Chapter 3

Explain the electronegativity trends across a row and down a column of the periodic table. Compare these trends with those of ionization energies and atomic radii. How are they related?

The ionic compound \(A B\) is formed. The charges on the ions may be +1,\(-1 ;+2,-2 ;+3,-3 ;\) or even larger. What are the factors that determine the charge for an ion in an ionic compound?

Using only the periodic table, predict the most stable ion for Na, Mg. Al, \(\mathrm{S}, \mathrm{Cl}, \mathrm{K}, \mathrm{Ca},\) and Ga. Arrange these from largest to smallest radius, and explain why the radius varies as it does. Compare your predictions with Fig. \(3-5\).

The bond energy for a \(\mathrm{C}-\mathrm{H}\) bond is about \(413 \mathrm{kJ} / \mathrm{mol}\) in \(\mathrm{CH}_{4}\) but \(380 \mathrm{kJ} / \mathrm{mol}\) in \(\mathrm{CHBr}_{3}\). Although these values are relatively close in magnitude, they are different. Explain why they are different. Does the fact that the bond energy is lower in CHBr \(_{3}\) make any sense? Why?

Consider the following statement: "Because oxygen wants to have a negative two charge, the second electron affinity is more negative than the first." Indicate everything that is correct in this statement. Indicate everything that is incorrect. Correct the incorrect statements and explain.

Which has the greater bond lengths: \(\mathrm{NO}_{2}^{-}\) or \(\mathrm{NO}_{3}^{-} ?\) Explain.

The second electron affinity values for both oxygen and sulfur are unfavorable (positive). Explain.

What is meant by a chemical bond? Why do atoms form bonds with each other? Why do some elements exist as molecules in nature instead of as free atoms?

Why are some bonds ionic and some covalent?

How does a bond between Na and Cl differ from a bond between \(\mathrm{C}\) and \(\mathrm{O}\) ? What about a bond between \(\mathrm{N}\) and \(\mathrm{N} ?\)

Does a Lewis structure tell which electrons come from which atoms? Explain.

Evaluate each of the following as an acceptable name for Water: a. dihydrogen oxide b. hydroxide hydride c. hydrogen hydroxide d. oxygen dihydride

Why do we call \(\mathrm{Ba}\left(\mathrm{NO}_{3}\right)_{2}\) barium nitrate, but we call \(\mathrm{Fe}\left(\mathrm{NO}_{3}\right)_{2}\) iron(II) nitrate?

Why is calcium dichloride not the correct systematic name for \(\mathrm{CaCl}_{2} ?\)

The common name for \(\mathrm{NH}_{3}\) is ammonia. What would be the systematic name for \(\mathrm{NH}_{3}\) ? Support your answer.

Compare and contrast the bonding found in the \(\mathrm{H}_{2}(g)\) and HF \((g)\) molecules with that found in NaF(s).

Describe the type of bonding that exists in the \(\mathrm{Cl}_{2}(g)\) molecule. How does this type of bonding differ from that found in the HCl(g) molecule? How is it similar?

Some of the important properties of ionic compounds are as follows: i. Iow electrical conductivity as solids and high conductivity in solution or when molten ii. relatively high melting and boiling points iii. brittleness iv. solubility in polar solvents How does the concept of ionic bonding discussed in this chapter account for these properties?

Distinguish between the following terms. a. molecule versus ion b. covalent bonding versus ionic bonding c. molecule versus compound d. anion versus cation

What is the electronegativity trend? Where does hydrogen fit into the electronegativity trend for the other elements in the periodic table?

When comparing the size of different ions, the general radii trend discussed in Chapter 2 is usually not very useful. What do you concentrate on when comparing sizes of ions to each other or when comparing the size of an ion to its neutral atom?

In general, the higher the charge on the ions in an ionic compound, the more favorable the lattice energy. Why do some stable ionic compounds have +1 charged ions even though \(+4,+5,\) and +6 charged ions would have a more favorable lattice energy?

Combustion reactions of fossil fuels provide most of the energy needs of the world. Why do the combustion reactions of fossil fuels produce so much energy?

Which of the following statements is/are true? Correct the false statements. a. It is impossible to satisfy the octet rule for all atoms in \(\mathrm{XeF}_{2}\) b. Because \(\mathrm{SF}_{4}\) exists, OF, should also exist, because oxygen is in the same family as sulfur. c. The bond in \(\mathrm{NO}^{+}\) should be stronger than the bond in \(\mathrm{NO}^{-}\) d. As predicted from the two Lewis structures for ozone, one oxygen-oxygen bond is stronger than the other oxygen-oxygen bond.

Three resonance structures can be drawn for \(\mathrm{CO}_{2}\). Which resonance structure is best from a formal charge standpoint?

By analogy with phosphorus compounds, name the following: \(\mathrm{Na}_{3} \mathrm{AsO}_{4}, \mathrm{H}_{3} \mathrm{AsO}_{4}, \mathrm{Mg}_{3}\left(\mathrm{SbO}_{4}\right)_{2}\)

Each of the following compounds has three possible names listed for it. For each compound, what is the correct name and why aren't the other names used? a. \(\mathrm{N}_{2} \mathrm{O}\) : nitrogen oxide, nitrogen(I) oxide, dinitrogen monoxide b. \(\mathrm{Cu}_{2} \mathrm{O}:\) copper oxide, copper(I) oxide, dicopper monoxide c. \(\mathrm{Li}_{2} \mathrm{O}:\) lithium oxide, lithium(I) oxide, dilithium monoxide

Without using Fig. \(3-4,\) predict the order of increasing electronegativity in each of the following groups of elements. a. \(C, N, O\) b. \(\mathbf{s}, \mathbf{S e}, \mathbf{C l}\) \(\mathbf{c}_{*} \mathrm{Si}, \mathrm{Ge}, \mathrm{Sn}\) d. \(\mathrm{TI}, \mathrm{S}, \mathrm{Ge}\)

Without using Fig. \(3-4,\) predict the order of increasing electronegativity in each of the following groups of elements. a. \(\mathrm{Na}, \mathrm{K}, \mathrm{Rb}\) \(\mathbf{b} . \mathbf{B}, \mathbf{O}, \mathbf{G a}\) c. \(F, C\), \(B r\) d. \(s, O, F\)

Without using Fig. \(3-4,\) predict which bond in each of the following groups will be the most polar. a. \(C-F, S i-F, G e-F\) b. \(P-C\) or \(S-C\) \(\mathbf{c} . \mathbf{S}-\mathbf{F}, \mathbf{S}-\mathbf{C} \mathbf{l}, \mathbf{S}-\mathbf{B r}\) d. \(\mathrm{Ti}-\mathrm{Cl}, \mathrm{Si}-\mathrm{Cl}, \mathrm{Ge}-\mathrm{Cl}\)

Without using Fig. \(3-4,\) predict which bond in each of the following groups will be the most polar. a. \(\mathrm{C}-\mathrm{H}, \mathrm{Si}-\mathrm{H}, \mathrm{Sn}-\mathrm{H}\) \(\mathbf{b .}\) Al- \(\mathbf{B r}, \mathbf{G a}-\mathbf{B r}, \operatorname{In}-\mathbf{B r}, \mathbf{T}-\mathbf{B r}\) c. \(C-O\) or \(S i-O\) d. \(\mathrm{O}-\mathrm{F}\) or \(\mathrm{O}-\mathrm{Cl}\)

Indicate the bond polarity (show the partial positive and partial negative ends) in the following bonds. a. \(C-O\) b. \(P-H\) \(\mathbf{c} . \quad \mathbf{H}-\mathbf{C l}\) d. \(\mathrm{Br}-\mathrm{Te}\) \(\mathbf{e} . \mathbf{S e}-\mathbf{S}\)

Predict the type of bond (ionic, covalent, or polar covalent) one would expect to form between the following pairs of elements. a. \(\mathrm{Rb}\) and \(\mathrm{Cl}\) b. \(S\) and \(S\) c. \(C\) and \(F\) d. Ba and S e. \(\mathrm{N}\) and \(\mathrm{P}\) f. \(B\) and \(H\)

List all the possible bonds that can occur between the elements \(\mathrm{P}, \mathrm{Cs}, \mathrm{O},\) and \(\mathrm{H} .\) Predict the type of bond (ionic, covalent, or polar covalent) one would expect to form for each bond.

Hydrogen has an electronegativity value between boron and carbon and identical to phosphorus. With this in mind, rank the following bonds in order of decreasing polarity: \(\mathrm{P}-\mathrm{H}\) \(\mathbf{O}-\mathbf{H}, \mathbf{N}-\mathbf{H}, \mathbf{F}-\mathbf{H}, \mathbf{C}-\mathbf{H}\)

Rank the following bonds in order of increasing ionic character: \(\mathrm{N}-\mathrm{O}, \mathrm{Ca}-\mathrm{O}, \mathrm{C}-\mathrm{F}, \mathrm{Br}-\mathrm{Br}, \mathrm{K}-\mathrm{F}\)

Would you expect each of the following atoms to gain or lose electrons when forming ions? What ion is the most likely in each case? a. Ra b. In c. \(P\) d. \(T e\) e. \(\mathrm{Br}\) f. \(\mathrm{Rb}\)

For each of the following atomic numbers, use the periodic table to write the formula (including the charge) for the simple ion that the element is most likely to form in ionic compounds. a. 13 b. 34 c. 56 d. 7 e. 87 f. 35

Write electron configurations for the most stable ion formed by each of the elements Al, Ba, Se, and I (when in stable ionic compounds).

Write electron configurations for the most stable ion formed by each of the elements Te, \(\mathrm{Cl}, \mathrm{Sr},\) and \(\mathrm{Li}\) (when in stable ionic compounds).

Predict the empirical formulas of the ionic compounds formed from the following pairs of elements. Name each compound. a. Li and \(N\) b. Ga and 0 c. Rb and Cl d. Ba and S

Predict the empirical formulas of the ionic compounds formed from the following pairs of elements. Name each compound. a. Al and Cl b. Na and O c. Sr and F d. Ca and Se

Write electron configurations for a. the cations \(\mathrm{Mg}^{2+}, \mathrm{K}^{+},\) and \(\mathrm{Al}^{3+}\) b. the anions \(\mathrm{N}^{3-}, \mathrm{O}^{2-}, \mathrm{F}^{-},\) and \(\mathrm{Te}^{2-}\)

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

Which of the following ions have noble gas electron configurations? a. \(\mathrm{Fe}^{2+}, \mathrm{Fe}^{3+}, \mathrm{Sc}^{3+}, \mathrm{Co}^{3+}\) b. \(\mathrm{Tl}^{+}, \mathrm{Te}^{2-}, \mathrm{Cr}^{3+}\) c. \(\mathrm{Pu}^{4+}, \mathrm{Ce}^{4+}, \mathrm{Ti}^{4+}\) d. \(\mathrm{Ba}^{2+}, \mathrm{Pt}^{2+}, \mathrm{Mn}^{2+}\)

What noble gas has the same electron configuration as each of the ions in the following compounds? a. cesium sulfide b. strontium fluoride c. calcium nitride d. aluminum bromide

Give the formula of a negative ion that would have the same number of electrons as each of the following positive ions. a. \(\mathrm{Na}^{+}\) b. \(\mathrm{Ca}^{2+}\) \(\mathbf{c} . \mathrm{Al}^{3+}\) d. \(\mathbf{R} \mathbf{b}^{+}\)

Give an example of an ionic compound where both the anion and the cation are isoelectronic with each of the following noble gases. a. Ne b. Arr c. Kr d. Xe

Give three ions that are isoelectronic with neon. Place these ions in order of increasing size.

For each of the following groups, place the atoms and/or ions in order of decreasing size. a. \(\mathrm{Cu}, \mathrm{Cu}^{+}, \mathrm{Cu}^{2+}\) b. \(\mathrm{Ni}^{2+}, \mathrm{Pd}^{2+}, \mathrm{Pt}^{2+}\) c. \(\mathrm{O}, \mathrm{O}^{-}, \mathrm{O}^{2-}\) d. \(\mathrm{La}^{3+}, \mathrm{Eu}^{3+}, \mathrm{Gd}^{3+}, \mathrm{Yb}^{3+}\) e. \(\mathrm{Te}^{2-}, \mathrm{I}^{-}, \mathrm{Cs}^{+}, \mathrm{Ba}^{2+}, \mathrm{La}^{3+}\)