# anode

(noun)

## Definition of anode

The electrode of an electrochemical cell at which oxidation occurs.

Source: Wiktionary - CC BY-SA 3.0

## Examples of anode in the following topics:

• ### Voltaic Cells

• Electrochemical cells have two conductive electrodes (the anode and the cathode).
• The anode is defined as the electrode where oxidation occurs.
• One electrode will undergo oxidation (the anode) and the other will undergo reduction (the cathode).
• The anode attracts anions.
• For the cell depicted in the figure, the anode is zinc and the cathode is copper.
• ### Electrochemical Cell Notation

• The reaction conditions (pressure, temperature, concentration, etc), the anode, the cathode, and the electrode components are all described in this unique shorthand.
• Recall that oxidation takes place at the anode and reduction at the cathode.
• When the anode and cathode are connected by a wire, electrons flow from anode to cathode.
• The cadmium is oxidized (loses electrons) and is the anode.
• The anode half-cell is described first; the cathode half-cell follows.
• ### Electrolytic Properties

• Lone electrons normally cannot pass through the electrolyte; instead, a chemical reaction occurs at the cathode that consumes electrons from the anode.
• Another reaction occurs at the anode, producing electrons that are eventually transferred to the cathode.
• As a result, a negative charge cloud develops in the electrolyte around the cathode, and a positive charge develops around the anode.
• The anode reaction is:$2NaCl \rightarrow 2 Na^{+} + Cl_2 + 2e^{-}$and chlorine gas will be liberated.
• For example, it is possible to oxidize ferrous ions to ferric ions at the anode:$Fe^{2+}(aq) \rightarrow Fe^{3+} (aq) + e^{-}$Neutral molecules can also react at either electrode.
• ### Cathode Rays

• Since the electrons have a negative charge, they are repelled by the cathode and attracted to the anode.
• After the electrons reach the anode, they travel through the anode wire to the power supply and back to the cathode, so cathode rays carry electric current through the tube.
• But at the anode (positive) end of the tube, the glass of the tube itself began to glow.
• By the time the tube was dark, most of the electrons could travel in straight lines from the cathode to the anode end of the tube without a collision.
• When they reached the anode end of the tube, they were travelling so fast that, although they were attracted to it, they often flew past the anode and struck the back wall of the tube.
• ### The Lithium-Ion Battery

• Charging and Discharging The three participants in the electrochemical reactions in a lithium-ion battery are the anode, the cathode, and the electrolyte.
• Both the anode, which is a lithium-containing compound, and the cathode, which is a carbon-containing compound, are materials into which and from which lithium ions can migrate.
• When a lithium-based cell is discharging, the positive lithium ion is extracted from the cathode and inserted into the anode, releasing stored energy in the process.
• The anode is generally one of three materials: a layered oxide (such as lithium cobalt oxide), a polyanion (such as lithium iron phosphate), or a spinel (such as lithium manganese oxide).
• In a lithium-ion battery, the lithium ions are transported to and from the cathode or anode.
• ### Electrolysis of Molten Sodium Chloride

• Let's go through each of the methods to understand the different processes.If sodium chloride is melted (above 801°C) and two electrodes are inserted into the melt as shown in and an electric current is passed through the molten salt, then chemical reactions take place at the electrodes.Sodium ions migrate to the cathode, where electrons enter the melt and are reduced to sodium metal:${Na}^{+} + {e}^{-} \rightarrow Na$Chloride ions migrate the other way, toward the anode, give up their electrons to the anode, and are oxidized to chlorine gas:${Cl}^{-} \rightarrow \frac{1}{2}{Cl}_{2} + {e}^{-}$The overall reaction is the breakdown of sodium chloride into its elements:$2NaCl \rightarrow Na(s) + {Cl}_{2}(g)$Now what happens when we have an aqueous solution of sodium chloride as shown here ?
• The reaction at the cathode is:${H}_{2}O (l) + 2 {e}^{–} \rightarrow {H}_{2}(g) + 2{ OH}^{–}$and at the anode:${Cl}^{–} \rightarrow \frac{1}{2} {Cl}_{2}(g)$The overall reaction is as follows:$NaCl(aq) + {H}_{2}O(l) \rightarrow {Na}^{+}(aq) + {OH}^{-}(aq) + {H}_{2}(g) + \frac{1}{2}{Cl}_{2}(g)$Reduction of Na+ (E° = –2.7 v) is energetically more difficult than the reduction of water (–1.23 v), so in aqueous solution the latter will prevail.
• ### Predicting the Products of Electrolysis

• Positively charged ions (cations) move toward the electron-providing (negative) cathode, whereas negatively charged ions (anions) move toward the positive anode.
• You may have noticed that this is the opposite of a galvanic cell, where the anode is negative and the cathode is positive.
• Oxidation and Reduction Oxidation of ions or neutral molecules occurs at the anode, and reduction of ions or neutral molecules occurs at the cathode.
• For example, it is possible to oxidize ferrous ions to ferric ions at the anode: $Fe^{2+} (aq) \rightarrow Fe^{3+} (aq) + e^-$ It is also possible to reduce ferricyanide ions to ferrocyanide ions at the cathode: $Fe(CN)^{3-}_6 + e^- \rightarrow Fe(CN)^{4-}_6$ Neutral molecules can also react at either electrode.
• The reaction at this electron is: $Cu^{2+} (aq) + 2e^- \rightarrow Cu (s)$ At the positive anode, copper metal is oxidized to form Cu2+ ions.
• ### Free Energy and Cell Potential

• Free Energy and Cell PotentialThe basis for an electrochemical cell, such as the galvanic cell, is always a redox reaction that can be broken down into two half-reactions: oxidation at anode (loss of electron) and reduction at cathode (gain of electron).
•  As such, the following rules apply:If E°cell > 0, then the process is spontaneous (galvanic cell)If E°cell < 0, then the process is nonspontaneous (electrolytic cell)Therefore, in order to have a spontaneous reaction, E°cell must be positive, where:E°cell = E°cathode − E°anodewhere E°anode is the standard potential at the anode and E°cathode is the standard potential at the cathode as given in the table of standard electrode potentials.
• ### Preventing Corrosion

• Sacrificial Anode Protection Using the same principle as sacrificial film coating, a sacrificial anode, made of a metal more active than the metal you are trying to protect, can be used to prevent corrosion on submerged or buried metal structures.
• The sacrificial anode will corrode before the metal it is protecting does.
• However, once the sacrificial anode corrodes, it must be replaced; otherwise, the metal it is protecting will begin to corrode as well.
• ### Electrolysis of Water

• Electrolysis of a solution of sulfuric acid or of a salt such as NaNO3 results in the decomposition of water at both electrodes:cathode: ${H}_{2}O + 2 {e}^{–} \rightarrow {H}_{2}(g) + 2 {OH}^{–}$E =+0.41 v ([OH–] = 10-7 M)anode: $2 {H}_{2}O \rightarrow {O}_{2}(g) + 4 {H}^{+} + 2 {e}^{–}$E° = -0.82 vnet:  $2 {H}_{2}O(l) \rightarrow 2 {H}_{2}(g) + {O}_{2}(g)$ E = -1.23 vHydrogen will appear at the cathode (the negatively charged electrode, where electrons enter the water) and oxygen will appear at the anode (the positively charged electrode).