In an acidic solution, O21g2 oxidizes the substance.
The cell does electrical work when a reaction occurs.
Think of it as the work of moving electric charges.
The constant is equal to 96,485 coulombs per mole of electrons.
The unit of welec in equation (19.13) is joules.
If the cell operates reversibly, expression applies.
The work that can be derived from a process is equal to C/.
The number with no units is called the electron number.
The number of electrons transferred in the reaction is written.
The reaction on the left and the reaction on the right are shown.
In considering an overall cell reaction, we must balance the electrons.
Edeg1right2 and Edeg1 left2 include systematic studies of electrolysis.
The meaning of a reversible process was illustrated by Figure 7 on page 262.
The electric current must be drawn from the cell only very slowly in order for the cell to be reversible.
When a mole of Cu2+ is reduced or a mole of H+ is produced, 68.6 kJ of energy is generated.
There are two moles of electrons around the outer circuit.
The value of Edegcell is the same as before, but the electron number is one-half of what was calculated.
The result supports the fact that the standard reduction potential is an intensive property.
The reaction tells us that when Cu2+ is reduced, 32.8 kJ of energy is released and one mole of electrons passes from the anode to the cathode.
The overall equation needs to be separated into two half-cell equations.
The combination of elements to form a compound is what the overall cell reaction is all about.
The equation can be applied to half-cell reac tions and half-cell potentials.
Fe1s2; C/rGdeg is 10.880F2 V - 10.771F2 V.
We can use equation (19.15) again to solve for E Fe deg 3+>Fe.
C/rG 6 0 is our main criterion for change.
If C/rG 6 0 and Ecell 7 0 have the same property, then it's a redox reaction.
The reaction is at equilibrium if Ecell is 0
This is an important point to remember.
There are answers to questions about redox reactions that can be found without going through a complete calculation of Ecell deg.
Most metals do not react with an acid, but a few do, according to the discussion in Chapter 5.
This observation can be explained.
The reduction involves H+ M2+ when a metal reacts with an acid.
Reaction of Al and being reduced to H21g2.
There is a place where Al(s) has dissolved.
H21g2 should be displaced from acidic solutions by these metals.
HNO31aq2 reacts with nitric acid.
The species reduced in the reaction is what we need to identify.
We calculate Ecell deg.
The reaction will occur if Ecell deg is positive.
Under standard-state conditions, Al(s) will remove Cu2+ from the solution.
Their values are unaffected by the choice of coefficients used to balance the equation for the cell reaction because they do not depend on the quantities of materials involved.
No magnesium metal is obtained when salt is added to the water.
Na2S2O8 is an oxidizer used in bleaching.
Laboratory oxidizers such as K2Cr2O7 have been used.
The better the oxidizing agent is, the less the oxidizing agent is reduced.
The Edeg value is used to measure the reduction tendency.
We can qualitatively assess the spontaneity of a particular redox reaction by inspecting standard reduction potentials.
An inexpensive way to produce peroxodisulfates would be to pass O21g2 through an acidic solution.
The reduction half-cell reaction is difficult to maintain because atmospheric oxygen oxidizes Sn2+ to Sn4+.
The three quantities are related to each other.
The constant has a value of 0.025693 J>C.
Figure 19-8 summarizes several important relationships.
We identify the reactant that is reduced and the reactant that is oxidation.
The data in Table 19.1 and Appendix D can be used to get the standard reduction potentials.
We used equation (19.17) to get K from Ecell deg.
The equilibrium constant is greater than one if the cell potential is positive.
We can expect this reaction to go to completion because the equilibrium constant is very large.
The forward reaction should be favored by Zn2+1aq2 +.
The relationship between the cell potential, Ecell, and the concentrations of reactants and products is easy to establish.
The equation was first proposed in 1889.
The third law of thermodynamics is the nernst equation.
Gases and molarities are obtained for the activities of solution components.
We can insert the concentration of the species once we have determined the form of the Nernst equation.
The Edeg is using the Nernst equation.
The Ecell is positive if the reaction is in the direction of the reduction of silver.
Qualitative conclusions reached with Edegcell values often hold over a broad range of nonstandard conditions as well.
We need to calculate Ecell by using equation (19.18) with the concentrations given.
We look up the appropriate standard half-cell potentials after identifying the oxidation and reduced species.
We conclude that the reaction as written is spontaneously.
As long as we balance charge and electron number correctly, we will always get the correct result.
Under standard-state conditions, the following cell is set up.
The cell is made of hydrogen.
There is a SHE on the right.
The pH of the solution in the anode compartment is proportional to Pt on the voltmeter.
The concentration cell always changes when the concentrated solution becomes more concentrated.
The solutions were simply mixed.
In a concentration cell, the natural tendency to increase in a mixing process is used as a means of generating electricity.
It is difficult to make and use a hydrogen electrode.
The Pt metal surface needs to be prepared and maintained, gas pressure needs to be controlled, and the electrode can't be used in the presence of reducing agents.
The solution to these problems will be discussed later in the chapter.
It gives a basis for determining Ksp values for ionic compounds.
Consider the concentration cell.
A saturated solution of silver iodide is placed at the anode.
The measured cell voltage is 0.417 V, and the second silver electrode is placed in a solution with 3 Ag+4.
It is not the most convenient to use the standard hydrogen electrode because it requires highly flammable hydrogen gas to be bubbled over.
A saturated AgI(satd aq) Ag1( 0.
100 M) solution of AgI is in contact with the silver electrode in the anode compartment.
Ksp for AgI is calculated with the data given for the reaction.
We can use the expression for the solubility product to calculate the equilibrium constant once we have determined the concentration of Ag+ ion in the cell.
The realization that the only source of Ag+ and I- is from the AgI present is one of the essential aspects.
The following concentration cell information should be used to calculate the Ksp for PbI2.
A silver wire is covered with a layer of solid silver.
The standard hydrogen electrode has a potential of 0.22233 V, and this one has a potential of 0.22233 V.
The silver-silver chloride electrode has a standard potential of 0.22233 V at 25 degC, since all components are in their standard states.
The whole setup is immersed in either a 1.0 M solution of potassium chloride or a saturated solution, as mercurous chloride is mixed with mercury to form a paste, which is in contact with liquid mercury, Hg(l).
The silver wire is immersed in a solution of KCl.
A porous disc at the bottom of the tube allows contact with a solution of interest.
The inner tube has a Pt wire inserted in it and the outer tube has a small hole in it.
When the glass is dipped into a solution, there is an interaction between the ion and the membrane.
The silver wire's potential is dependent on the solution being tested.
There is a small sintered disc in the side of the outer tube that acts as a salt bridge.
The reduction potential is 0.2412 V if a saturated solution of KCl is used.
Reduction potentials are quoted with respect to a specific reference because a variety of reference electrodes are used.
To measure the pH of a solution, we need a response to changes in 3H.
The standard hydrogen electrode is difficult to use for this purpose and so a simpler and safer one is needed.
A potential develops when the bulb is placed in a solution of unknown pH and the concentration difference across the membranes is similar to a concentration cell.
The cell can be represented as Ag1s2fAgCl1s2fCl-11.0 M2, H11.0 M2fglass, andunknown2f.
The half-cell potentials of the two half-cell reactions do not make a difference to the cell potential.
The source of the potential difference between the two half-cells and the unknown solution is the difference in the molar Gibbs energy between the two half-cells.
After converting the logarithm to base 10 and using the definition of pH as -log 3unknown4 we get Ecell, which is 0.0592 V. The cell potential is measured with a pH meter, a device that converts Ecell to pH and displays the result in units.
The prototype for a large number of membrane electrodes that areselective for a particular ion, such as the ion K+, NH + 4 and many others, was created by German Biologist Max Cremer in 1906.
An example of a flashlight cell is a single voltaic cell with two electrodes and an appropriate electrolyte.
Other batteries have two or more voltaic cells joined in a fashion that increases the total voltage.
The batteries and voltaic cells will be considered in this section.
The cell reaction in a primary cell can't be reversed.
Electricity is not used when 10 batteries per person are converted to products.
A battery with secondary cells can be used for several hundred or more cycles.
The rest of the battery is used to reduce self-discharge or the chemical degradation of rechargeable batteries.
The battery is designed to be used for a long time to deliver high power over a relatively short period of time.
The materials that pass through the battery are used to convert power for portable electronic chemical energy to electric energy.
These types of batteries can be used.
Review flashcards and saved quizzes
Getting your flashcards
You're all caught up!
Looks like there aren't any notifications for you to check up on. Come back when you see a red dot on the bell!