Periodic Patterns in the Main-Group Elements

(a) Define the four atomic properties that are reviewed in the Interchapter.

(b) To what do the three parts of the electron configuration (  ) correlate?

(c) Which part of the electron configuration is primarily associated with the size of the atom? 

(d) What is the major distinction between the outer electron configurations within a group and those within a period? 

(e) What correlation, if any, exists between the group number and the number of valence electrons?

How are covalent and metallic bondings similar? How are they different?

If the leftmost element in a period combined with each of the others in the period, how would the type of bonding change from left to right? Explain in terms of atomic properties.

How do each of the following correlate with the electronegativity of the elements: 

(a) the type of bonding in element oxides; 

(b) the acid-base behaviour of element oxides?

How are covalent and metallic bondings similar? How are they different?

If the leftmost element in a period combined with each of the others in the period, how would the type of bonding change from left to right? Explain in terms of atomic properties.

How do the physical properties of a network covalent solid and a molecular covalent solid differ? Why?

(a) What trends, if any, exist for Zeff across a period and down a group? 

(b) How does Zeff influence atomic size, IE1, and EN across a period?

Iodine monochloride and elemental bromine have nearly the same molar mass and liquid density but very different boiling points. 

(a) What molecular property is primarily responsible for this difference in boiling point? What atomic property gives rise to it? Explain. 

(b) Which substance has a higher boiling point? Why?

How does the trend in atomic size differ from the trend in ionization energy? Explain.

How are bond energy, bond length, and reactivity related to similar compounds?

Why is rotation about the bond axis possible for single-bonded atoms but not double-bonded atoms?

Explain the horizontal irregularity in the size of the most common ions of Period 3 elements.

(a) How does the metallic character of an element correlate with the acidity of its oxide?

 (b) What trends, if any, exist in oxide basicity across a period and down a group

How are atomic size, IE1, and EN related to the redox behaviour of the elements in Groups 1A(1), 2A(2), 6A(16), and 7A(17)?

Rank the following elements in order of increasing 

(a) Atomic size: Ba, Mg, Sr 

(b) IE1: P, Na, Al 

(c) EN: Br, Cl, Se 

(d) No. of valence electrons: Bi, Ga, Sn

Which member of each pair gives the more acidic solution in water: (a) CO₂ or SrO; (b) SnO or SnO₂; (c) Cl₂O or Na₂O; (d) SO₂ or MgO? For the pair in part (a), write an equation for the dissolving of each oxide to support your answer.

What oxidation slates does $\mathbf{\text{Xe}}$ show in its compounds?

Why do the noble gases have such low boiling points?

Explain why $\mathbf{\text{Xe}}$, and to a limited extent $\mathbf{Kr}$, form compounds, whereas $\mathbf{\text{He}}$, $\mathbf{Ne}$, and $\mathbf{Ar}$ do not.

(a) Why do stable xenon fluorides have an even number of $\mathbf{\text{F}}$ atoms?

(b) Why do the ionic species  ${{\mathbf{\text{XeF}}}_{\mathbf{3}}}^{\mathbf{+}}$ and ${{\mathbf{\text{XeF}}}_{\mathbf{7}}}^{\mathbf{-}}$  have odd numbers of $\mathbf{\text{F}}$ atoms? 

(c) Predict the shape of ${{\mathbf{\text{XeF}}}_{\mathbf{3}}}^{\mathbf{+}}$.

Rank the following in order of decreasing

(a) Boiling point: O₂, Br₂, As(s)

(b) ${\mathbf{\Delta H}}_{\mathbf{vap}}$: Cl₂, Ar, I₂

Sketch a periodic table, and label the areas containing elements that give rise to the three types of hydrides.

Draw Lewis structures for the following compounds, and predict which member of each pair will form hydrogen bonds:

(a) NF₃ or NH₃ (b) CH₃OCH₃ or CH₃CH₂OH

Draw Lewis structures for the following compounds and predict which member of each pair will form hydrogen bonds:

(a) NH₃ or AsH₃    

 (b) CH₄ or H₂O

Complete and balance the following equations:

(a) An active metal reacting with acid:

$\mathbf{Al}\mathbf{\left(}\mathbf{s}\mathbf{\right)}\mathbf{+}\mathbf{HCl}\mathbf{\left(}\mathbf{aq}\mathbf{\right)}\mathbf{\to }$ 

(b) A saltlike (alkali metal) hydride reacting with water:

$\mathbf{LiH}\mathbf{\left(}\mathbf{s}\mathbf{\right)}\mathbf{+}{\mathbf{H}}_{\mathbf{2}}\mathbf{O}\mathbf{\left(}\mathbf{l}\mathbf{\right)}\mathbf{\to }$ 

Fluorine lies between oxygen and neon in Period 2. Whereas atomic sizes and ionization energies of these three elements change smoothly, their electronegativities display a dramatic change. What is this change, and how do their electron configurations explain it?

Hydrogen has only one proton, but its IE₁is much greater than that of lithium, which has three protons. Explain.

Lithium salts are often much less soluble in water than the corresponding salts of other alkali metals. For example, at 18°C, the concentration of a saturated LiF solution is $\mathbf{1}\mathbf{.}\mathbf{0}\mathbf{\times }{\mathbf{10}}^{\mathbf{-}\mathbf{2}}$M, whereas that of a saturated KF solution is 1.6 M. How would you explain this behaviour?

Which member of each pair has more ionic character in its bonds: 

(a) BeF₂ or CaF₂; (b) PbF₂ or PbF₄; (c) GeF₄ or PF₃?

Which member of each pair has more covalent character in

its bonds: (a) LiCl or KCl; (b) AlCl₃ or PCl₃; (c) NCl₃ or AsCl₃?

Which member of each pair has more ionic character in its bonds: 

(a) BeF₂ or CaF₂; (b) PbF₂ or PbF₄; (c) GeF₄ or PF₃?

Give the name and symbol or formula of a Group 3A (13) element or compound that fits each description or use: 

(a) Component of heat-resistant (Pyrex-type) glass 

(b) Manufacture of high-speed computer chips 

(c) Largest temperature range for liquid state of an element 

(d) Elemental substance with three-center, two-electron bonds 

(e) Metal protected from oxidation by adherent oxide coat 

(f) Mild antibacterial agent (e.g., for eye infections) 

(g) Toxic metal that lies between two other toxic metals

Rank the following oxides in the order of increasing aqueous acidity: $\mathit{G}{\mathit{a}}_{\mathbf{2}}{\mathit{O}}_{\mathbf{3}}\mathbf{,}\mathbf{ }\mathit{A}{\mathit{l}}_{\mathbf{2}}{\mathit{O}}_{\mathbf{3}}\mathbf{,}\mathbf{ }\mathit{I}{\mathit{n}}_{\mathbf{2}}{\mathit{O}}_{\mathbf{3}}$.

Rank the following hydroxides in order of increasing aqueous

basicity: $\mathit{A}\mathit{l}{\left(OH\right)}_{\mathbf{3}}\mathbf{,}\mathit{B}{\left(OH\right)}_{\mathbf{3}}\mathbf{,}\mathit{I}\mathit{n}{\left(OH\right)}_{\mathbf{3}}$

Thallium forms the compound $\mathit{T}\mathit{l}{\mathit{I}}_{\mathbf{3}}$. What is the apparent oxidation state of Tl in this compound? Given that the anion is ${{\mathbf{I}}_{\mathbf{3}}}^{\mathbf{-}}$, what is the actual oxidation state of Tl? Draw the shape of the anion, giving its VSEPR class and bond angles. Propose a reason why the compound does not exist as $\left(T{l}^{3+}\right){\left(I^{-}\right)}_{\mathbf{3}}$.

Indium (In) reacts with HCl to form a diamagnetic solid with the formula $\mathit{I}\mathit{n}\mathit{C}{\mathit{l}}_{\mathbf{2}}$ . 

(a) Write condensed electron configurations for $\mathit{I}\mathit{n}\mathbf{,}\mathbf{ }\mathit{I}{\mathit{n}}^{\mathbf{+}}\mathbf{,}\mathbf{ }\mathit{I}{\mathit{n}}^{\mathbf{2}\mathbf{+}}$and $\mathit{I}{\mathit{n}}^{\mathbf{3}\mathbf{+}}$. 

(b) Which of these species is (are) diamagnetic and which paramagnetic? 

(c) What is the apparent oxidation state of In in $\mathit{I}\mathit{n}\mathit{C}{\mathit{l}}_{\mathbf{2}}$?

(d) Given your answers to parts (b) and (c), explain how $\mathit{I}\mathit{n}\mathit{C}{\mathit{l}}_{\mathbf{2}}$can be diamagnetic.

Nearly every compound of silicon has the element in the $\mathbf{+}\mathbf{4}$ oxidation state. In contrast, most compounds of lead have the element in the $\mathbf{+}\mathbf{2}$ state.

(a) What general observation do these facts illustrate?

(b) Explain in terms of atomic and molecular properties.

(c) Give an analogous example from Group $\mathbf{3}\mathit{A}\mathbf{\left(}\mathbf{13}\mathbf{\right)}$.

Boron nitride (BN) has a structure similar to graphite, but is a white insulator rather than a black conductor. It is synthesized by heating diboron trioxide with ammonia at about $\mathbf{1000}\mathbf{°}\mathit{C}$.

(a) Write a balanced equation for the formation of BN; water forms also.

(b) Calculate $\mathbf{\Delta }{\mathbf{H}\mathbf{°}}_{\mathbf{rxn}}$ for the production of BN ($\mathbf{\Delta }{\mathbf{H}\mathbf{°}}_{\mathbf{f}}$ of BN is $\mathbf{-}\mathbf{254}{\mathbf{kJmol}}^{\mathbf{-}\mathbf{1}}$).

(c) Boron is obtained from the mineral borax, ${\mathbf{Na}}_{\mathbf{2}}{\mathbf{B}}_{\mathbf{4}}{\mathbf{O}}_{\mathbf{7}}\mathbf{.}\mathbf{10}{\mathbf{H}}_{\mathbf{2}}\mathbf{O}$. How much borax is needed to produce 0.1kg of BN, assuming 72% yield?

The sum of ${\mathbf{IE}}_{\mathbf{1}}$ through ${\mathbf{IE}}_{\mathbf{4}}$ for group 4A (14) elements shows a decrease from C to Si,  a slight increase from Si to Ge, a decrease from Ge to Sn and an increase from Sn to Pb.

1. What is the expected trend for IEs down a group?
2. Suggest a reason for the deviations in group 4 A (14).
3. Which group might show even greater deviations?

Give explanations for the large drops in melting point from C to Si and from Ge to Sn.

What is an allotrope? Name two Group 4 A (14) elements that exhibit allotropism and identify two of their allotropes.

Even though EN values vary relatively little down a group 4A (14), the elements change from nonmetal to metal. Explain.

How do atomic properties account for the enormous number of carbon compounds? Why don’t other group 4A(14) elements behave similarly?

Question: (a) What is the range of oxidation states shown by the elements of Group 5A (15) as you move down the group? 

(b) How does this range illustrate the general rule for the range of oxidation states in groups on the right side of the periodic table?

One similarity between B and Si is the explosive combustion of their hydrides in air. Write balanced equations for the combustion of ${\mathbf{B}}_{\mathbf{2}}{\mathbf{H}}_{\mathbf{6}}$ and  ${\mathbf{Si}}_{\mathbf{4}}{\mathbf{H}}_{\mathbf{10}}$ .

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

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

(a) What is the range of oxidation states shown by the elements of Group 5A (15) as you move down the group? 

(b) How does this range illustrate the general rule for the range of oxidation states in groups on the right side of the periodic table?

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?