Periodic Patterns in the Main-Group Elements

Rank the following oxides in order of increasing acidity in water: ${\mathbf{Sb}}_{\mathbf{2}}{\mathbf{O}}_{\mathbf{3}}$, ${\mathbf{Bi}}_{\mathbf{2}}{\mathbf{O}}_{\mathbf{3}}$, ${\mathbf{P}}_{\mathbf{4}}{\mathbf{O}}_{\mathbf{10}}$, ${\mathbf{Sb}}_{\mathbf{2}}{\mathbf{O}}_{\mathbf{5}}$ .

Assuming acid strength relates directly to the electronegativity of the central atom, rank ${\mathbf{H}}_{\mathbf{3}}{\mathbf{PO}}_{\mathbf{4}}$, ${\mathbf{HNO}}_{\mathbf{3}}$, and ${\mathbf{H}}_{\mathbf{3}}{\mathbf{AsO}}_{\mathbf{4}}$ in order of increasing acid strength.

Question: As you move down Group 5A (15), the melting points of the elements increases and then decreases. Explain

Question: Assuming acid strength relates directly to the electronegativity of the central atom, rank H₃PO₄,HNO₃ , and H₃ASO₄ in order of increasing acid strength.

Question: One similarity between B and Si is the explosive combustion of their hydrides in air. Write balanced equations for the combustion of B₂H₆ and of Si₄H₁₀.

Question: Complete and balance the following: 

$$\begin{gathered} \mathbf{a}\mathbf{.}\mathbf{ }\mathbf{As}\left(\mathbf{s}\right)\mathbf{+}{\mathbf{excessO}}_{\mathbf{2}}\left(\mathbf{g}\right)\to \\ b. \mathbf{Bi}\left(\mathbf{s}\right)\mathbf{+}{\mathbf{excessF}}_{\mathbf{2}}\left(\mathbf{g}\right)\to \\ c. {\mathbf{Ca}}_{\mathbf{3}}{\mathbf{As}}_{\mathbf{2}}\left(\mathbf{s}\right)\mathbf{+}{\mathbf{H}}_{\mathbf{2}}\mathbf{O}\left(\mathbf{l}\right)\to \end{gathered}$$

Question: Which Group 5A (15) elements form trihalides? Pentahalides? Explain.

Hydrogen peroxide can act as either an oxidizing agent or a reducing agent. (a) When ${\mathbf{H}}_{\mathbf{2}}{\mathbf{O}}_{\mathbf{2}}$   is treated with aqueous KI, ${\mathbf{I}}_{\mathbf{2}}$ forms. In which role  ${\mathbf{H}}_{\mathbf{2}}{\mathbf{O}}_{\mathbf{2}}$ is acting? What is the oxygen- containing product formed? (b) When ${\mathbf{H}}_{\mathbf{2}}{\mathbf{O}}_{\mathbf{2}}$ is treated with acidic ${\mathbf{KMnO}}_{\mathbf{4}}$, the purple color of $\mathbf{Mn}{{\mathbf{O}}_{\mathbf{4}}}^{\mathbf{-}}$disappears and a gas forms . In which role is ${\mathbf{H}}_{\mathbf{2}}{\mathbf{O}}_{\mathbf{2}}$  acting? What is the oxygen-containing product formed?

a) Give the physical state and color of each halogen at STP.

Question: Draw a Lewis structure for 

1. The cyclic silicate ion ${\mathbf{Si}}_{\mathbf{4}}{{\mathbf{O}}_{\mathbf{12}}}^{\mathbf{-}\mathbf{8}}$ 
2. A cyclic hydrocarbon with formula ${\mathbf{C}}_{\mathbf{4}}{\mathbf{H}}_{\mathbf{8}}$ 

Draw a Lewis structure for 

1. The cyclic silicate ion ${\mathbf{Si}}_{\mathbf{6}}{{\mathbf{O}}_{\mathbf{18}}}^{\mathbf{-}\mathbf{12}}$   
2. A cyclic hydrocarbon with formula ${\mathbf{C}}_{\mathbf{6}}{\mathbf{H}}_{\mathbf{12}}$  

Show two chains of three units of each of a sheet silicon polymer made from $\mathbf{(}\mathbf{C}{\mathbf{H}}_{\mathbf{3}}{\mathbf{)}}_{\mathbf{2}}\mathbf{Si}{\left(\mathbf{OH}\right)}_{\mathbf{2}}$ with $\mathbf{(}\mathbf{C}{\mathbf{H}}_{\mathbf{3}}{\mathbf{)}}_{\mathbf{2}}\mathbf{Si}{\left(\mathbf{OH}\right)}_{\mathbf{3}}$ added to crosslink the chains.

Bismuth(V) compounds are such powerful oxidizing agents that they have not been prepared in pure form. How is this fact consistent with the location of Bi in the periodic table?

Complete and balance the following: 

(a) ${\mathbf{\text{AlCl}}}_{\mathbf{\text{3}}}\mathbf{\left(}\mathbf{\text{I}}\mathbf{\right)}{\mathbf{\text{+H}}}_{\mathbf{\text{2}}}\mathbf{\text{O}}\mathbf{\left(}\mathbf{\text{I}}\mathbf{\right)}\mathbf{\to }$

(b) ${\mathbf{\text{Sb}}}_{\mathbf{\text{2}}}{\mathbf{\text{O}}}_{\mathbf{\text{3}}}\mathbf{\left(}\mathbf{\text{s}}\mathbf{\right)}\mathbf{\text{+NaOH}}\mathbf{\left(}\mathbf{\text{aq}}\mathbf{\right)}\mathbf{\to }$

Based on the relative sizes of F and Cl, predict the structure of  ${\mathbf{PF}}_{\mathbf{2}}{\mathbf{Cl}}_{\mathbf{3}}$

Use the VSEPR model to predict the structure of the cyclic ion ${\mathbf{P}}_{\mathbf{3}}{{\mathbf{O}}_{\mathbf{9}}}^{\mathbf{3}\mathbf{-}}$

In addition to those in Table 14.3, other less stable nitrogen oxides exist. Draw a Lewis structure for each of the following: 

(a) ${\mathbf{\text{N}}}_{\mathbf{\text{2}}}{\mathbf{\text{O}}}_{\mathbf{\text{2}}}$, a dimer of nitrogen monoxide with an $\mathbf{\text{N-N}}$ bond.

(b) ${\mathbf{\text{N}}}_{\mathbf{\text{2}}}{\mathbf{\text{O}}}_{\mathbf{\text{2}}}$, a dimer of nitrogen monoxide with no $\mathbf{\text{N-N}}$ bond.

(c) ${\mathbf{\text{N}}}_{\mathbf{\text{2}}}{\mathbf{\text{O}}}_{\mathbf{\text{3}}}$ with no $\mathbf{\text{N-N}}$ bond.

(d)  ${\mathbf{\text{NO}}}^{\mathbf{+}}$ and  ${{\mathbf{\text{NO}}}_{\mathbf{\text{3}}}}^{\mathbf{-}}$ , products of the ionization of liquid  ${\mathbf{N}}_2{\mathbf{O}}_4$.

Write balanced chemical equations for the thermal decomposition of potassium nitrate (oxygen is also formed in both the cases): (a) at low temperature to the nitrite; (b) at high temperature to the metal oxide and nitrogen.

Give the name and symbol or formula of a Group 6A(16) element or compound that fits each description or use: (a) Unreactive gas used as an electrical insulator (b) Unstable allotrope of oxygen (c) Oxide having sulfur with the same O.N. as in sulfuric acid (d) Air pollutant produced by burning sulfur-containing coal (e) Powerful dehydrating agent (f) Compound used in solution in the photographic process (g) Gas in trace amounts in air that tarnishes silver

 Give the oxidation state of sulfur in (a)${\mathbf{S}}_{\mathbf{8}}$ ; (b)${\mathbf{SF}}_{\mathbf{4}}$   ; (c)${\mathbf{SF}}_{\mathbf{6}}$  ; (d)${\mathbf{H}}_{\mathbf{2}}\mathbf{S}$  ; (e)${\mathbf{FeS}}_{\mathbf{2}}$; (f)${\mathbf{H}}_{\mathbf{2}}{\mathbf{SO}}_{\mathbf{4}}$  ; (g) ${\mathbf{Na}}_{\mathbf{2}}{\mathbf{S}}_{\mathbf{2}}{\mathbf{O}}_{\mathbf{3}}\mathbf{·}\mathbf{5}{\mathbf{H}}_{\mathbf{2}}\mathbf{O}$ .

Bromine monofluoride (BF) disproportionates to bromine gas and bromine tri- and pentafluorides. Use the following to find Δ${{\mathbf{H}}^{\mathbf{o}}}_{\mathbf{rxn}}$for the decomposition of BrF to its elements:

$$\begin{gathered} \mathbf{3}{\mathbf{BrF}}_{\left(g\right)}\mathbf{\to }\mathbf{B}{{\mathbf{r}}_{\mathbf{2}}}_{\left(g\right)}\mathbf{+}{\mathbf{BrF}}_{{\mathbf{3}}_{\left(l\right)}}\mathbf{ }\mathbf{ }\mathbf{ }\mathbf{ }\mathbf{ }\mathbf{\Delta }{{\mathbf{H}}^{\mathbf{o}}}_{\mathbf{rxn}}\mathbf{=}\mathbf{-}\mathbf{125}\mathbf{.}\mathbf{3}\mathbf{KJ}\mathbf{\backslash }\mathbf{hfill} \\ \mathbf{5}{\mathbf{BrF}}_{\left(g\right)}\mathbf{\to }\mathbf{2}\mathbf{B}{{\mathbf{r}}_{\mathbf{2}}}_{\left(g\right)}\mathbf{+}{\mathbf{BrF}}_{{\mathbf{5}}_{\left(l\right)}}\mathbf{ }\mathbf{ }\mathbf{ }\mathbf{\Delta }{{\mathbf{H}}^{\mathbf{o}}}_{\mathbf{rxn}}\mathbf{=}\mathbf{-}\mathbf{166}\mathbf{.}\mathbf{1}\mathbf{KJ}\mathbf{\backslash }\mathbf{hfill} \\ {\mathbf{BrF}}_{{\mathbf{3}}_{\left(l\right)}}\mathbf{+}{\mathbf{F}}_{\mathbf{2}}\mathbf{\to }\left(g\right){\mathbf{BrF}}_{{\mathbf{5}}_{\left(l\right)}}\mathbf{ }\mathbf{ }\mathbf{ }\mathbf{ }\mathbf{ }\mathbf{ }\mathbf{\Delta }{{\mathbf{H}}^{\mathbf{o}}}_{\mathbf{rxn}}\mathbf{=}\mathbf{-}\mathbf{158}\mathbf{.}\mathbf{0}\mathbf{KJ}\mathbf{\backslash }\mathbf{hfill} \end{gathered}$$

Which noble gas is the most abundant in the universe? In Earth’s atmosphere?

Rank the following acids in order of decreasing acid strength: $HBr{O}_3$ , $HBr{O}_4$ , $HI{O}_3$ , $HCl{O}_4$ 

Complete and balance the following equations. If no reaction occurs, write NR:(a)$Rb\left(s\right)+B{r}_2\left(l\right)\to$

 (b) $I_2\left(s\right)+H_{2}O\left(l\right)\to$ 

 (c) $B{r}_2\left(l\right)+I^{-}\left(aq\right)\to$ 

 (d)   $Ca{F}_2\left(s\right)+H_{2}S{O}_4\left(l\right)\to$

In addition to interhalogen compounds, many interhalogen ions exist. Would you expect interhalogen ions with a 1+ or a 1- charge to have an even or odd number of atoms? Explain.

Select the stronger bond in each pair: 

(a) Cl-Cl or Br-Br 

(b) Br-Br or I-I 

(c) F-F or Cl-Cl. 

Why doesn't the F-F bond strength follow the group trend?

An industrial chemist treats solid NaCl with concentrated  $H_{2}S{O}_4$ and obtains gaseous HCl and $NaHS{O}_4$ . When she substitutes solid NaI for NaCl, gaseous $H_{2}S$  , solid $I_2$ , and $S_8$  are obtained but no HI. (a) What type of reaction did the $H_{2}S{O}_4$  undergo with NaI? (b) Why does NaI, but not NaCl, cause this type of reaction? (c) To produce HI(g) by the reaction of NaI with an acid, how does the acid have to differ from sulfuric acid?

(a) Give the physical state and color of each halogen at STP. (b) Explain the change in physical state down Group 7A(17) in terms of molecular properties.

(a) What are the common oxidation states of the halogens? (b) Give an explanation based on electron configuration for the range and values of the oxidation states of chlorine. (c) Why is fluorine an exception to the pattern of oxidation states found for the other group members?

Two substances with empirical formula HNO are hyponitrous acid  $\left( \mu =62.04g/\mathrm{mol}\right)$ and nitroxyl $\left( \mu =31.02g/\mathrm{mol}\right)$ . 

(a) What is the molecular formula of each species? 

(b) For each species, draw the Lewis structure having the lowest formal charges. (Hint: Hyponitrous acid has an N N bond.) 

(c) Predict the shape around the N atoms of each species.

(d) When hyponitrous acid loses two protons, it forms the hyponitrite ion. Draw cis and trans forms of this ion.

The main reason alkali metal dihalides (MX2) do not form is the high  IE₂ of the metal.

 (a) Why is IE₂ so high for alkali metals?

 (b) The IE₂ for Cs is 2255 kJ/mol, low enough for CsF2 to form exothermically $\left(\Delta {\mathbf{H}}^0\mathbf{f} =-\mathbf{125}\mathbf{kJ}/\mathbf{mol}\right)$ . This compound cannot be synthesized, however, because CsF forms with a much greater release of heat $\left(\Delta H^{0}f =-530\mathrm{kJ}/\mathrm{mol}\right)$  . Thus, the breakdown of CsF2 to CsF happens readily. Write the equation for this breakdown, and calculate the heat of reaction per mole of CsF.

Thionyl chloride (SOCl₂) is a sulfur oxohalide used industrially to dehydrate metal halide hydrates. 

(a) Write a balanced equation for its reaction with magnesium chloride hexahydrate, in which SO₂  and HCl form along with the metal halide. 

(b) Draw a Lewis structure of (SOCl₂)  with minimal formal charges.

Cyclopropane (C₃H₆), the smallest cyclic hydrocarbon, is reactive for the same reason white phosphorus is. Use bond properties and valence bond theory to explain its high reactivity.

The interhalogen IF undergoes the reaction depicted below (I is purple and F is green):

(a) Write the balanced equation. 

(b) Name the interhalogen product. 

(c) What type of reaction is shown? 

(d) If each molecule of IF represents $\mathbf{2}.\mathbf{50}\times {\mathbf{10}}^{-\mathbf{3}}$ mol, what mass of each product forms?

Unlike other Group 2A(2) metals, beryllium reacts like aluminum and zinc with concentrated aqueous base to release hydrogen gas and form oxoanions of formula $\mathbf{M}{\left(\mathbf{OH}\right)}_{\mathbf{4}}{\mathbf{n}}^{-}$ . Write equations for the reactions of these three metals with NaOH.

The electronic transition in Na from 3p¹ to 3s¹ gives rise to a bright yellow-orange emission at 589.2 nm. What is the energy of this transition?

Xenon tetrafluoride reacts with antimony pentafluoride to form the following ionic complex: ${\mathbf{\left[}\mathbf{XeF}\mathbf{3}\mathbf{\right]}}^{\mathbf{+}}{\mathbf{\left[}\mathbf{SbF}\mathbf{6}\mathbf{\right]}}^{\mathbf{–}}$.

(a) Which of the following illustrates the molecular shapes of the reactants and product? 

(b) How, if at all, does the hybridization of xenon change in the reaction?

Given the following information,

$$\begin{gathered} {\mathbf{H}}^{+}\mathbf{\left(}\mathit{g}\mathbf{\right)}\mathbf{+}{\mathbf{H}}_2\mathbf{O}\mathbf{\left(}\mathit{g}\mathbf{\right)}\mathbf{\to }{\mathbf{H}}_3{\mathbf{O}}^{+}\mathbf{\left(}\mathit{g}\mathbf{\right)}\mathbf{\Delta H}\mathbf{=}\mathbf{-}\mathbf{720}\mathbf{\text{ kJ}} \\ {\mathbf{H}}^{+}\mathbf{\left(}\mathit{g}\mathbf{\right)}\mathbf{+}{\mathbf{H}}_2\mathbf{O}\mathbf{\left(}\mathit{l}\mathbf{\right)}\mathbf{\to }{\mathbf{H}}_3{\mathbf{O}}^{+}\mathbf{\left(}\mathit{a}\mathit{q}\mathbf{\right)}\mathbf{\Delta {\rm H}}\mathbf{=}\mathbf{-}\mathbf{1090}\mathbf{\text{ kJ}} \\ {\mathbf{H}}_2\mathbf{O}\mathbf{\left(}\mathit{l}\mathbf{\right)}\mathbf{\to }{\mathbf{H}}_2\mathbf{O}\mathbf{\left(}\mathit{g}\mathbf{\right)}\mathbf{\Delta H}\mathbf{=}\mathbf{40}\mathbf{.}\mathbf{7}\mathbf{\text{ kJ}} \end{gathered}$$ 

Calculate the heat of solution of the hydronium ion:

${\mathbf{H}}_3{\mathbf{O}}^{+}\mathbf{\left(}\mathbf{g}\mathbf{\right)}\stackrel{{\mathbf{H}}_2\mathbf{O}}{\mathbf{\to }}{\mathbf{H}}_3{\mathbf{O}}^{+}\mathbf{\left(}\mathbf{a}\mathbf{q}\mathbf{\right)}$

Element E forms an oxide of general structure A and a chloride of general structure B:

                      A                                                                                       B

For the anion $\mathbf{E}{{\mathbf{F}}_{\mathbf{5}}}^{\mathbf{-}}$ , what is 

(a) the molecular shape; 

(b) the hybridization of E; 

(c) the O.N. of E?

H2 may act as a reducing agent or an oxidizing agent, depending on the substance reacting with it. Using, in turn, sodium or chlorine as the other reactant, write balanced equations for the two reactions and characterize the redox role of hydrogen.

Cake alum (aluminum sulfate) is used as a flocculating agent in water purification. The “floc” is a gelatinous precipitate of aluminium hydroxide that carries small suspended particles and bacteria out of the solution for removal by filtration. 

(a) The hydroxide is formed by the reaction of the aluminium ion with water. Write the equation for this reaction. 

(b) The acidity from this reaction is partially neutralized by the sulfate ion. Write an equation for this reaction.

Account for the following facts: 

(a)   ${\mathbf{Ca}}^{\mathbf{+}\mathbf{2}}\mathbf{ }\mathbf{and}\mathbf{ }{\mathbf{Na}}^{\mathbf{+}}$ have very nearly the same radii.

(b) ${\mathbf{CaF}}_{\mathbf{2}}$ is insoluble in water, but NaF is quite soluble. 

(c) Molten ${\mathbf{BeCl}}_{\mathbf{2}}$ is a poor electrical conductor, whereas molten   is an excellent one

An important starting material for the manufacture of polyphosphazene is the cyclic molecule (NPCl₂)₃. The molecule has a symmetrical six-membered ring of alternating N and P atoms, with the Cl atoms bonded to the P atoms. The nitrogen phosphorus bond length is significantly less than that expected for an NP single bond. 

(a) Draw a likely Lewis structure for the molecule. 

(b) How many lone pairs of electrons do the ring atoms have? 

(c) What is the order of the nitrogen-phosphorus bond?

From its formula, one might expect CO to be quite polar, but its dipole moment is low (0.11 D). 

(a) Draw the Lewis structure for CO. 

(b) Calculate the formal charges. 

(c) Based on your answers to parts (a) and (b), explain why the dipole moment is so low