A piece of magnesium ribbon and a copper wire are partially immersed in a \(0.1 \mathrm{M} \mathrm{HCl}\) solution in a beaker. The metals are joined externally by another piece of metal wire. Bubbles are seen to evolve at both the \(\mathrm{Mg}\) and \(\mathrm{Cu}\) surfaces. (a) Write equations representing the reactions occurring at the metals. (b) What visual evidence would you seek to show that \(\mathrm{Cu}\) is not oxidized to \(\mathrm{Cu}^{2+} ?\) (c) At some stage, \(\mathrm{NaOH}\) solution is added to the beaker to neutralize the HCl acid. Upon further addition of \(\mathrm{NaOH},\) a white precipitate forms. What is it?

When magnesium (\textbf{Mg}) comes into contact with hydrochloric acid (\textbf{HCl}), a chemical reaction occurs where \textbf{Mg} is oxidized, resulting in the liberation of hydrogen gas (\textbf{H}\textsubscript{2}) and the formation of magnesium chloride (\textbf{MgCl}\textsubscript{2}). This is described by the equation:
\[ Mg + 2HCl \rightarrow MgCl_2 + H_2 \]
The reaction is quite vigorous, and the evolution of \textbf{H}\textsubscript{2} gas can be visibly observed as bubbles. The hydrogen gas is formed due to Mg donating its two valence electrons to the two \textbf{H}\textsuperscript{+} ions in \textbf{HCl}. This transformation is a textbook example of a single displacement reaction where magnesium displaces hydrogen from hydrochloric acid forming a salt and releasing hydrogen gas. It’s important to monitor the temperature and color changes during this type of reaction, as they are indicative of the reaction's progression and intensity.

If a student is having trouble understanding this concept, it might be helpful to visualize the reaction by drawing the exchange of electrons and emphasizing the idea that \textbf{Mg} is actually donating electrons to the \textbf{H}\textsuperscript{+} ions, thus reducing them to \textbf{H}\textsubscript{2} gas.

• Copper (\textbf{Cu}) does not typically react with \textbf{HCl} to produce \textbf{Cu}\textsuperscript{2+} ions because it is less reactive than many other metals.
• When \textbf{Cu} is seemingly reacting with \textbf{HCl}, it may be due to the presence of an oxidizing agent, which is not specified in this exercise scenario.
• In this case, the \textbf{Cu} wire remains largely unchanged and the absence of blue color in the solution indicates that \textbf{Cu}\textsuperscript{2+} ions are not formed.
• If bubbles are observed at the \textbf{Cu} surface, the reaction is likely generating hydrogen gas due to a secondary influence such as an interconnected metal wire that can carry electrons from magnesium's reaction.

The fact that the copper remains unchanged can be demonstrated by the solution's lack of color change. A stable copper color suggests that oxidation to \textbf{Cu}\textsuperscript{2+} has not occurred:

\[ \text{(no reaction with Cu and HCl under normal conditions)} \]
To enhance understanding, it’s crucial to explain that oxidation involves the loss of electrons, and in the case of copper oxidation, the metal would need to lose two electrons to form \textbf{Cu}\textsuperscript{2+}. Without an appropriate oxidizing agent, copper remains in its elemental form.

In the scenario where sodium hydroxide (\textbf{NaOH}) is added to the solution containing \textbf{Mg}\textsuperscript{2+} ions, a white precipitate forms upon continued addition beyond the point of neutralizing the \textbf{HCl}. This precipitate is magnesium hydroxide (\textbf{Mg(OH)\textsubscript{2}}), a result of the following precipitation reaction:
\[ Mg^{2+} + 2OH^- \rightarrow Mg(OH)_2 \]
A precipitation reaction takes place when two soluble substances react to form an insoluble solid, known as the precipitate. The insoluble \textbf{Mg(OH)\textsubscript{2}} emerges due to the \textbf{Mg}\textsuperscript{2+} ions from dissolved \textbf{Mg} reacting with \textbf{OH}\textsuperscript{-} ions from the excess \textbf{NaOH}. This is useful to demonstrate the solubility rules and the concept of ionic equilibria in aqueous solutions. Students struggling with this might benefit from conducting a solubility test experiment or observing a visible precipitation reaction to grasp the concept better. Additionally, understanding the stoichiometry of the reactants and the charge balance can clarify why two hydroxide ions are needed to react with one magnesium ion to form the precipitate.