electron-electron repulsion.

Question 6

Why do successive ionization energies increase?

Accepted Answer

The remaining electrons are closer to the nucleus and experience greater attraction.

Question 7

Using sodium as an example, differentiate between electron affinity and ionization energy.

Accepted Answer

Electron affinity is the energy released when an electron is added; ionization energy is the energy required to remove an electron.

Question 8

Which atom in each group has the smallest atomic radius? (I) K, Zn, Br (II) Al, Ga, In

Accepted Answer

Zn and Ga

Question 9

Explain why the lattice energy of $$\mathrm{MgS} is $$approximately 4 times as large as that of $$\mathrm{NaCl}.$$

Accepted Answer

MgS has higher charges on its ions, increasing the electrostatic attraction.

Question 10

What is the ground-state electron configuration of $$\mathrm{Se$$^{2-}$$}$$?

Accepted Answer

$$[\mathrm{Ar}]\ 4s$$^{2}$$\ 3d$$^{10}$$\ 4p$$^{6}$$

Question 11

Choose the compound below that should have the highest melting point according to the ionic bonding model. Explain.

Accepted Answer

MgO, because it has ions with the highest charges and smallest radii.

Question 12

Estimate $$\Delta H^\circ_{rxn}$$ for the reaction $$2\ \mathrm{Br_2$(l)} + \mathrm{C_2H_2(g)} \rightarrow \mathrm{C_2H_2Br_4(l)}$$ using the following bond energies: $$\mathrm{Br-Br} = 193\ \mathrm{kJ/mol}$$, $$\mathrm{C\equiv C} = 837\ \mathrm{kJ/mol}$$, $$\mathrm{C-C} = 347\ \mathrm{kJ/mol}$$, $$\mathrm{C-Br} = 276\ \mathrm{kJ/mol}$$, $$\mathrm{C-H} = 414\ \mathrm{kJ/mol}. $$What is the estimated $$\Delta H^\circ_{rxn}$$?

Accepted Answer

$$-484\ \mathrm{kJ/mol}$$

Question 13

Estimate $$\Delta H^\circ_{rxn}$$ for the reaction $$\mathrm{$$XeF_2$$} + 2\ \mathrm{F_2$} \rightarrow \mathrm{$$XeF_6$$}$$ using the following bond energies: $$\mathrm{Xe-F} = 147\ \mathrm{kJ/mol}$$, $$\mathrm{F-F} = 159\ \mathrm{kJ/mol}. $$What is the estimated $$\Delta H^\circ_{rxn}$$?

Accepted Answer

$$+318\ \mathrm{kJ/mol}$$

Question 14

Using the provided data, calculate the energy change for the formation of $$\mathrm{LiCl(s)}$$ from its elements in their standard states: $$\mathrm{Li(s) + \frac{1}{2}Cl_2$(g) \rightarrow LiCl(s)}. $$Use the following values: $$\mathrm{Li^+(g) + Cl^-(g) \rightarrow LiCl(s)} = -853\ \mathrm{kJ/mol}$$, $$\mathrm{Li(g) \rightarrow Li^+(g) + e^-} = +520.2\ \mathrm{kJ/mol}$$, $$\mathrm{Li(s) \rightarrow Li(g)} = +159.4\ \mathrm{kJ/mol}$$, $$\mathrm{\frac{1}{2}Cl_2$(g) \rightarrow Cl(g)} = +121.7\ \mathrm{kJ/mol}$$, $$\mathrm{Cl(g) + e^- \rightarrow Cl^-(g)} = -348.6\ \mathrm{kJ/mol}. $$What is the total energy change?

Accepted Answer

$$-409.9\ \mathrm{kJ/mol}$$

Question 15

Which of the following best describes the formal charge calculation for an atom in a Lewis structure?

Accepted Answer

Formal charge = valence electrons minus non-bonding electrons minus half the bonding electrons.

Question 16

According to Coulomb’s law, which factor most strongly affects the lattice energy of an ionic compound?

Accepted Answer

The distance between the ions and the magnitude of their charges.

Calculating the Energy Change for the Formation of LiCl(s) from Its Elements

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