Chapter 10

In carbon dioxide, \(\mathrm{CO}_{2}\) (a) Assign an oxidation state to each atom in the molecule. (Hint: Draw a dot diagram first.) (b) How many electrons does the \(C\) atom own by oxidation-state electron bookkeeping? (c) How many more or fewer valence electrons are assigned to the \(\mathrm{C}\) atom here than are present in a free \(C\) atom? (d) Is it correct or incorrect to say that the \(\mathrm{C}\) atom in \(\mathrm{CO}_{2}\) has a complete octet of valence electrons? Explain.

In methane, \(\mathrm{CH}_{4}\), (a) Assign an oxidation state to each atom in the molecule. (Hint: Draw a dot diagram first.) (b) How many electrons does the \(C\) atom own by oxidation-state electron bookkeeping? (c) How many more or fewer valence electrons are assigned to the \(\mathrm{C}\) atom here than are present in a free \(C\) atom?

In chloroform, \(\mathrm{CHCl}_{3}\) (a) Assign an oxidation state to each atom in the molecule. (Hint: Draw a dot diagram first.) (b) How many electrons does the \(\mathrm{C}\) atom own by oxidation-state electron bookkeeping? (c) How many more or fewer valence electrons are assigned to the \(\mathrm{C}\) atom here than are present in a free \(C\) atom? (d) Is it correct or incorrect to say that the \(\mathrm{C}\) atom in \(\mathrm{CHCl}_{3}\) has a complete octet of valence electrons? Explain.

Use the shortcut rules to assign oxidation states to all atoms. \(\mathrm{COCl}_{2}\) (oxygen and chlorine bonded to the central carbon) Answer: \(\mathrm{O}\) is \(-2\) (Rule 2). \(\mathrm{Cl}\) is \(-1\) (halide rule applies because \(\mathrm{Cl}\) is more electronegative than C). Therefore C is \(+4\) (the sum of the oxidation numbers for the \(\mathrm{O}\) and two \(\mathrm{Cl}^{\prime} \mathrm{s}\) is \(-4)\)

Use the shortcut rules to assign oxidation states to all atoms. \(\mathrm{MgBr}_{2}\)

Use the shortcut rules to assign oxidation states to all atoms. \(\mathrm{Fe}_{2} \mathrm{O}_{3}\)

Use the shortcut rules to assign oxidation states to all atoms. \(\mathrm{CO}_{3}^{2-}\)

Use the shortcut rules to assign oxidation states to all atoms. \(\mathrm{Cu}^{2+}\)

Indicate whether each reaction is a redox reaction. If it is, which atom gets oxidized and which atom gets reduced? Consult the shortcut rules. \(\mathrm{P}_{4}+6 \mathrm{Br}_{2} \rightarrow 4 \mathrm{PBr}_{3}\) (Hint: \(\mathrm{Br}\) is more electronegative than P.)

Indicate whether each reaction is a redox reaction. If it is, which atom gets oxidized and which atom gets reduced? Consult the shortcut rules. \(4 \mathrm{Fe}+3 \mathrm{O}_{2} \rightarrow 2 \mathrm{Fe}_{2} \mathrm{O}_{3}\)

Indicate whether each reaction is a redox reaction. If it is, which atom gets oxidized and which atom gets reduced? Consult the shortcut rules. \(\mathrm{Ca}+2 \mathrm{H}^{+} \rightarrow \mathrm{Ca}^{2+}+\mathrm{H}_{2}\)

Indicate whether each reaction is a redox reaction. If it is, which atom gets oxidized and which atom gets reduced? Consult the shortcut rules. \(2 \mathrm{NaBr}+\mathrm{MgO} \rightarrow \mathrm{MgBr}_{2}+\mathrm{Na}_{2} \mathrm{O}\)

For each redox reaction, indicate which substance is the oxidizing agent and which is the reducing agent. \(\mathrm{P}_{4}+6 \mathrm{Br}_{2} \rightarrow 4 \mathrm{PBr}_{3}\)

For each redox reaction, indicate which substance is the oxidizing agent and which is the reducing agent. \(4 \mathrm{Fe}+3 \mathrm{O}_{2} \rightarrow 2 \mathrm{Fe}_{2} \mathrm{O}_{3}\)

\(\mathrm{Ca}+2 \mathrm{H}^{+} \rightarrow \mathrm{Ca}^{2+}+\mathrm{H}_{2}\)

A battery is constructed from iron, Fe, and silver, \(\mathrm{Ag}\), by dipping a strip of each metal into a solution of its ions \(\left(\mathrm{Fe}^{3+}\right.\) and \(\mathrm{Ag}^{+}\), respectively). As the battery operates, the \(\mathrm{Fe}^{3+}\) concentration increases while the \(\mathrm{Ag}^{+}\) concentration decreases. (a) What is getting oxidized? (b) What is getting reduced? (c) Draw a battery similar to the one you drew for WorkPatch \(10.7\), label which way the electrons flow, and label cathode, \(+\), and anode, \(=\).

What will happen if a piece of Zn metal is placed in a solution of \(\mathrm{Al}^{3+}\) ions?

Which of the following reactions is spontaneous? (a) \(3 \mathrm{Zn}^{2+}+2 \mathrm{Al} \rightarrow 3 \mathrm{Zn}+2 \mathrm{Al}^{3+}\) (b) \(3 \mathrm{Zn}+2 \mathrm{Al}^{3+} \rightarrow 3 \mathrm{Zn}^{2+}+2 \mathrm{Al}\)

Consider a battery made from lead, \(\mathrm{Pb}\), and copper, \(\mathrm{Cu}\), along with \(\mathrm{Pb}^{2+}\) ions and \(\mathrm{Cu}^{2+}\) ions. In such a battery, what is the oxidizing agent, and what is the reducing agent? (Hint: Start by consulting the EMF series, and determine which metal oxidizes and which metal ion gets reduced.)

How is electricity similar to water flowing in hose?

How is electricity mechanically produced in a power plant?

What is an electron-transfer reaction?

Why did chemists find it necessary to invent oxidation states?

Which method of assigning shared electrons is correct-the double-counting method of the octet rule or the oxidation-state method?

When assigning an oxidation state to an atom, why do we need to know how many valence electrons are present on a free atom of that element?

Draw a Lewis dot diagram for \(\mathrm{NF}_{2} \mathrm{H}\), and use the oxidation-state method of electron bookkeeping to determine how many electrons each atom should be assigned.

Draw a Lewis dot diagram for \(\mathrm{H}_{2} \mathrm{O}_{2}\) (hydrogen peroxide), and use the oxidation-state method of electron bookkeeping to determine how many electrons each atom should be assigned.

Draw a Lewis dot diagram for \(\mathrm{NO}_{3}^{-}\), and use the oxidation- state method of electron bookkeeping to determine how many electrons each atom should be assigned.

Consider the following molecules. Without using the shortcut rules, determine the oxidation state for each atom in the molecules and show how you calculated it. (a) \(\mathrm{H}_{2}\) (b) \(\mathrm{O}_{2}\) (c) \(\mathrm{Cl}_{2}\) (d) IF

What must always be true when you add up all the oxidation states for the atoms in a molecule?

Suppose an atom has an oxidation state of \(+3\). (a) Would more or fewer electrons be assigned to this atom by oxidation-state bookkeeping than are present on a free atom of that element? (b) How many more or fewer?

Suppose the shortcut rules can determine the oxidation state of every atom in a compound except one. How can you find the oxidation state of the remaining atom?

Use the shortcut rules to assign an oxidation state to each atom in: (a) \(\mathrm{PCl}_{3}\) (b) \(\mathrm{H}_{2} \mathrm{~S}\) (c) \(\mathrm{MnO}_{4}^{-}\) (d) \(\mathrm{HNO}_{3}\) (e) \(\mathrm{HCOOH}\) (f) \(\mathrm{S}_{2} \mathrm{O}_{3}^{2-}\)

Use the shortcut rules to assign an oxidation state to each atom in: (a) \(\mathrm{CO}\) (b) \(\mathrm{CH}_{2} \mathrm{Cl}_{2}\) (c) \(\mathrm{HCOO}^{-}\) (d) \(\mathrm{PtCl}_{6}^{2-}\)

Use the shortcut rules to assign an oxidation state to each atom in: (a) \(\mathrm{O}^{2-}\) (b) \(\mathrm{Li}_{3} \mathrm{~N}\) (c) \(\mathrm{MgSO}_{4}\) (d) \(\mathrm{MnO}_{2}\)

Consider \(\mathrm{ClO}^{-}\) and \(\mathrm{AlCl}_{3}\). For one of these substances, the halide shortcut rule works. For the other, it does not. (a) Which one does it work for, and why? (b) Why doesn't it work for the other? (c) Assign oxidation states to all the atoms in both substances.

WorkPatch \(10.2\) had you assign oxidation states to the oxygen atoms in ozone. The answer was \(-1\) and 0 for the terminal (end) oxygen atoms and \(+1\) for the central oxygen atom. But now, a student claims the oxidation states are \(+1\) for the central oxygen atom and \(-1 / 2\) for the terminal oxygen atoms. He is also right. Explain why.

What do the terms oxidation and reduction mean with regard to valence electrons?

When is an oxidizing agent also the species that gets oxidized? Explain.

How are oxidation states useful in determining whether a reaction is a redox reaction?

What happens to an atom's oxidation state when the atom is reduced?

What happens to an atom's oxidation state when the atom is oxidized?

What does the word transfer imply about an electron-transfer reaction?

Why can we always call an electron-transfer reaction a redox reaction?

Which of the following are redox reactions? (a) \(2 \mathrm{Na}+2 \mathrm{H}_{2} \mathrm{O} \rightarrow 2 \mathrm{NaOH}+\mathrm{H}_{2}\) (b) \(\mathrm{MgBr}_{2}+2 \mathrm{NaF} \rightarrow \mathrm{MgF}_{2}+2 \mathrm{NaBr}\) (c) \(2 \mathrm{CO}+\mathrm{O}_{2} \rightarrow 2 \mathrm{CO}_{2}\) (d) \(\mathrm{SO}_{2}+\mathrm{H}_{2} \mathrm{O} \rightarrow \mathrm{H}_{2} \mathrm{SO}_{3}\)

Which of the following are electron-transfer reactions? (a) \(2 \mathrm{CrO}_{4}^{2-}+2 \mathrm{H}^{+} \rightarrow \mathrm{Cr}_{2} \mathrm{O}_{7}^{2-}+\mathrm{H}_{2} \mathrm{O}\) (b) \(\mathrm{Fe}+\mathrm{NO}_{3}^{-}+4 \mathrm{H}^{+} \rightarrow \mathrm{Fe}^{3+}+\mathrm{NO}+2 \mathrm{H}_{2} \mathrm{O}\) (c) \(2 \mathrm{C}_{2} \mathrm{H}_{6}+7 \mathrm{O}_{2} \rightarrow 4 \mathrm{CO}_{2}+6 \mathrm{H}_{2} \mathrm{O}\) (d) \(2 \mathrm{AgBr} \rightarrow 2 \mathrm{Ag}+\mathrm{Br}_{\dot{2}}\)

The following reaction is responsible for producing electricity in your car battery (often called a lead storage battery): \(\mathrm{Pb}+\mathrm{PbO}_{2}+2 \mathrm{H}_{2} \mathrm{SO}_{4} \rightarrow 2 \mathrm{PbSO}_{4}+2 \mathrm{H}_{2} \mathrm{O}\) (a) Assign an oxidation state to each atom. (Hint: For \(\mathrm{PbSO}_{4}\), you can figure out the charge of the \(\mathrm{Pb}\) if you remember that the charge of the sulfate ion is \(-2\left(\mathrm{SO}_{4}^{2-}\right)\). Then use shortcut Rule 7 to get the oxidation state of the \(\mathrm{Pb}\) in \(\mathrm{PbSO}_{4}\).) (b) Identify the atom that gets oxidized and the atom that gets reduced. (c) Identify the reactant that is the oxidizing agent and the reactant that is the reducing agent.

Hydrogen gas burns very well in the presence of oxygen to give water: \(2 \mathrm{H}_{2}+\mathrm{O}_{2} \rightarrow 2 \mathrm{H}_{2} \mathrm{O}\) In principle, should it be possible to use this chemical reaction to produce electricity? Explain.

In the upper atmosphere, sunlight can convert oxygen to ozone: \(2 \mathrm{O}_{2} \rightarrow \mathrm{O}_{3}+\mathrm{O}\) Ozone Is this a redox reaction? Completely justify your answer.

If you put a piece of iron in an aqueous solution of blue \(\mathrm{Cu}^{2+}\) ions, the spontaneous redox reaction \(\mathrm{Fe}+\mathrm{Cu}^{2+} \rightarrow \mathrm{Fe}^{2+}+\mathrm{Cu}\) will occur. An aqueous solution of \(\mathrm{Fe}^{2+}\) ions is red-brown. (a) What is oxidized? (b) What is reduced? (c) What is the oxidizing agent? (d) What is the reducing agent? (e) What will happen to the piece of iron over time? (f) Do the electrons move from the oxidizing agent to the reducing agent or from the reducing agent to the oxidizing agent? (g) How could you use this reaction to make a battery? (Explain and show your battery in a diagram, using a nail and a penny as your source of iron and copper, respectively.)