Balancing the number of electrons by multiplying the oxidation reaction by 3,. We must now check to make sure the charges and atoms on each side of the equation balance:.
We can also use the alternative procedure, which does not require the half-reactions listed in Table P1. Dividing the reaction into two half-reactions,. Step 2: Balancing the atoms other than oxygen and hydrogen,. We now balance the O atoms by adding H 2 O—in this case, to the right side of the reduction half-reaction.
Because the oxidation half-reaction does not contain oxygen, it can be ignored in this step. Again, we can ignore the oxidation half-reaction.
Step 3: We must now add electrons to balance the charges. Step 4: To have the same number of electrons in both half-reactions, we must multiply the oxidation half-reaction by Step 5: Adding the two half-reactions and canceling substances that appear in both reactions,.
Step 6: This is the same equation we obtained using the first method. Thus the charges and atoms on each side of the equation balance. This is the same value that is observed experimentally. As we shall see in Section With a sufficient input of electrical energy, virtually any reaction can be forced to occur. A galvanic cell with a measured standard cell potential of 0. One beaker contains a strip of gallium metal immersed in a 1 M solution of GaCl 3 , and the other contains a piece of nickel immersed in a 1 M solution of NiCl 2.
The half-reactions that occur when the compartments are connected are as follows:. Given: galvanic cell, half-reactions, standard cell potential, and potential for the oxidation half-reaction under standard conditions. Asked for: standard electrode potential of reaction occurring at the cathode.
With three electrons consumed in the reduction and two produced in the oxidation, the overall reaction is not balanced. Recall, however, that standard potentials are independent of stoichiometry. When the compartments are connected, a potential of 3. When using a galvanic cell to measure the concentration of a substance, we are generally interested in the potential of only one of the electrodes of the cell, the so-called indicator electrode , whose potential is related to the concentration of the substance being measured.
To ensure that any change in the measured potential of the cell is due to only the substance being analyzed, the potential of the other electrode, the reference electrode , must be constant. You are already familiar with one example of a reference electrode: the SHE. The potential of a reference electrode must be unaffected by the properties of the solution, and if possible, it should be physically isolated from the solution of interest. To measure the potential of a solution, we select a reference electrode and an appropriate indicator electrode.
Whether reduction or oxidation of the substance being analyzed occurs depends on the potential of the half-reaction for the substance of interest the sample and the potential of the reference electrode.
The potential of any reference electrode should not be affected by the properties of the solution to be analyzed, and it should also be physically isolated. There are many possible choices of reference electrode other than the SHE.
The SHE requires a constant flow of highly flammable hydrogen gas, which makes it inconvenient to use. Consequently, two other electrodes are commonly chosen as reference electrodes. One is the silver—silver chloride electrode , which consists of a silver wire coated with a very thin layer of AgCl that is dipped into a chloride ion solution with a fixed concentration.
The cell diagram and reduction half-reaction are as follows:. If a saturated solution of KCl is used as the chloride solution, the potential of the silver—silver chloride electrode is 0. That is, 0. A second common reference electrode is the saturated calomel electrode SCE , which has the same general form as the silver—silver chloride electrode. The SCE consists of a platinum wire inserted into a moist paste of liquid mercury Hg 2 Cl 2 ; called calomel in the old chemical literature and KCl.
Although it sounds and looks complex, this cell is actually easy to prepare and maintain, and its potential is highly reproducible. The SCE cell diagram and corresponding half-reaction are as follows:. The glass membrane absorbs protons, which affects the measured potential. This indicates that it is a highly corrosive substance, and only a few materials are resistant.
When trying to understand the degradation of materials in an HCl environment, it becomes very complex because small amounts of metal chlorides and impurities can vastly affect the corrosive nature of the solution.
Oxidizing agents greatly increase the corrosion rate of metals and alloys, and also organic contaminates detrimentally affect plastics and elastomers. The table below is a general guideline of some alloys to consider based on their performance in an HCl environment. As previously mentioned, however, impurities in HCl can drastically alter these guidelines.
Metals such as aluminum, cast iron, steel, copper, and titanium will suffer rapid attack from HCl at all concentrations and temperatures. Most stainless steel grades will be subject to attack, because their chromium content is not sufficient in forming a protective passive layer. Without the passive layer the stainless steel will then begin to corrode actively, which leads to rapid corrosion rates, pitting, and stress corrosion cracking.
The figure below left is a corrosion curve of nickel and reactive alloys with respect to temperature and concentration in HCl. In general, nickel alloys have a better resistance than stainless steel, but still not completely corrosion resistant for most concentrations.
Only C is able handle most concentrations of HCl at room temperature. Similar comparisons of other metals have made it possible to arrange them in order of their increasing electron-donating, or reducing, power. This sequence is known as the electromotive, or activity, series of the metals. The activity series has long been used to predict the direction of oxidation-reduction reactions.
Consider, for example, the oxidation of copper by metallic zinc mentioned above. Zinc is near the top of the activity series, meaning that this metal has a strong tendency to lose electrons. Copper, on the other hand, is a poorer electron donor, and therefore its oxidized form, Cu, is a fairly good electron acceptor. We would therefore expect the following reaction to proceed in the direction already indicated, rather than in the reverse direction:.
Note that the table also takes the replacement of hydrogen H 2 into account. Each half-cell is associated with a potential difference whose magnitude depends on the nature of the particular electrode reaction and on the concentrations of the dissolved species.
The sign of this potential difference depends on the direction oxidation or reduction in which the electrode reaction proceeds.
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