Question

(a) Based on standard reduction potentials, would you expect copper metal to oxidize under standard conditions in the presence of oxygen and hydrogen ions? (b) When the Statue of Liberty was refurbished, Teflon spacers were placed between the iron skeleton and the copper metal on the surface of the statue. What role do these spacers play?

Short answer

(a) By calculating the overall standard reduction potential (Eº) for the net redox reaction between copper and oxygen in the presence of hydrogen ions, we find that Eº(net) = +0.89 V. Since this value is positive, the reaction is spontaneous under standard conditions, meaning that copper metal will oxidize in the presence of oxygen and hydrogen ions. (b) The Teflon spacers between the iron skeleton and the copper surface of the Statue of Liberty help to prevent galvanic corrosion. Galvanic corrosion occurs when two dissimilar metals are in direct contact, causing the less noble metal (iron, in this case) to corrode. The Teflon spacers act as insulators, preventing electrical contact between the two metals and interrupting the electrical path for the galvanic cell, thus preventing corrosion and preserving the structural integrity of the statue.

Step-by-step solution

Part (a): Using Standard Reduction Potentials

First, we need to find the standard reduction potentials for the relevant half-reactions: 1. \(Cu^{2+}(aq) + 2e^- \rightarrow Cu(s)\) has a standard reduction potential Eº = +0.34 V. 2. \(O_2(g) + 4H^+(aq) + 4e^- \rightarrow 2H_2O(l)\) has a standard reduction potential Eº = +1.23 V. To determine whether copper will oxidize in the presence of oxygen and hydrogen ions under standard conditions, we must determine whether a net spontaneous redox reaction can occur between them. To do this, we need to first write the oxidation half-reaction for copper. By reversing the first half-reaction, we get: \(Cu(s) \rightarrow Cu^{2+}(aq) + 2e^-\), which has a standard reduction potential of -0.34 V (enthalpy changes its sign since the reaction is reversed). Now, we can combine the two half-reactions (copper's oxidation half-reaction and oxygen's reduction half-reaction): \(Cu(s) + 1/2 O_2(g) + 2H^+(aq) \rightarrow Cu^{2+}(aq) + H_2O(l)\) Next, we'll need to calculate the overall standard reduction potential (Eº) for the net redox reaction using the individual standard reduction potentials: Eº(net) = Eº(reduction) + Eº(oxidation) Eº(net) = (+1.23 V) + (-0.34 V) = +0.89 V Since Eº(net) is positive, the reaction is spontaneous under standard conditions. Therefore, copper metal will oxidize in the presence of oxygen and hydrogen ions under standard conditions.

Part (b): Role of Teflon Spacers

The Teflon spacers placed between the iron skeleton and the copper surface of the Statue of Liberty play a crucial role in preventing galvanic corrosion. Galvanic corrosion occurs when two dissimilar metals are in contact with each other and an electrolyte, forming a galvanic cell. The less noble metal will act as an anode and corrode, while the nobler metal will act as a cathode and remain unchanged. In this case, iron is the less noble metal and would have been more prone to corrosion if it was in direct contact with the copper surface. The Teflon spacers act as an insulator, preventing the electrical contact between the two metals. By doing so, they interrupt the electrical path for the galvanic cell to form, thus preventing the galvanic corrosion from taking place. This helps to maintain the structural integrity of the Statue of Liberty by preventing damage due to corrosion.

Key concepts

Redox Reactions

In the world of chemistry, redox reactions involve the transfer of electrons between two species. The term "redox" comes from the words "reduction" and "oxidation." In these reactions, one element is oxidized, meaning it loses electrons, while the other is reduced, meaning it gains electrons. This exchange is fundamental to numerous chemical processes, including those that power batteries and support biological functions.

To understand redox reactions, it's essential to know about half-reactions, which show either the oxidation or the reduction part separately. For the exercise given, copper oxidizes by losing electrons, shown as:

• Copper oxidation: \(Cu(s) \rightarrow Cu^{2+}(aq) + 2e^-\).

Oxygen, on the other hand, reduces by gaining electrons:

• Oxygen reduction: \(O_2(g) + 4H^+(aq) + 4e^- \rightarrow 2H_2O(l)\).

By combining these half-reactions, you create a full redox equation that helps predict whether a reaction is spontaneous. The spontaneity is indicated by the net standard reduction potential \(E°(net)\). When \(E°(net)\) is positive, the reaction is spontaneous, as in the case with copper in the presence of oxygen and hydrogen ions under standard conditions.

Galvanic Corrosion

Galvanic corrosion occurs when two different metals are electrically connected in the presence of an electrolyte. This condition sets up a galvanic cell where one of the metals acts as an anode (corrodes) and the other as a cathode (remains less affected). This process is harmful because it can quickly deteriorate structures or objects composed of more than one type of metal.

The physics behind this phenomenon involve electric potentials. The metal with the higher tendency to lose electrons will oxidize more easily and suffer corrosion. This makes it act as the anode. In contrast, the metal with a lesser tendency to oxidize becomes the cathode and is protected. For example, in the Statue of Liberty, iron acts as the anode, and copper as the cathode.

To prevent galvanic corrosion, methods like using an insulating material between the metals are employed. In the refurbishment of the Statue of Liberty, Teflon spacers were used between copper and iron surfaces. These spacers break the electrical circuit and prevent the electrochemical reaction from occurring, thus stopping galvanic corrosion.

Electrochemistry

Electrochemistry is the branch of chemistry that deals with the relationships between electrical energy and chemical changes. It covers a broad spectrum of processes, from simple reactions to complex ones that power our technology. One key aspect of electrochemistry is the study of how chemical energy is converted to electrical energy and vice versa.

Standard reduction potentials are crucial in electrochemistry. They help us determine the potential difference between substances in a reaction, predicting if a reaction will occur spontaneously. The greater the positive standard reduction potential, the stronger the substance’s tendency to gain electrons and undergo reduction. The opposite is true for oxidation.

In practice, electrochemical principles are employed in batteries, fuel cells, and electroplating, among other applications. In the given exercise, knowing the standard reduction potentials of the involved reactions was necessary to calculate whether copper would oxidize in the presence of oxygen and hydrogen ions. This showcases how theoretical concepts translate into practical applications.