Chapter 3

Arrange the following in order of increasing radius and increasing ionization energy. a. \(\mathrm{N}^{+}, \mathrm{N}, \mathrm{N}^{-}\) b. \(\mathrm{Se}, \mathrm{Se}^{-}, \mathrm{Cl}, \mathrm{Cl}^{+}\) c. \(\mathrm{Br}^{-}, \mathrm{Rb}^{+}, \mathrm{Sr}^{2+}\)

For each of the following, write an equation that corresponds to the energy given. a. lattice energy of \(\mathrm{NaCl}\) b. lattice energy of \(\mathrm{NH}_{4} \mathrm{Br}\) c. lattice energy of \(\mathrm{MgS}\) d. \(O=O\) double bond energy beginning with \(O_{2}(g)\) as a reactant

Write Lewis structures for \(\mathrm{CO}_{3}^{2-}, \mathrm{HCO}_{3}^{-},\) and \(\mathrm{H}_{2} \mathrm{CO}_{3}\). When acid is added to an aqueous solution containing carbonate or bicarbonate ions, carbon dioxide gas is formed. We generally say that carbonic acid \(\left(\mathrm{H}_{2} \mathrm{CO}_{3}\right)\) is unstable. Use bond energies to estimate \(\Delta E\) for the reaction (in the gas phase) $$\mathrm{H}_{2} \mathrm{CO}_{3} \longrightarrow \mathrm{CO}_{2}+\mathrm{H}_{2} \mathrm{O}$$Specify a possible cause for the instability of carbonic acid.

Which member of the following pairs would you expect to be more energetically stable? Justify each choice. a. NaBr or \(\mathrm{NaBr}_{2}\) b. \(\mathrm{ClO}_{4}\) or \(\mathrm{ClO}_{4}^{-}\) c. \(\mathrm{SO}_{4}\) or \(\mathrm{XeO}_{4}\) d. \(\mathrm{OF}_{4}\) or \(\mathrm{SeF}_{4}\)

What do each of the following sets of compounds/ions have in common? a. \(\mathrm{SO}_{3}, \mathrm{NO}_{3}^{-}, \mathrm{CO}_{3}^{2-}\) b. \(\mathrm{O}_{3}, \mathrm{SO}_{2}, \mathrm{NO}_{2}^{-}\)

Look up the energies for the bonds in CO and \(\mathrm{N}_{2}\). Although the bond in CO is stronger, CO is considerably more reactive than \(\mathrm{N}_{2}\). Give a possible explanation.

The formulas and common names for several substances are given below. Give the systematic names for these substances. a. sugar of lead b. blue vitrol c. quicklime d. Epsom salts e. milk of magnesia f. gypsum g. laughing gas \(\mathrm{Pb}\left(\mathrm{C}_{2} \mathrm{H}_{3} \mathrm{O}_{2}\right)_{2}\) \(\mathrm{CuSO}_{4}\) CaO \(\mathrm{MgSO}_{4}\) \(\mathrm{Mg}(\mathrm{OH})_{2}\) \(\mathrm{CaSO}_{4}\) \(\mathrm{N}_{2} \mathrm{O}\)

Identify each of the following elements: a. a member of the same family as oxygen whose most stable ion contains 54 electrons b. a member of the alkali metal family whose most stable ion contains 36 electrons c. a noble gas with 18 protons in the nucleus d. a halogen with 85 protons and 85 electrons

A certain element has only two naturally occurring isotopes: one with 18 neutrons and the other with 20 neutrons. The element forms \(1-\) charged ions when in ionic compounds. Predict the identity of the element. What number of electrons does the \(1-\) charged ion have?

The designations \(1 \mathrm{A}\) through \(8 \mathrm{A}\) used for certain families of the periodic table are helpful for predicting the charges on ions in binary ionic compounds. In these compounds, the metals generally take on a positive charge equal to the family number, while the nonmetals take on a negative charge equal to the family number minus \(8 .\) Thus the compound between sodium and chlorine contains \(\mathrm{Na}^{+}\) ions and \(\mathrm{Cl}^{-}\) ions and has the formula NaCl. Predict the formula and the name of the binary compound formed from the following pairs of elements. a. Ca and N b. \(K\) and 0 c. \(\mathrm{Rb}\) and \(\mathrm{F}\) d. \(\mathrm{Mg}\) and \(\mathrm{S}\) e. Ba and I f. Al and Se g. Cs and \(P\) h. In and Br

When molten sulfur reacts with chlorine gas, a vile-smelling orange liquid forms that has a formula of \(\mathrm{S}_{2} \mathrm{Cl}_{2}\). The Lewis structure of this compound has a formal charge of zero on all elements in the compound. Draw the Lewis structure for the vile-smelling orange liquid.

Which of the following statements is(are) correct? a. The symbols for the clements magnesium, aluminum, and xenon are \(\mathrm{Mn}, \mathrm{Al},\) and \(\mathrm{Xe},\) respectively. b. The elements \(P, A s,\) and \(B i\) are in the same family on the periodic table. c. All of the following elements are expected to gain electrons to form ions in ionic compounds: Ga, Se, and Br. d. The elements \(\mathrm{Co}, \mathrm{Ni},\) and \(\mathrm{Hg}\) are all transition elements. e. The correct name for \(\mathrm{TiO}_{2}\) is titanium dioxide.

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

List the bonds \(\mathbf{P}-\mathbf{C l}, \mathbf{P}-\mathbf{F}, \mathbf{O}-\mathbf{F},\) and \(\mathbf{S i}-\mathbf{F}\) from least polar to most polar.

Arrange the atoms and/or ions in the following groups in order of decreasing size. a. \(\mathrm{O}, \mathrm{O}^{-}, \mathrm{O}^{2-}\) b. \(\mathrm{Fe}^{2+}, \mathrm{Ni}^{2+}, \mathrm{Zn}^{2+}\) c. \(\mathrm{Ca}^{2+}, \mathrm{K}^{+}, \mathrm{Cl}^{-}\)

Use the following data to estimate \(\Delta E\) for the reaction: $$\mathrm{Ba}(s)+\mathrm{Br}_{2}(g) \longrightarrow \mathrm{BaBr}_{2}(s) \quad \Delta E=?$$ Lattice energy First ionization energy of Ba Second ionization energy of Ba Electron affinity of Br Bond energy of \(\mathrm{Br}_{2}\) Enthalpy of sublimation of Ba \(-1985 \mathrm{kJ} / \mathrm{mol}\) \(503 \mathrm{kJ} / \mathrm{mol}\) \(965 \mathrm{kJ} / \mathrm{mol}\) -325 kJ/mol \(193 \mathrm{kJ} / \mathrm{mol}\) \(178 \mathrm{kJ} / \mathrm{mol}\)

Which of the following compounds or ions exhibit resonance? a. \(\mathrm{O}_{3}\) b. CNO- c. \(A s I_{3}\) d. \(\mathrm{CO}_{3}^{2-}\) \(\mathbf{e}, \quad A s F_{3}\)

Use Coulomb's law, $$V=\frac{Q_{1} Q_{2}}{4 \pi \epsilon_{0} r}=2.31 \times 10^{-19} \mathrm{J} \cdot \mathrm{nm}\left(\frac{Q_{1} Q_{2}}{r}\right)$$ to calculate the energy of interaction, \(V\), for the following two arrangements of charges, each having a magnitude equal to the electron charge.

An alternative definition of electronegativity is Electronegativity \(=\) constant (I.E. \(-\) E.A.) where I.E. is the ionization energy and E.A. is the electron affinity using the sign conventions of this book. Use data in Chapter 2 to calculate the (I.E. \(-\) E.A.) term for \(\mathbf{F}, \mathbf{C l}, \mathbf{B r}\) and I. Do these values show the same trend as the electronegativity values given in this chapter? The first ionization energies of the halogens are \(1678,1255,1138,\) and \(1007 \mathrm{kJ} / \mathrm{mol},\) respectively. (Hint: Choose a constant so that the electronegativity of fluorine equals \(4.0 .\) Using this constant, calculate relative electronegativities for the other halogens and compare to values given in the text.)

Given the following information: Energy of sublimation of \(\mathrm{Li}(s)=166 \mathrm{kJ} / \mathrm{mol}\) Bond energy of \(\mathrm{HCl}=427 \mathrm{kJ} / \mathrm{mol}\) Ionization energy of \(\mathrm{Li}(g)=520 . \mathrm{kJ} / \mathrm{mol}\) Electron affinity of \(\mathrm{Cl}(g)=-349 \mathrm{kJ} / \mathrm{mol}\) Lattice energy of \(\mathrm{LiCl}(s)=-829 \mathrm{kJ} / \mathrm{mol}\) Bond energy of \(\mathrm{H}_{2}=432 \mathrm{kJ} / \mathrm{mol}\) Calculate the net change in energy for the following reaction: $$ 2 \mathrm{Li}(s)+2 \mathrm{HCl}(g) \longrightarrow 2 \mathrm{LiCl}(s)+\mathrm{H}_{2}(g) $$

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 \(\mathrm{MgO}\) compared to \(\mathrm{MgO}_{2} ?\) \(\mathrm{Mg}_{2} \mathrm{O} ?\)

Use data in this chapter (and Chapter 2 ) to discuss why \(\mathrm{MgO}\) is an ionic compound but CO is not an ionic compound.

Draw a Lewis structure for the \(N, N\) -dimethylformamide molecule. The skeletal structure is Various types of evidence lead to the conclusion that there is some double bond character to the \(\mathrm{C}-\mathrm{N}\) bond. Draw one or more resonance structures that support this observation.

For each of the following ions, indicate the total number of protons and electrons in the ion. For the positive ions in the list, predict the formula of the simplest compound formed between each positive ion and the oxide ion. Name the compounds. For the negative ions in the list, predict the formula of the simplest compound formed between each negative ion and the aluminum ion. Name the compounds. a. \(\mathrm{Fe}^{2+}\) b. \(\mathrm{Fe}^{3+}\) c. \(B a^{2+}\) d. \(C s^{+}\) e. \(S^{2-}\) f. \(P^{3-}\) g. \(\mathrm{Br}^{-}\) \(\mathbf{h} . \mathbf{N}^{3-}\)

Identify the following elements based on their electron configurations and rank them in order of increasing electronegativity: \([\mathrm{Ar}] 4 s^{1} 3 d^{3} ;[\mathrm{Ne}] 3 s^{2} 3 p^{3} ;[\mathrm{Ar}] 4 s^{2} 3 d^{10} 4 p^{3} ;[\mathrm{Ne}] 3 s^{2} 3 p^{5}\)

A polyatomic ion is composed of \(\mathrm{C}, \mathrm{N},\) and an unknown element \(X\). The skeletal Lewis structure of this polyatomic ion is \([\mathrm{X}-\mathrm{C}-\mathrm{N}]^{-} .\) The ion \(\mathrm{X}^{2-}\) has an electron configuration of [Ar]4s^{2} 3 d ^ { 1 0 } 4 p ^ { 6 } . \text { What is element X? Knowing the identity of } X, complete the Lewis structure of the polyatomic ion, including all important resonance structures.