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Chapter 10: Problem 106

You are trying to determine the identity of an unknown metal. You place a strip of it in a solution of \(\mathrm{Mg}^{2+}\), and no reaction occurs. You then place a strip of it in a solution of \(\mathrm{Zn}^{2+}\), and zinc metal plates out on the strip. Name one possibility for the identity of the unknown metal.

Short Answer

Expert verified
One possibility for the identity of the unknown metal is aluminum (Al), as it is more reactive than zinc but less reactive than magnesium.

Step by step solution

01

Understand the reactivity series

The reactivity series is a list of metals arranged in order of their reactivity, from the most reactive to the least reactive. Reactivity here means how easily a metal loses electrons to form positive ions (cations). Metals with a higher reactivity will displace metals with a lower reactivity from their salts in solution.
02

Evaluate the reaction with \(\mathrm{Mg}^{2+}\)

Since there is no reaction when the unknown metal is placed in a solution of \(\mathrm{Mg}^{2+}\), this means that the unknown metal is less reactive than magnesium. Otherwise, it would have displaced \(\mathrm{Mg}^{2+}\), causing a reaction.
03

Evaluate the reaction with \(\mathrm{Zn}^{2+}\)

In the second reaction, the unknown metal is placed in a solution of \(\mathrm{Zn}^{2+}\), and zinc metal plates out on the strip. This means that the unknown metal is more reactive than zinc and displaces \(\mathrm{Zn}^{2+}\) in the solution.
04

Identify a possible unknown metal using the reactivity series

Using the reactivity series, we can find a metal that is less reactive than magnesium but more reactive than zinc. One possibility for the unknown metal is aluminum (Al). Aluminum is more reactive than zinc but less reactive than magnesium, fitting the conditions given in the problem. Therefore, one possibility for the identity of the unknown metal is aluminum (Al).

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Most popular questions from this chapter

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.)

The spontaneous redox reaction \(\mathrm{Mn}+\mathrm{Cd}^{2+} \rightarrow \mathrm{Mn}^{2+}+\mathrm{Cd}\) takes place in a battery. (a) What is the oxidizing agent? (b) What is the reducing agent? (c) Which metal is the cathode? (d) Which metal is the anode?

What happens when you place a less active metal in a solution of ions of a more active metal?

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

Using the EMF series on page 391, decide which of the following redox reactions is spontaneous. Explain your answer. (a) \(3 \mathrm{Ag}+\mathrm{Au}^{3+} \rightarrow \mathrm{Au}+3 \mathrm{Ag}^{+}\) (b) \(\mathrm{Au}+3 \mathrm{Ag}^{+} \rightarrow 3 \mathrm{Ag}+\mathrm{Au}^{3+}\)
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