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CHEMISTRY: The Molecular Nature of Matter and Change 2016

Martin S. Silberberg, Patricia G. Amateis

Chapter 14

Periodic Patterns in the Main-Group Elements

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Problem 1

Hydrogen has only one proton, but its $\mathrm{IE}_{1}$ is much greater than that of lithium, which has three protons. Explain.

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Problem 2

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

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Problem 3

Draw Lewis structures for the following compounds, and predict which member of each pair will form hydrogen bonds:
(a) $\mathrm{NF}_{3}$ or $\mathrm{NH}_{3} \quad$ (b) $\mathrm{CH}_{3} \mathrm{OCH}_{3}$ or $\mathrm{CH}_{3} \mathrm{CH}_{2} \mathrm{OH}$

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Problem 4

Draw Lewis structures for the following compounds, and predict which member of each pair will form hydrogen bonds:
(a) $\mathrm{NH}_{3}$ or $\mathrm{AsH}_{3} \quad$ (b) $\mathrm{CH}_{4}$ or $\mathrm{H}_{2} \mathrm{O}$

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Problem 5

Complete and balance the following equations:
(a) An active metal reacting with acid,
$$\mathrm{Al(s)} +\mathrm{HCl}(a q) \longrightarrow$$
(b) A saltlike (alkali metal) hydride reacting with water,
$$\mathrm{LiH}(s)+\mathrm{H}_{2} \mathrm{O}(l) \longrightarrow$$

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Problem 6

Complete and balance the following equations:
(a) A saltlike (alkaline earth metal) hydride reacting with water,
$$\mathrm{CaH}_{2}(s)+\mathrm{H}_{2} \mathrm{O}(l) \longrightarrow$$
(b) Reduction of a metal halide by hydrogen to form a metal,
$$\mathrm{PdCl}_{2}(a q)+\mathrm{H}_{2}(g) \longrightarrow$$

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Problem 7

Compounds such as $\mathrm{NaBH}_{4},$ Al $\left(\mathrm{BH}_{4}\right)_{3},$ and LiAlH $_{4}$ are complex hydrides used as reducing agents in many syntheses.
(a) Give the oxidation state of each element in these compounds.
(b) Write a Lewis structure for the polyatomic anion in $\mathrm{NaBH}_{4}$ , and predict its shape.

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Problem 8

Unlike the $\mathrm{F}^{-}$ ion, which has an ionic radius close to 133 $\mathrm{pm}$ in all alkali metal fluorides, the ionic radius of $\mathrm{H}^{-}$ varies from 137 $\mathrm{pm}$ in LiH to 152 $\mathrm{pm}$ in CsH. Suggest an explanation for the large variability in the size of $\mathrm{H}^{-}$ but not $\mathrm{F}^{-}$.

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Problem 9

How does the maximum oxidation number vary across a period in the main groups? Is the pattern in Period 2 different?

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Problem 10

What correlation, if any, exists for the Period 2 elements between group number and the number of covalent bonds the element typically forms? How is the correlation different for elements in Periods 3 to 6?

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Problem 11

Each of the chemically active Period 2 elements forms stable compounds in which it has bonds to fluorine.
(a) What are the names and formulas of these compounds?
(b) Does $\Delta \mathrm{EN}$ increase or decrease left to right across the period?
(c) Does percent ionic character increase or decrease left to right?
(d) Draw Lewis structures for these compounds.

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Problem 12

Period 6 is unusual in several ways.
(a) It is the longest period in the table. How many elements belong to Period 6? How many metals?
(b) It contains no metalloids. Where is the metal/nonmetal boundary in Period 6?

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Problem 13

An element forms an oxide, $\mathrm{E}_{2} \mathrm{O}_{3},$ and a fluoride, $\mathrm{EF}_{3}$ .
(a) Of which two groups might $\mathrm{E}$ be a member?
(b) How does the group to which E belongs affect the properties of the oxide and the fluoride?

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Problem 14

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?

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Problem 15

Lithium salts are often much less soluble in water than the corresponding salts of other alkali metals. For example, at $18^{\circ} \mathrm{C},$ the concentration of a saturated LiF solution is $1.0 \times 10^{-2} \mathrm{M}$ , whereas that of a saturated KF solution is 1.6 $\mathrm{M} .$ How would you explain this behavior?

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Problem 16

The alkali metals play virtually the same general chemical role in all their reactions. (a) What is this role? (b) How is it based on atomic properties? (c) Using sodium, write two balanced equations that illustrate this role.

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Problem 17

How do atomic properties account for the low densities of the Group 1A(1) elements?

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Problem 18

Each of the following properties shows a regular trend in Group 1 $\mathrm{A}(1) .$ Predict whether each increases or decreases down the group: (a) density; (b) ionic size; (c) E-E bond energy; (d) $\mathrm{IE}_{1} ;(\mathrm{e})$ magnitude of $\Delta H_{\mathrm{hydr}}$ of $\mathrm{E}^{+}$ ion.

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Problem 19

Each of the following properties shows a regular trend in Group 1 $\mathrm{A}(1) .$ Predict whether each increases or decreases $u p$ the group: (a) melting point; (b) E- $-\mathrm{E}$ bond length; (c) hardness; (d) molar volume; (e) lattice energy of EBr.

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Problem 20

Write a balanced equation for the formation from its elements of sodium peroxide, an industrial bleach.

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Problem 21

Write a balanced equation for the formation of rubidium bromide through a reaction of a strong acid and a strong base.

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Problem 22

Although the alkali metal halides can be prepared directly from the elements, the far less expensive industrial route is treatment of the carbonate or hydroxide with aqueous hydrohalic acid (HX) followed by recrystallization. Balance the reaction between potassium carbonate and aqueous hydriodic acid.

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Problem 23

Lithium forms several useful organolithium compounds. Calculate the mass percent of Li in the following:
(a) Lithium stearate $\left(\mathrm{C}_{17} \mathrm{H}_{35} \text { COOLi), a water-resistant grease used }\right.$ in cars because it does not harden at cold temperatures
(b) Butyllithium (LiC, $_{4} \mathrm{H}_{9} )$ , a reagent in organic syntheses

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Problem 24

How do Groups 1A(1) and 2A(2) compare with respect to reaction of the metals with water?

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Problem 25

Alkaline earth metals are involved in two key diagonal relationships in the periodic table.
(a) Give the two pairs of elements in these diagonal relationships.
(b) For each pair, cite two similarities that demonstrate the relationship.
(c) Why are the members of each pair so similar in behavior?

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Problem 26

The melting points of alkaline earth metals are many times higher than those of the alkali metals. Explain this difference on the basis of atomic properties. Name three other physical properties for which Group 2A(2) metals have higher values than the corresponding 1A(1) metals.

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Problem 27

Write a balanced equation for each reaction:
(a) “Slaking” of lime (treatment with water)
(b) Combustion of calcium in air

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Problem 28

Write a balanced equation for each reaction:
(a) Thermal decomposition of witherite (barium carbonate)
(b) Neutralization of stomach acid ($\mathrm{HCl}$) by milk of magnesia (magnesium hydroxide)

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Problem 29

Lime $(\mathrm{CaO})$ is one of the most abundantly produced chemicals in the world. Write balanced equations for these reactions:
(a) The preparation of lime from natural sources
(b) The use of slaked lime to remove $\mathrm{SO}_{2}$ from flue gases
(c) The reaction of lime with arsenic acid $\left(\mathrm{H}_{3} \mathrm{AsO}_{4}\right)$ to manufacture the insecticide calcium arsenate
(d) The regeneration of $\mathrm{NaOH}$ in the paper industry by reaction of lime with aqueous sodium carbonate

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Problem 30

In some reactions, Be behaves like a typical alkaline earth metal; in others, it does not. Complete and balance the following:
(a) $\mathrm{BeO}(s)+\mathrm{H}_{2} \mathrm{O}(l) \longrightarrow$
(b) $\mathrm{BeCl}_{2}(l)+\mathrm{Cl}^{-}(l ; \text { from molten } \mathrm{NaCl}) \longrightarrow$
In which reaction does Be behave like the other Group 2 $\mathrm{A}(2)$ members?

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Problem 31

How do the transition metals in Period 4 affect the pattern of ionization energies in Group 3A(13)? How does this pattern compare with that in Group 3B(3)?

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Problem 32

How do the acidities of aqueous solutions of $\mathrm{Tl}_{2} \mathrm{O}$ and $\mathrm{TI}_{2} \mathrm{O}_{3}$ compare with each other? Explain.

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Problem 33

Despite the expected decrease in atomic size, there is an unexpected drop in the first ionization energy between Groups 2A(2) and 3A(13) in Periods 2 through 4. Explain this pattern in terms of electron configurations and orbital energies.

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Problem 34

Many compounds of Group 3A(13) elements have chemical behavior that reflects an electron deficiency.
(a) What is the meaning of electron deficiency?
(b) Give two reactions that illustrate this behavior.

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Problem 35

Boron’s chemistry is not typical of its group.
(a) Cite three ways in which boron and its compounds differ significantly from the other 3A(13) members and their compounds.
(b) What is the reason for these differences?

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Problem 36

Rank the following oxides in order of increasing aqueous acidity: $\mathrm{Ga}_{2} \mathrm{O}_{3}, \mathrm{Al}_{2} \mathrm{O}_{3}, \mathrm{In}_{2} \mathrm{O}_{3}$.

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Problem 37

Rank the following hydroxides in order of increasing aqueous basicity: $\mathrm{Al}(\mathrm{OH})_{3}, \mathrm{B}(\mathrm{OH})_{3}, \operatorname{In}(\mathrm{OH})_{3}$.

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Problem 38

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

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Problem 39

Very stable dihalides of the Group 3A(13) metals are known. What is the apparent oxidation state of $\mathrm{Ga}$ in $\mathrm{GaCl}_{2}$? Given that $\mathrm{GaCl}_{2}$ consists of a $\mathrm{Ga}^{+}$ cation and a $\mathrm{GaCl}_{4}^{-}$ anion, what are the actual oxidation states of $\mathrm{Ga}$? Draw the shape of the anion, giving its VSEPR class and bond angles.

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Problem 40

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) Largest temperature range for liquid state of an element
(c) Elemental substance with three-center, two-electron bonds
(d) Metal protected from oxidation by adherent oxide coat
(e) Toxic metal that lies between two other toxic metals

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Problem 41

Indium (In) reacts with $\mathrm{HCl}$ to form a diamagnetic solid with the formula $\mathrm{In} \mathrm{Cl}_{2}$ (a) Write condensed electron configurations for $\mathrm{In}, \mathrm{In}^{+}, \mathrm{In}^{2+},$ and $\mathrm{In}^{3+} .$ (b) Which of these species is (are) diamagnetic and which paramagnetic? (c) What is the apparent oxidation state of In in $\mathrm{In} \mathrm{Cl}_{2}$? (d) Given your answers to parts (b) and (c), explain how $\mathrm{In} \mathrm{Cl}_{2}$2 can be diamagnetic.

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Problem 42

Use VSEPR theory to draw structures, with ideal bond angles, for boric acid and the anion it forms when it reacts with water.

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Problem 43

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 $1000^{\circ} \mathrm{C}$ .
(a) Write a balanced equation for the formation of BN; water also forms.
(b) Calculate $\Delta H_{\mathrm{rxn}}^{\circ}$ for the production of $\mathrm{BN}\left(\Delta H_{\mathrm{f}}^{\circ} \text { of } \mathrm{BN} \text { is }\right.$ $-254 \mathrm{kJ} / \mathrm{mol} )$.
(c) Boron is obtained from the mineral borax, $\mathrm{Na}_{2} \mathrm{B}_{4} \mathrm{O}_{7} \cdot 10 \mathrm{H}_{2} \mathrm{O}$. How much borax is needed to produce 1.0 $\mathrm{kg}$ of $\mathrm{BN}$ , assuming 72$\%$ yield?

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Problem 44

How does the basicity of $\mathrm{SnO}_{2}$ in water compare with that of $\mathrm{CO}_{2}$ ? Explain.

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Problem 45

Nearly every compound of silicon has the element in the $+4$ oxidation state. In contrast, most compounds of lead have the element in the $+2$ 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 3A(13).

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Problem 46

The sum of $\mathrm{IE}_{1}$ through $\mathrm{IE}_{4}$ for Group 4 $\mathrm{A}(14)$ elements shows a decrease from $\mathrm{C}$ to $\mathrm{Si},$ a slight increase from Si to Ge, a decrease from Ge to $\mathrm{Sn},$ and an increase from $\mathrm{Sn}$ to $\mathrm{Pb}$ .
(a) What is the expected trend for IEs down a group?
(b) Suggest a reason for the deviations in Group 4A(14).
(c) Which group might show even greater deviations?

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Problem 47

Give explanations for the large drops in melting point from $\mathrm{C}$ to $\mathrm{Si}$ and from Ge to $\mathrm{Sn}$ .

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Problem 48

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

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Problem 49

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

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Problem 50

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

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Problem 51

Draw a Lewis structure for each species:
(a) The cyclic silicate ion $\mathrm{Si}_{4} \mathrm{O}_{12}^{8-}$
(b) A cyclic hydrocarbon with formula $\mathrm{C}_{4} \mathrm{H}_{8}$

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Problem 52

Draw a Lewis structure for each species:
(a) The cyclic silicate ion $\mathrm{Si}_{6} \mathrm{O}_{18} \mathrm{1} 2-$
(b) A cyclic hydrocarbon with formula $\mathrm{C}_{6} \mathrm{H}_{12}$

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Problem 53

Zeolite $A, \mathrm{Na}_{12}\left[\left(\mathrm{AlO}_{2}\right)_{12}\left(\mathrm{SiO}_{2}\right)_{12}\right] \cdot 27 \mathrm{H}_{2} \mathrm{O},$ is used to soften water by replacing $\mathrm{Ca}^{2+}$ and $\mathrm{Mg}^{2+}$ with $\mathrm{Na}^{+} .$ Hard water from a certain source is $4.5 \times 10^{-3} M \mathrm{Ca}^{2+}$ and $9.2 \times 10^{-4} M \mathrm{Mg}^{2+},$ and a pipe delivers $25,000 \mathrm{L}$ of this hard water per day. What mass ( in kg) of zeolite A is needed to soften a week’s supply of the water? (Assume zeolite A loses its capacity to exchange ions when 85 mol $\%$ of its $\mathrm{Na}^{+}$ has been lost.)

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Problem 54

Give the name and symbol or formula of a Group 4A(14) element or compound that fits each description or use:
(a) Hardest known natural substance
(b) Medicinal antacid
(c) Atmospheric gas implicated in the greenhouse effect
(d) Gas that binds to Fe(II) in blood
(e) Element used in the manufacture of computer chips

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Problem 55

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

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Problem 56

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

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Problem 57

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

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Problem 58

(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?

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Problem 59

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?

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Problem 60

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

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Problem 61

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

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Problem 62

Assuming acid strength relates directly to number of $\mathrm{O}$ atoms bonded to the central atom, rank $\mathrm{H}_{2} \mathrm{N}_{2} \mathrm{O}_{2}\left[\text { or }(\mathrm{HON})_{2}\right] ;$
$\mathrm{HNO}_{3}\left(\text { or } \mathrm{HONO}_{2}\right) ;$ and $\mathrm{HNO}_{2}(\text {or } \mathrm{HONO})$ in order of decreasing acid strength.

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Problem 63

Complete and balance the following:
(a) As $(s)+$ excess $\mathrm{O}_{2}(g) \longrightarrow$
(b) $\mathrm{Bi}(s)+$ excess $\mathrm{F}_{2}(g) \longrightarrow$
(c) $\mathrm{Ca}_{3} \mathrm{As}_{2}(s)+\mathrm{H}_{2} \mathrm{O}(l) \longrightarrow$

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Problem 64

Complete and balance the following:
(a) Excess $\mathrm{Sb}(s)+\mathrm{Br}_{2}(l) \longrightarrow$
(b) $\mathrm{HNO}_{3}(a q)+\mathrm{MgCO}_{3}(s) \longrightarrow$
(c) $\mathrm{K}_{2} \mathrm{HPO}_{4}(s) \stackrel{\Delta}{\longrightarrow}$

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Problem 65

Complete and balance the following:
(a) $\mathrm{N}_{2}(g)+\mathrm{Al}(s) \stackrel{\Delta}{\longrightarrow} \quad$ (b) $\mathrm{PF}_{5}(g)+\mathrm{H}_{2} \mathrm{O}(l) \rightarrow$

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Problem 66

Complete and balance the following:
(a) $\mathrm{AsCl}_{3}(l)+\mathrm{H}_{2} \mathrm{O}(l) \longrightarrow \quad$ (b) $\mathrm{Sb}_{2} \mathrm{O}_{3}(s)+\mathrm{NaOH}(a q) \rightarrow$

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Problem 67

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

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Problem 68

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

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Problem 69

The pentafluorides of the larger members of Group 5 $\mathrm{A}(15)$ have been prepared, but $\mathrm{N}$ can have only eight electrons. $\mathrm{A}$ claim has been made that, at low temperatures, a compound with the empirical formula $\mathrm{NF}_{5}$ forms. Draw a possible Lewis structure for this compound. (Hint: $\mathrm{NF}_{5}$ is ionic.)

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Problem 70

Give the name and symbol or formula of a Group 5A(15) element or compound that fits each description or use:
(a) Hydride that exhibits hydrogen bonding
(b) Compound used in “strike-anywhere” match heads
(c) Oxide used as a laboratory drying agent
(d) Odd-electron molecule (two examples)
(e) Compound used as an additive in soft drinks

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Problem 71

In addition to those in Table 14.3, other less stable nitrogen oxides exist. Draw a Lewis structure for each of the following:
(a) $\mathrm{N}_{2} \mathrm{O}_{2},$ a dimer of nitrogen monoxide with an $\mathrm{N}-\mathrm{N}$ bond
(b) $\mathrm{N}_{2} \mathrm{O}_{2},$ a dimer of nitrogen monoxide with no $\mathrm{N}-\mathrm{N}$ bond
(c) $\mathrm{N}_{2} \mathrm{O}_{3}$ with no $\mathrm{N}-\mathrm{N}$ bond
(d) $\mathrm{NO}^{+}$ and $\mathrm{NO}_{3}^{-},$ products of the ionization of liquid $\mathrm{N}_{2} \mathrm{O}_{4}$

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Problem 72

Nitrous oxide $\left(\mathrm{N}_{2} \mathrm{O}\right),$ the "laughing gas" used as an anesthetic by dentists, is made by thermal decomposition of solid $\mathrm{NH}_{4} \mathrm{NO}_{3}$ . Write a balanced equation for this reaction. What are the oxidation states of $\mathrm{N}$ in $\mathrm{NH}_{4} \mathrm{NO}_{3}$ and in $\mathrm{N}_{2} \mathrm{O} ?$

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Problem 73

Write balanced equations for the thermal decomposition of potassium nitrate ($\mathrm{O}_{2}$ is also formed in both cases): (a) at low temperature to the nitrite; (b) at high temperature to the metal oxide and nitrogen.

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Problem 74

Rank the following in order of increasing electrical conductivity, and explain your ranking: Po, S, Se.

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Problem 75

The oxygen and nitrogen families have some obvious similarities and differences.
(a) State two general physical similarities between Group 5A(15) and 6A(16) elements.
(b) State two general chemical similarities between Group 5A(15) and 6A(16) elements.
(c) State two chemical similarities between P and S.
(d) State two physical similarities between N and O.
(e) State two chemical differences between N and O.

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Problem 76

A molecular property of the Group 6A(16) hydrides changes abruptly down the group. This change has been explained in terms of a change in orbital hybridization.
(a) Between what periods does the change occur?
(b) What is the change in the molecular property?
(c) What is the change in hybridization?
(d) What other group displays a similar change?

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Problem 77

Complete and balance the following:
(a) $\mathrm{NaHSO}_{4}(a q)+\mathrm{NaOH}(a q) \longrightarrow$
(b) $\mathrm{S}_{8}(s)+$ excess $\mathrm{F}_{2}(g) \longrightarrow$
(c) $\mathrm{FeS}(s)+\mathrm{HCl}(a q) \rightarrow$
(d) $\mathrm{TeS}(s)+\mathrm{I}_{2}(s) \longrightarrow$

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Problem 78

Complete and balance the following:
(a) $\mathrm{H}_{2} \mathrm{S}(g)+\mathrm{O}_{2}(g) \longrightarrow$
(b) $\mathrm{SO}_{3}(g)+\mathrm{H}_{2} \mathrm{O}(l) \longrightarrow$
(c) $\mathrm{SF}_{4}(g)+\mathrm{H}_{2} \mathrm{O}(l) \longrightarrow$
(d) $\mathrm{Al}_{2} \mathrm{Se}_{3}(s)+\mathrm{H}_{2} \mathrm{O}(l) \longrightarrow$

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Problem 79

Is each oxide basic, acidic, or amphoteric in water: (a) $\mathrm{SeO}_{2}$ (b) $\mathrm{N}_{2} \mathrm{O}_{3} ;(\mathrm{c}) \mathrm{K}_{2} \mathrm{O} ;(\mathrm{d}) \mathrm{BeO} ;(\mathrm{e}) \mathrm{BaO}$ ?

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Problem 80

Is each oxide basic, acidic, or amphoteric in water: (a) $\mathrm{MgO}$ ; (b) $\mathrm{N}_{2} \mathrm{O}_{5} ;(\mathrm{c}) \mathrm{CaO} ;(\mathrm{d}) \mathrm{CO}_{2} ;(\mathrm{e}) \mathrm{TeO}_{2} ?$

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Problem 81

Rank the following hydrides in order of increasing acid strength: $\mathrm{H}_{2} \mathrm{S}, \mathrm{H}_{2} \mathrm{O}, \mathrm{H}_{2} \mathrm{Te} .$

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Problem 82

Rank the following species in order of decreasing acid strength: $\mathrm{H}_{2} \mathrm{SO}_{4}, \mathrm{H}_{2} \mathrm{SO}_{3}, \mathrm{HSO}_{3}^{-}$.

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Problem 83

Selenium tetrafluoride reacts with more fluorine to form selenium hexafluoride.
(a) Draw Lewis structures for both selenium fluorides and predict any deviations from ideal bond angles.
(b) Describe the change in orbital hybridization of the central Se during the reaction.

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Problem 84

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

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Problem 85

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

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Problem 86

Disulfur decafluoride is intermediate in reactivity between $\mathrm{SF}_{4}$ and $\mathrm{SF}_{6}$ . It disproportionates at $150^{\circ} \mathrm{C}$ to these monosulfur fluorides. Write a balanced equation for this reaction, and give the oxidation state of $\mathrm{S}$ in each compound.

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Problem 87

(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.

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Problem 88

(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?

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Problem 89

How many electrons does a halogen atom need to complete its octet? Give examples of the ways a Cl atom can do so.

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Problem 90

Select the stronger bond in each pair:
(a) $\mathrm{Cl}-\mathrm{Cl}$ or $\mathrm{Br}-\mathrm{Br} \quad$ (b) $\mathrm{Br}-$ Br or $\mathrm{I}-\mathrm{I}$
(c) $\mathrm{F}-\mathrm{F}$ or $\mathrm{Cl}-\mathrm{Cl}$ . Why doesn't the $\mathrm{F}-\mathrm{F}$ bond strength follow the group trend?

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Problem 91

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.

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Problem 92

(a) A halogen $\left(\mathrm{X}_{2}\right)$ disproportionates in base in several steps to $\mathrm{X}^{-}$ and $\mathrm{XO}_{3}$ . Write the overall equation for the disproportionation of $\mathrm{Br}_{2}$ in excess $\mathrm{OH}^{-}$ to $\mathrm{Br}^{-}$ and $\mathrm{BrO}_{3}^{-}$
(b) Write a balanced equation for the reaction of $\mathrm{CIF}_{5}$ with aqueous base (Hint: Add base to the reaction of $\mathrm{BrF}_{5}$ on p. 604 ).

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Problem 93

Complete and balance the following equations. If no reaction occurs, write NR:
(a) $\operatorname{Rb}(s)+\mathrm{Br}_{2}(l) \longrightarrow \quad$ (b) $\mathrm{I}_{2}(s)+\mathrm{H}_{2} \mathrm{O}(l) \longrightarrow$
(c) $\mathrm{Br}_{2}(l)+\mathrm{I}^{-}(a q) \longrightarrow \quad$ (d) $\mathrm{CaF}_{2}(s)+\mathrm{H}_{2} \mathrm{SO}_{4}(l) \longrightarrow$

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Problem 94

Complete and balance the following equations. If no reaction occurs, write NR:
(a) $\mathrm{H}_{3} \mathrm{PO}_{4}(l)+\mathrm{NaI}(s) \longrightarrow \quad$ (b) $\mathrm{Cl}_{2}(g)+\mathrm{I}^{-}(a q) \longrightarrow$
(c) $\mathrm{Br}_{2}(l)+\mathrm{Cl}^{-}(a q) \longrightarrow \quad$ (d) $\mathrm{ClF}(g)+\mathrm{F}_{2}(g) \rightarrow$

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Problem 95

Rank the following acids in order of increasing acid strength: $\mathrm{HClO}, \mathrm{HClO}_{2}, \mathrm{HClO}_{2}, \mathrm{HBrO}, \mathrm{HIO}$ .

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Problem 96

Rank the following acids in order of decreasing acid strength: $\mathrm{HBrO}_{3}, \mathrm{HBrO}_{4}, \mathrm{HIO}_{3}, \mathrm{HClO}_{4}$ .

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Problem 97

Give the name and symbol or formula of a Group 7A(17) element or compound that fits each description or use:
(a) Used in etching glass
(b) Compound used in household bleach
(c) Weakest hydrohalic acid
(d) Element that is a liquid at room temperature
(e) Organic chloride used to make PVC

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Problem 98

An industrial chemist treats solid $\mathrm{NaCl}$ with concentrated $\mathrm{H}_{2} \mathrm{SO}_{4}$ and obtains gaseous $\mathrm{HCl}$ and $\mathrm{NaHSO}_{4}$ . When she substitutes solid Nal for NaCl, gaseous $\mathrm{H}_{2} \mathrm{S}$ , solid $\mathrm{I}_{2},$ and $\mathrm{S}_{8}$ are obtained but no HI. (a) What type of reaction did the $\mathrm{H}_{2} \mathrm{SO}_{4}$ undergo with Nal? (b) Why does Nal, but not NaCl, cause this type of reaction? (c) To produce HI(g) by the reaction of Nal with an acid, how does the acid have to differ from sulfuric acid?

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Problem 99

Rank the halogens $\mathrm{Cl}_{2}, \mathrm{Br}_{2},$ and $\mathrm{I}_{2}$ in order of increasing oxidizing strength based on their products with metallic Re: $\mathrm{ReCl}_{6}$, $\mathrm{ReBr}_{5},\mathrm{Rel}_{4}$ . Explain your ranking.

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Problem 100

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

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Problem 101

What oxidation states does Xe show in its compounds?

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Problem 102

Why do the noble gases have such low boiling points?

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Problem 103

Explain why Xe, and to a limited extent Kr, form compounds, whereas He, Ne, and Ar do not.

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Problem 104

(a) Why do stable xenon fluorides have an even number of Fatoms? (b) Why do the ionic species $\mathrm{XeF}_{3}+$ and $\mathrm{XeF}_{7}$ - have odd numbers of $\mathrm{F}$ atoms? (c) Predict the shape of $\mathrm{XeF}_{3}^{+}$ .

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Problem 105

Xenon tetrafluoride reacts with antimony pentafluoride to form the ionic complex [XeF $_{3} ]^{+}\left[\operatorname{SbF}_{6}\right]^{-} .$ (a) Which depiction shows the molecular shapes of the reactants and product? (b) How, if at all, does the hybridization of xenon change in the reaction?

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Problem 106

Given the following information,
$$\begin{array}{ll}{\mathrm{H}^{+}(g)+\mathrm{H}_{2} \mathrm{O}(g) \longrightarrow \mathrm{H}_{3} \mathrm{O}^{+}(g)} & {\Delta H=-720 \mathrm{kJ}} \\ {\mathrm{H}^{+}(g)+\mathrm{H}_{2} \mathrm{O}(l) \longrightarrow \mathrm{H}_{3} \mathrm{O}^{+}(a q)} & {\Delta H=-1090 \mathrm{kJ}}\end{array}$$
$$\quad \mathrm{H}_{2} \mathrm{O}(l) \rightarrow \mathrm{H}_{2} \mathrm{O}(g) \quad \Delta H=40.7 \mathrm{kJ}$$
calculate the heat of solution of the hydronium ion:
$$\mathrm{H}_{3} \mathrm{O}^{+}(g) \stackrel{\mathrm{H}_{2} \mathrm{O}}{\longrightarrow} \mathrm{H}_{3} \mathrm{O}^{+}(a q)$$

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Problem 107

The electronic transition in Na from 3$p^{1}$ to 3$s^{1}$ gives rise to a bright yellow-orange emission at 589.2 nm. What is the energy of this transition?

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Problem 108

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

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Problem 109

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 $2.50 \times 10^{-3}$ mol, what mass of each product forms?

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Problem 110

The main reason alkali metal dihalides $\left(\mathrm{MX}_{2}\right)$ do not form is the high $\mathrm{IE}_{2}$ of the metal. (a) Why is $\mathrm{IE}_{2}$ so high for alkali metals? (b) The IE, for Cs is $2255 \mathrm{kJ} / \mathrm{mol},$ low enough for $\mathrm{CsF}_{2}$ to form exothermically $\left(\Delta H_{\mathrm{f}}^{\circ}=-125 \mathrm{kJ} / \mathrm{mol}\right) .$ This compound cannot be synthesized, however, because $\mathrm{CsF}$ forms with a much greater release of heat $\left(\Delta H_{\mathrm{f}}^{\circ}=-530 \mathrm{kJ} / \mathrm{mol}\right) .$ Thus, the breakdown of $\mathrm{CsF}_{2}$ to CsF happens readily. Write the equation for this breakdown, and calculate the enthalpy of reaction per mole of $\mathrm{CsF}$ .

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Problem 111

Semiconductors made from elements in Groups 3A(13) and 5A(15) are typically prepared by direct reaction of the elements at high temperature. An engineer treats 32.5 g of molten gallium with 20.4 L of white phosphorus vapor at 515 K and 195 kPa. If purification losses are 7.2% by mass, how many grams of gallium phosphide will be prepared?

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Problem 112

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

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Problem 113

The species $\mathrm{CO}, \mathrm{CN}^{-},$ and $\mathrm{C}_{2}^{2-}$ are isoelectronic.
(a) Draw their Lewis structures.
(b) Draw their MO diagrams (assume $2 s-2 p$ mixing, as in $\mathrm{N}_{2}$ ), and give the bond order and electron configuration for each.

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Problem 114

The Ostwald process is a series of three reactions used for the industrial production of nitric acid from ammonia.
(a) Write a series of balanced equations for the Ostwald process.
(b) If NO is not recycled, how many moles of $\mathrm{NH}_{3}$ are consumed per mole of $\mathrm{HNO}_{3}$ produced?
(c) In a typical industrial unit, the process is very efficient, with a 96$\%$ yield for the first step. Assuming 100$\%$ yields for the subsequent steps, what volume of concentrated aqueous nitric acid $(60 . \% \text { by mass; } d=1.37 \mathrm{g} / \mathrm{mL})$ can be prepared for each cubic meter of a gas mixture that is $90 . \%$ air and $10 . \% \mathrm{NH}_{3}$ by volume at the industrial conditions of 5.0 atm and $850 .^{\circ} \mathrm{C} ?$

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Problem 115

All common plant fertilizers contain nitrogen compounds. Determine the mass % of N in (a) ammonia; (b) ammonium nitrate; (c) ammonium hydrogen phosphate.

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Problem 116

Producer gas is a fuel formed by passing air over red-hot coke (amorphous carbon). What mass of a producer gas that consists of 25% CO, 5.0% CO2, and 70.% N2 by mass can be formed from 1.75 metric tons of coke, assuming an 87% yield?

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Problem 117

Gaseous $\mathrm{F}_{2}$ reacts with water to form HF and $\mathrm{O}_{2}$ . In $\mathrm{NaOH}$ solution, $\mathrm{F}_{2}$ forms $\mathrm{F}^{-},$ water, and oxygen difluoride $\left(\mathrm{OF}_{2}\right), \mathrm{a}$ highly toxic gas and powerful oxidizing agent. The $\mathrm{OF}_{2}$ reacts with excess $\mathrm{OH}^{-}$ , forming $\mathrm{O}_{2},$ water, and $\mathrm{F}^{-}$
(a) For each reaction, write a balanced equation, give the oxidation state of $\mathrm{O}$ in all compounds, and identify the oxidizing and reducing agents.
(b) Draw a Lewis structure for $\mathrm{OF}_{2},$ and predict its shape.

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Problem 118

What is a disproportionation reaction, and which of the following fit the description?
(a) $\mathrm{I}_{2}(s)+\mathrm{KI}(a q) \longrightarrow \mathrm{KI}_{3}(a q)$
(b) $2 \mathrm{ClO}_{2}(g)+\mathrm{H}_{2} \mathrm{O}(l) \longrightarrow \mathrm{HClO}_{3}(a q)+\mathrm{HClO}_{2}(a q)$
(c) $\mathrm{Cl}_{2}(g)+2 \mathrm{NaOH}(a q) \longrightarrow$
NaCl $(a q)+\mathrm{NaClO}(a q)+\mathrm{H}_{2} \mathrm{O}(l)$
(d) $\mathrm{NH}_{4} \mathrm{NO}_{2}(s) \longrightarrow \mathrm{N}_{2}(g)+2 \mathrm{H}_{2} \mathrm{O}(g)$
(e) $3 \mathrm{MnO}_{4}^{2-}(a q)+2 \mathrm{H}_{2} \mathrm{O}(l) \longrightarrow$
$2 \mathrm{MnO}_{4}-(a q)+\mathrm{MnO}_{2}(s)+4 \mathrm{OH}^{-}(a q)$
(f) $3 \mathrm{AuCl}(s) \longrightarrow \operatorname{AuCl}_{3}(s)+2 \mathrm{Au}(s)$

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Problem 119

Explain the following observations:
(a) In reactions with $\mathrm{Cl}_{2},$ phosphorus forms $\mathrm{PCl}_{5}$ in addition to the expected $\mathrm{PCl}_{3},$ but nitrogen forms only $\mathrm{NCl}_{3}$ .
(b) Carbon tetrachloride is unreactive toward water, but silicon tetrachloride reacts rapidly and completely. (To give what?)
(c) The sulfur-oxygen bond in $\mathrm{SO}_{4}^{2-}$ is shorter than expected for an $\mathrm{S}-$ O single bond.
(d) Chlorine forms $\mathrm{ClF}_{3}$ and $\mathrm{ClF}_{5},$ but $\mathrm{ClF}_{4}$ is unknown.

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Problem 120

Which group(s) of the periodic table is (are) described by each of the following general statements?
(a) The elements form compounds of VSEPR class $\mathrm{AX}_{3} \mathrm{E}$.
(b) The free elements are strong oxidizing agents and form monatomic ions and oxoanions.
(c) The atoms form compounds by combining with two other atoms that donate one electron each.
(d) The free elements are strong reducing agents, show only one nonzero oxidation state, and form mainly ionic compounds.
(e) The elements can form stable compounds with only three bonds, but as a central atom, they can accept a pair of electrons from a fourth atom without expanding their valence shell.
(f) Only larger members of the group are chemically active.

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Problem 121

Diiodine pentoxide (I $_{2} \mathrm{O}_{5} )$ was discovered by Joseph Gay-Lussac in 1813 , but its structure was unknown until $1970 !$ Like $\mathrm{Cl}_{2} \mathrm{O}_{7},$ it can be prepared by the dehydration-condensation of the corresponding oxoacid.
(a) Name the precursor oxoacid, write a reaction for formation of the oxide, and draw a likely Lewis structure.
(b) Data show that the bonds to the terminal O are shorter than the bonds to the bridging O. Why?
(c) $\mathrm{I}_{2} \mathrm{O}_{5}$ is one of the few chemicals that can oxidize CO rapidly and completely; elemental iodine forms in the process. Write a balanced equation for this reaction.

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Problem 122

Bromine monofluoride (BrF) disproportionates to bromine gas and bromine tri- and pentafluorides. Use the following to find $\Delta H_{\mathrm{rxn}}^{\mathrm{o}}$ for the decomposition of BrF to its elements:
$$3 \mathrm{BrF}(g) \rightarrow \mathrm{Br}_{2}(g)+\mathrm{BrF}_{3}(l) \quad \Delta H_{\mathrm{ran}}=-125.3 \mathrm{kJ}$$
$$5 \mathrm{BrF}(g) \rightarrow 2\mathrm{Br}_{2}(g)+\mathrm{BrF}_{5}(l) \quad \Delta H_{\mathrm{ran}}=-1661 \mathrm{kJ}$$
$$\mathrm{Br} \mathrm{F}_{3}(l)+\mathrm{F}_{2}(g) \longrightarrow \operatorname{Br} \mathrm{F}_{5}(l) \qquad \Delta H_{\mathrm{rxn}}=-158.0 \mathrm{kJ}$$

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Problem 123

White phosphorus is prepared by heating phosphate rock [principally $\mathrm{Ca}_{3}\left(\mathrm{PO}_{4}\right)_{2} ]$ with sand and coke:
$\mathrm{Ca}_{3}\left(\mathrm{PO}_{4}\right)_{2}(s)+\mathrm{SiO}_{2}(s)+\mathrm{C}(s) \longrightarrow$
$$\mathrm{CaSiO}_{3}(s)+\mathrm{CO}(g)+\mathrm{P}_{4}(g)[\text { unbalanced }]$$
How many kilograms of phosphate rock are needed to produce 315 mol of $\mathrm{P}_{4},$ assuming that the conversion is $90 . \%$ efficient?

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Problem 124

Element E forms an oxide of general structure $A$ and a chloride of general structure $B,$ shown at right.For the anion $\mathrm{EF}_{5}^{-},$ what is (a) the molecular shape; (b) the hybridization of $\mathrm{E} ;(\mathrm{c})$ the O.N. of $\mathrm{E}$ ?

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Problem 125

From its formula, one might expect CO to be quite polar, but its dipole moment is actually 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.

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Problem 126

When an alkaline earth carbonate is heated, it releases $\mathrm{CO}_{2},$ leaving the metal oxide. The temperature at which each Group 2 $\mathrm{A}(2)$ carbonate yields a $\mathrm{CO}_{2}$ partial pressure of 1 $\mathrm{atm}$ is
$$\begin{array}{ll}{\text { Carbonate }} & {\text { Temperature }\left(^{\circ} \mathrm{C}\right)} \\ \hline \quad {\mathrm{MgCO}_{3}} & \quad {542} \\ \quad {\mathrm{CaCO}_{3}} & \quad {882} \\ \quad {\mathrm{SrCO}_{3}} & \quad {1155} \\ \quad {\mathrm{BACO}_{3}} & \quad {1360} \end{array}$$
(a) Suggest a reason for this trend.
(b) Mixtures of $\mathrm{CaCO}_{3}$ and MgO are used to absorb dissolved silicates from boiler water. How would you prepare a mixture of $\mathrm{CaCO}_{3}$ and MgO from dolomite, which contains $\mathrm{CaCO}_{3}$ and $\mathrm{MgCO}_{3} ?$

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Problem 127

The bond angles in the nitrite ion, nitrogen dioxide, and the nitronium ion $\left(\mathrm{NO}_{2}^{+}\right)$ are $115^{\circ}, 134^{\circ},$ and $180^{\circ},$ respectively. Explain these values using Lewis structures and VSEPR theory.

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Problem 128

A common method for producing a gaseous hydride is to treat a salt containing the anion of the volatile hydride with a strong acid.
(a) Write an equation for each of the following examples:
(1) the production of HF from $\mathrm{CaF}_{2} ;(2)$ the production of $\mathrm{HCl}$ from $\mathrm{NaCl} ;(3)$ the production of $\mathrm{H}_{2}$ S from FeS.
(b) In some cases, even as weak an acid as water can be used for this preparation if the anion of the salt has a sufficiently strong attraction for protons. An example is the production of $\mathrm{PH}_{3}$ from $\mathrm{Ca}_{3} \mathrm{P}_{2}$ and water. Write the equation for this reaction.
(c) By analogy, predict the products and write the equation for the reaction of $\mathrm{Al}_{4} \mathrm{C}_{3}$ with water.

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Problem 129

Chlorine trifluoride was formerly used in the production of uranium hexafluoride for the U.S. nuclear industry:
$$\mathrm{U}(s)+3 \mathrm{ClF}_{3}(l) \longrightarrow \mathrm{UF}_{6}(l)+3 \mathrm{ClF}(g)$$
How many grams of $\mathrm{UF}_{6}$ can form from 1.00 metric ton of uranium ore that is 1.55% by mass uranium and 12.75 L of chlorine trifluoride $(d=1.88 \mathrm{g} / \mathrm{mL}) ?$

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Problem 130

Chlorine is used to make bleach solutions containing5.25$\% \mathrm{NaClO}$ (by mass). Assuming 100$\%$ yield in the reaction producing NaClO from $\mathrm{Cl}_{2},$ how many liters of $\mathrm{Cl}_{2}(g)$ at STP will be needed to make $1000 .$ L of bleach solution $(d=1.07 \mathrm{g} / \mathrm{mL}) ?$

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Problem 131

The triatomic molecular ion $\mathrm{H}_{3}^{+}$ was first detected and characterized by $\mathrm{J}$ . J. Thomson using mass spectrometry. Use the bond energy of $\mathrm{H}_{2}(432 \mathrm{kJ} / \mathrm{mol})$ and the proton affinity of $\mathrm{H}_{2}$ $\left(\mathrm{H}_{2}+\mathrm{H}^{+} \longrightarrow \mathrm{H}_{3}^{+} ; \Delta H=-337 \mathrm{kJ} / \mathrm{mol}\right)$ to calculate the enthalpy of reaction for $\mathrm{H}+\mathrm{H}+\mathrm{H}^{+} \longrightarrow \mathrm{H}_{3}^{+}$

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Problem 132

An atomic hydrogen torch is used for cutting and welding thick sheets of metal. When $\mathrm{H}_{2}$ passes through an electric arc, the molecules decompose into atoms, which react with $\mathrm{O}_{2}$ Temperatures over $5000^{\circ} \mathrm{C}$ are reached, which can melt all metals. Write equations for the breakdown of $\mathrm{H}_{2}$ to $\mathrm{H}$ atoms and for the subsequent overall reaction of the $\mathrm{H}$ atoms with oxygen. Use Appendix $\mathrm{B}$ to find the standard enthalpy of each reaction per mole of product.

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Problem 133

Which of the following oxygen ions are paramagnetic: $\mathrm{O}^{+}$ , $\mathrm{O}^{-}, \mathrm{O}^{2-}, \mathrm{O}^{2+} ?$

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Problem 134

Copper(II) hydrogen arsenite $\left(\mathrm{CuHAsO}_{3}\right)$ is a green pigment once used in wallpaper. In damp conditions, mold metabolizes this compound to trimethylarsine $\left[\left(\mathrm{CH}_{3}\right)_{3} \mathrm{As}\right]$, a highly toxic gas.
(a) Calculate the mass percent of As in each compound.
(b) How much CuHAsO, must react to reach a toxic level in a room that measures 12.35 $\mathrm{m} \times 7.52 \mathrm{m} \times 2.98 \mathrm{m}$ (arsenic is toxic at 0.50 $\mathrm{mg} / \mathrm{m}^{3} ) ?$

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Problem 135

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

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