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Chapter 4: Problem 68

Write balanced net ionic equations for the following reactions in acid solution. (a) Nitrogen oxide and hydrogen gases react to form ammonia gas and steam. (b) Hydrogen peroxide reacts with an aqueous solution of sodium hypochlorite to form oxygen and chlorine gases. (c) Zinc metal reduces the vanadyl ion \(\left(\mathrm{VO}^{2+}\right)\) to vanadium(III) ions. Zinc ions are also formed.

Short Answer

Expert verified
Question: Provide the balanced net ionic equation for the reaction between nitrogen oxide and hydrogen gases to form ammonia gas and steam in acidic solution. Answer: The balanced net ionic equation for the reaction is: $$2\,H_2(\text{g}) + 3\text{NO(g)} + 4\,H^+(aq) \rightarrow 2\text{NH}_3(\text{g})+\text{ H}_2\text{O}(\text{l})$$

Step by step solution

01

Write the unbalanced equation

The unbalanced equation for the reaction is: $$\text{NO}(\text{g}) + \text{H}_2(\text{g}) \rightarrow \text{NH}_3(\text{g}) + \text{H}_2\text{O}(\text{l})$$
02

Determine oxidation states of elements

Assign oxidation states to all the elements in the unbalanced equation: N: +2 (in NO), -3 (in NH3), O: -2 (in H2O), H: +1 (in H2 and H2O, and NH3).
03

Identify the redox half-reactions

The oxidation half-reaction is H2(g) -> H2O(l). The reduction half-reaction is NO(g) -> NH3(g).
04

Balance atoms and charges in half-reactions

Balance the two half-reactions, separately: Oxidation half-reaction: $$2\,H_2(\text{g}) \rightarrow 4\,\text{H}^+(aq) + 4\text{ e}^{-}$$ Reduction half-reaction: $$3\text{NO (g)} + 8\text{ H}^{+}(aq) + 3\text{ e}^{-}\rightarrow 2\text{NH}_3(\text{g})+\text{ H}_2\text{O}(\text{l})$$
05

Combine the balanced half-reactions and cancel spectator ions

Combine the balanced half-reactions to obtain the balanced net ionic equation: $$2\,H_2(\text{g}) + 3\text{NO(g)} + 8\text{ H}^{+}(aq) \rightarrow 4\,\text{H}^+(aq) + 2\text{NH}_3(\text{g})+\text{ H}_2\text{O}(\text{l})$$ Cancel out the spectator ions (4H+) to get the final balanced net ionic equation: $$2\,H_2(\text{g}) + 3\text{NO(g)} + 4\,H^+(aq) \rightarrow 2\text{NH}_3(\text{g})+\text{ H}_2\text{O}(\text{l})$$ (b) Hydrogen peroxide reacts with an aqueous solution of sodium hypochlorite to form oxygen and chlorine gases.
06

Write the unbalanced equation

The unbalanced equation for the reaction is: $$\text{H}_2\text{O}_2(\text{aq}) + \text{NaOCl}(\text{aq}) \rightarrow \text{O}_2(\text{g}) + \text{Cl}_2(\text{g}) + \text{Na}^{+}(\text{aq}) + \text{H}^{+}(\text{aq})$$
07

Determine oxidation states of elements

Assign oxidation states to all the elements in the unbalanced equation: O: -1 (in H2O2), -2 (in NaOCl), 0 (in O2), Cl: +1 (in NaOCl), 0 (in Cl2), Na: +1 (in NaOCl), H: +1 (in H2O2 and H+).
08

Identify the redox half-reactions

The oxidation half-reaction is H2O2(aq) -> O2(g). The reduction half-reaction is NaOCl(aq) -> Cl2(g).
09

Balance atoms and charges in half-reactions

Balance the two half-reactions, separately: Oxidation half-reaction: $$2\,\text{H}_2\text{O}_2(aq) \rightarrow 2\,\text{O}_2(\text{g}) + 4\text{ H}^{+}(aq) + 4\text{ e}^{-}$$ Reduction half-reaction: $$2\,\text{OCl}^{-}(aq) + 2\,H^{+}(aq) + 2\,\text{e}^{-} \rightarrow \text{Cl}_2(\text{g}) + 2\,H_2\text{O}(l)$$
10

Combine the balanced half-reactions and cancel spectator ions

Combine the balanced half-reactions to obtain the balanced net ionic equation: $$2\,\text{H}_2\text{O}_2(aq) + 2\,\text{OCl}^{-}(aq) + 2\,\text{H}^{+}(aq) \rightarrow 2\,\text{O}_2(\text{g}) + 4\text{ H}^{+}(aq) + \text{Cl}_2(\text{g}) + 2\,H_2\text{O}(l)$$ Cancel out the spectator ions (2H+) to get the final balanced net ionic equation: $$2\,\text{H}_2\text{O}_2(aq) + 2\,\text{OCl}^{-}(aq) \rightarrow 2\,\text{O}_2(\text{g}) + 2\text{H}^{+}(aq) + \text{Cl}_2(\text{g}) + 2\,H_2\text{O}(l)$$ (c) Zinc metal reduces the vanadyl ion (\(\text{VO}^{2+}\)) to vanadium(III) ions. Zinc ions are also formed.
11

Write the unbalanced equation

The unbalanced equation for the reaction is: $$\text{Zn}(s) + \text{VO}^{2+}(aq) \rightarrow \text{V}^{3+}(aq) + \text{Zn}^{2+}(aq)$$
12

Determine oxidation states of elements

Assign oxidation states to all the elements in the unbalanced equation: Zn: 0 (in Zn), +2 (in Zn2+), V: +5 (in VO2+), +3 (in V3+), O: -2 (in VO2+).
13

Identify the redox half-reactions

The oxidation half-reaction is Zn(s) -> Zn2+(aq). The reduction half-reaction is VO2+(aq) -> V3+(aq).
14

Balance atoms and charges in half-reactions

Balance the two half-reactions, separately: Oxidation half-reaction: $$\text{Zn}(s) \rightarrow \text{Zn}^{2+}(aq) + 2\text{ e}^{-}$$ Reduction half-reaction: $$2\,\text{VO}^{2+}(aq) + 6\,\text{H}^{+}(aq) + 2\,\text{e}^{-} \rightarrow 2\,\text{V}^{3+}(aq) + \text{H}_2\text{O}(l)$$
15

Combine the balanced half-reactions and cancel spectator ions

Combine the balanced half-reactions to obtain the balanced net ionic equation: $$\text{Zn}(s) + 2\,\text{VO}^{2+}(aq) + 6\,\text{H}^{+}(aq) \rightarrow \text{Zn}^{2+}(aq) + 2\,\text{V}^{3+}(aq) + \text{H}_2\text{O}(l)$$ There are no spectator ions to cancel in this case. The final balanced net ionic equation is: $$\text{Zn}(s) + 2\,\text{VO}^{2+}(aq) + 6\,\text{H}^{+}(aq) \rightarrow \text{Zn}^{2+}(aq) + 2\,\text{V}^{3+}(aq) + \text{H}_2\text{O}(l)$$

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Most popular questions from this chapter

Consider the following generic equation $$ \mathrm{OH}^{-}(a q)+\mathrm{HB}(a q) \longrightarrow \mathrm{B}^{-}(a q)+\mathrm{H}_{2} \mathrm{O} $$ For which of the following pairs would this be the correct prototype equation for the acid-base reaction in solution? If it is not correct, write the proper equation for the acid-base reaction between the pair. (a) hydrochloric acid and pyridine, \(\mathrm{C}_{5} \mathrm{H}_{5} \mathrm{~N}\) (b) sulfuric acid and rubidium hydroxide (c) potassium hydroxide and hydrofluoric acid (d) ammonia and hydriodic acid (e) strontium hydroxide and hydrocyanic acid

Wlassify each of the following half-equations as oxidation or reduction and balance. (a) (acidic) \(\quad \mathrm{Mn}^{2+}(a q) \longrightarrow \mathrm{MnO}_{4}^{-}(a q)\) (b) (basic) \(\quad \mathrm{CrO}_{4}^{2-}(a q) \longrightarrow \mathrm{Cr}^{3+}(a q)\) (c) (basic) \(\quad \mathrm{PbO}_{2}(s) \longrightarrow \mathrm{Pb}^{2+}(a q)\) (d) (acidic) \(\quad \mathrm{ClO}_{2}^{-}(a q) \longrightarrow \mathrm{ClO}^{-}(a q)\)

For an acid-base reaction, what is the reacting species, that is, the ion or molecule that appears in the chemical equation, in the following acids? (a) perchloric acid (b) hydriodic acid (c) nitrous acid (d) nitric acid (e) lactic acid \(\left(\mathrm{HC}_{3} \mathrm{H}_{5} \mathrm{O}_{3}\right)\)

Write a balanced net ionic equation for each of the following acid-base reactions in water. (a) acetic acid \(\left(\mathrm{HC}_{2} \mathrm{H}_{3} \mathrm{O}_{2}\right)\) with strontium hydroxide (b) diethylamine, \(\left(\mathrm{C}_{2} \mathrm{H}_{5}\right)_{2} \mathrm{NH}\), with sulfuric acid (c) aqueous hydrogen cyanide (HCN) with sodium hydroxide

Write net ionic equations to explain the formation of (a) a white precipitate when solutions of calcium sulfate and sodium carbonate are mixed. (b) two different precipitates formed when solutions of iron(III) sulfate and barium hydroxide are mixed.
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