8.3: Electrochemistry- Cells and Batteries

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Learning Objectives

Define electrochemistry.
Describe the basic components of electrochemical cells.
List some of the characteristics, applications and limitations of cells and batteries.
Know the difference between galvanic and electrolytic cells.
Define electrolysis and list several of its applications.

Metal exposed to the outside elements will usually corrode if not protected. The corrosion process is a series of redox reactions involving the metal of the sculpture. In some situations, the metals are deliberately left unprotected so that the surface will undergo changes that may enhance the esthetic value of the work.

Electrochemical Reactions

Chemical reactions either absorb or release energy, which can be in the form of electricity. Electrochemistry is a branch of chemistry that deals with the interconversion of chemical energy and electrical energy. Electrochemistry has many common applications in everyday life. All sorts of batteries, from those used to power a flashlight to a calculator to an automobile, rely on chemical reactions to generate electricity. Electricity is used to plate objects with decorative metals like gold or chromium. Electrochemistry is important in the transmission of nerve impulses in biological systems. Redox chemistry, the transfer of electrons, is behind all electrochemical processes.

An electrochemical cell is any device that converts chemical energy into electrical energy or electrical energy into chemical energy. There are three components that make up an electrochemical reaction. There must be a solution where redox reactions can occur. These reactions generally take place in water to facilitate electron and ion movement. A conductor must exist for electrons to be transferred. This conductor is usually some kind of wire so that electrons can move from one site to another. Ions also must be able to move through some form of salt bridge that facilitates ion migration.

Luigi Galvani

Luigi Galvani (1737 - 1798) was an Italian physician and scientist who did research on nerve conduction in animals. His accidental observation of the twitching of frog legs when they were in contact with an iron scalpel while the legs hung on copper hooks led to studies on electrical conductivity in muscles and nerves. He believed that animal tissues contained an "animal electricity" similar to the natural electricity that caused lightning to form.

Alessandro Volta developed the first "voltaic cell" in 1800. This battery consisted of alternating disks of zinc and silver with pieces of cardboard soaked in brine separating the disks. Since there were no voltmeters at the time (and no idea that the electric current was due to electron flow), Volta had to rely on another measure of battery strength: the amount of shock produced (it's never a good idea to test things on yourself). He found that the intensity of the shock increased with the number of metal plates in the system. Devices with twenty plates produced a shock that was quite painful. It's a good thing we have voltmeters today to measure electric current instead of the "stick your finger on this and tell me what you feel" method.

Galvanic (Voltaic) Cells

Galvanic cells, also known as voltaic cells, are electrochemical cells in which spontaneous oxidation-reduction reactions produce electrical energy. In writing the equations, it is often convenient to separate the oxidation-reduction reactions into half-reactions to facilitate balancing the overall equation and to emphasize the actual chemical transformations.

Consider what happens when a clean piece of copper metal is placed in a solution of silver nitrate (Figure \(\PageIndex{1}\)). As soon as the copper metal is added, silver metal begins to form and copper ions pass into the solution. The blue color of the solution on the far right indicates the presence of copper ions. The reaction may be split into its two half-reactions. Half-reactions separate the oxidation from the reduction, so each can be considered individually.

\[\begin{align*}
&\textrm{oxidation: }\ce{Cu}(s)⟶\ce{Cu^2+}(aq)+\ce{2e-}\\
&\underline{\textrm{reduction: }2×(\ce{Ag+}(aq)+\ce{e-}⟶\ce{Ag}(s))\hspace{40px}\ce{or}\hspace{40px}\ce{2Ag+}(aq)+\ce{2e-}⟶\ce{2Ag}(s)}\\
&\textrm{overall: }\ce{2Ag+}(aq)+\ce{Cu}(s)⟶\ce{2Ag}(s)+\ce{Cu^2+}(aq)
\end{align*} \nonumber \]

The equation for the reduction half-reaction had to be doubled so the number electrons “gained” in the reduction half-reaction equaled the number of electrons “lost” in the oxidation half-reaction.

Galvanic or voltaic cells involve spontaneous electrochemical reactions in which the half-reactions are separated (Figure \(\PageIndex{2}\)) so that current can flow through an external wire. The beaker on the left side of the figure is called a half-cell, and contains a 1 M solution of copper(II) nitrate [Cu(NO₃)₂] with a piece of copper metal partially submerged in the solution. The copper metal is an electrode. The copper is undergoing oxidation; therefore, the copper electrode is the anode. The anode is connected to a voltmeter with a wire and the other terminal of the voltmeter is connected to a silver electrode by a wire. The silver is undergoing reduction; therefore, the silver electrode is the cathode. The half-cell on the right side of the figure consists of the silver electrode in a 1 M solution of silver nitrate (AgNO₃). At this point, no current flows—that is, no significant movement of electrons through the wire occurs because the circuit is open. The circuit is closed using a salt bridge, which transmits the current with moving ions.

The salt bridge consists of a concentrated, nonreactive, electrolyte solution such as the sodium nitrate (NaNO₃) solution used in this example. As electrons flow from left to right through the electrode and wire, nitrate ions (anions) pass through the porous plug on the left into the copper(II) nitrate solution. This keeps the beaker on the left electrically neutral by neutralizing the charge on the copper(II) ions that are produced in the solution as the copper metal is oxidized. At the same time, the nitrate ions are moving to the left, sodium ions (cations) move to the right, through the porous plug, and into the silver nitrate solution on the right. These added cations “replace” the silver ions that are removed from the solution as they were reduced to silver metal, keeping the beaker on the right electrically neutral.

When the electrochemical cell is constructed in this fashion, a positive cell potential (0.46 V) indicates a spontaneous reaction and that the electrons are flowing from the left to the right.

Key Features of a Standard Galvanic Cell (as shown in Figure \(\PageIndex{2}\)

Electrons flow from the anode to the cathode: left to right in the standard galvanic cell.
The electrode in the left half-cell is the anode because oxidation occurs here. The name refers to the flow of anions in the salt bridge toward it.
The electrode in the right half-cell is the cathode because reduction occurs here. The name refers to the flow of cations in the salt bridge toward it.
Oxidation occurs at the anode (the left half-cell in the figure).
Reduction occurs at the cathode (the right half-cell in the figure).
The salt bridge must be present to close (complete) the circuit and both an oxidation and reduction must occur for current to flow.

Example \(\PageIndex{1}\): Chromium-Copper Galvanic Cell

Consider a galvanic cell consisting of

\(\ce{2Cr}(s)+\ce{3Cu^2+}(aq)⟶\ce{2Cr^3+}(aq)+\ce{3Cu}(s)\)

Write the oxidation and reduction half-reactions. Which reaction occurs at the anode? The cathode?

Solution

By inspection, Cr is oxidized when three electrons are lost to form Cr³⁺, and Cu²⁺ is reduced as it gains two electrons to form Cu. Balancing the charge gives

\(\begin{align*}
&\textrm{oxidation: }\ce{2Cr}(s)⟶\ce{2Cr^3+}(aq)+\ce{6e-}\\
&\textrm{reduction: }\ce{3Cu^2+}(aq)+\ce{6e-}⟶\ce{3Cu}(s)\\
&\overline{\textrm{overall: }\ce{2Cr}(s)+\ce{3Cu^2+}(aq)⟶\ce{2Cr^3+}(aq)+\ce{3Cu}(s)}
\end{align*}\)

Cell notation uses the simplest form of each of the equations, and starts with the reaction at the anode. No concentrations were specified so:

\(\ce{Cr}(s)│\ce{Cr^3+}(aq)║\ce{Cu^2+}(aq)│\ce{Cu}(s).\)

Oxidation occurs at the anode and reduction at the cathode.

Exercise \(\PageIndex{1}\) Copper-Zinc Galvanic Cell

Consider a galvanic cell consisting of

\(\ce{Zn}(s)+\ce{Cu^2+}(aq)⟶\ce{Zn^2+}(aq)+\ce{Cu}(s)\)

Write the oxidation and reduction half-reactions. Which reaction occurs at the anode? The cathode?

Answer: From the information given in the problem:
\(\begin{align*}
&\textrm{anode (oxidation): }\ce{Zn}(s)⟶\ce{Zn^2+}(aq)+\ce{2e-}\\
&\textrm{cathode (reduction): }\ce{Cu^2+}(aq)+\ce{2e-}⟶\ce{Cu}(s)\\
&\overline{\textrm{overall: }\ce{Zn}(s)+\ce{Cu^2+}(aq)⟶\ce{Zn^2+}(aq)+\ce{Cu}(s)}
\end{align*}\)

Batteries

A battery is an electrochemical cell or series of cells that produces an electric current. In principle, any galvanic cell could be used as a battery. An ideal battery would never run down, produce an unchanging voltage, and be capable of withstanding environmental extremes of heat and humidity. Real batteries strike a balance between ideal characteristics and practical limitations. For example, the mass of a car battery is about 18 kg or about 1% of the mass of an average car or light-duty truck. This type of battery would supply nearly unlimited energy if used in a smartphone, but would be rejected for this application because of its mass. Thus, no single battery is “best” and batteries are selected for a particular application, keeping things like the mass of the battery, its cost, reliability, and current capacity in mind. There are two basic types of batteries: primary and secondary. A few batteries of each type are described next.

Visit this site to learn more about batteries.

Dry Cells (Primary Batteries)

Primary batteries are single-use batteries because they cannot be recharged. A common primary battery is the dry cell (Figure \(\PageIndex{1}\)). The dry cell is a zinc-carbon battery. The zinc can serves as both a container and the negative electrode. The positive electrode is a rod made of carbon that is surrounded by a paste of manganese(IV) oxide, zinc chloride, ammonium chloride, carbon powder, and a small amount of water. The reaction at the anode can be represented as the ordinary oxidation of zinc:

\[\ce{Zn}(s)⟶\ce{Zn^2+}(aq)+\ce{2e-} \nonumber \]

The reaction at the cathode is more complicated, in part because more than one reaction occurs. The series of reactions that occurs at the cathode is approximately

\[\ce{2MnO2}(s)+\ce{2NH4Cl}(aq)+\ce{2e-}⟶\ce{Mn2O3}(s)+\ce{2NH3}(aq)+\ce{H2O}(l)+\ce{2Cl-} \nonumber \]

The overall reaction for the zinc–carbon battery can be represented as

\[\ce{2MnO2}(s) + \ce{2NH4Cl}(aq) + \ce{Zn}(s) ⟶ \ce{Zn^2+}(aq) + \ce{Mn2O3}(s) + \ce{2NH3}(aq) + \ce{H2O}(l) + \ce{2Cl-} \nonumber \]

with an overall cell potential which is initially about 1.5 V, but decreases as the battery is used. It is important to remember that the voltage delivered by a battery is the same regardless of the size of a battery. For this reason, D, C, A, AA, and AAA batteries all have the same voltage rating. However, larger batteries can deliver more moles of electrons.

Take Note

The dry cell is not very efficient in producing electrical energy because only the relatively small fraction of the \(MnO_2\) that is near the cathode is actually reduced and only a small fraction of the zinc cathode is actually consumed as the cell discharges. In addition, dry cells have a limited shelf life because the \(Zn\) anode reacts spontaneously with \(NH_4Cl\) in the electrolyte, causing the case to corrode and allowing the contents to leak out (Figure \(\PageIndex{2}\)).

Figure \(\PageIndex{4}\) A cross section of an alkaline battery.

Alkaline batteries (Figure \(\PageIndex{4}\)) were developed in the 1950s partly to address some of the performance issues with zinc–carbon dry cells. They are manufactured to be exact replacements for zinc-carbon dry cells. As their name suggests, these types of batteries use alkaline electrolytes, often potassium hydroxide.

An alkaline battery can deliver about three to five times the energy of a zinc-carbon dry cell of similar size. Alkaline batteries are prone to leaking potassium hydroxide, so these should also be removed from devices for long-term storage. While some alkaline batteries are rechargeable, most are not.