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Metals tend to be oxidized, not reduced. Oxidation happens at the anode (positive charged one during electrolysis, it wants electrons), thus the cathode is protected look into cathodic protection.

The long answer:

Metals tend to become positive ions thus they are easily oxidized not reduced. Oxidation happens at the anode, reduction at the cathode. During electrolysis the anode is connected to the positive terminal thus it wants electrons, though a lot of the electrons will come from the electrolyte Cl$^-$ and H$_2$O some electrons are also stripped from the metal electrode causing it to rust/corrode. We know that thermodynamically a reaction is spontaneous if its Gibbs free energy is negative. In electrochemistry the Gibbs free energy is given by the Nernst equation $$ \Delta G=-nfE $$ so in electrolysis we need to supply a voltage (E) just as high as those given in the standard reduction potential table to cause electrolysis/decomposition. We can tell which reactions are occurring using the standard reduction potential table and comparing them to the supplied voltage. Though we must consider kinetics as well not just thermodynamics for the real solution.

Metals tend to be oxidized, not reduced. Oxidation happens at the anode (positive charged one during electrolysis, it wants electrons), thus the cathode is protected look into cathodic protection.

The long answer:

Metals tend to become positive ions thus they are easily oxidized not reduced. Oxidation happens at the anode, reduction at the cathode. During electrolysis the anode is connected to the positive terminal thus it wants electrons, though a lot of the electrons will come from the electrolyte Cl$^-$ and H$_2$O some electrons are also stripped from the metal electrode causing it to rust/corrode. We know that thermodynamically a reaction is spontaneous if its Gibbs free energy is negative. In electrochemistry the Gibbs free energy is given by the Nernst equation $$ \Delta G=-nfE $$ so in electrolysis we need to supply a voltage (E) just as atleast as high as those calculated from the Nernst equation theoretically to cause electrolysis/decomposition. We can tell which reactions are occurring using the standard reduction potential table and comparing them to the supplied voltage. Though we must consider kinetics as well not just thermodynamics for the real solution.

Metals tend to be oxidized, not reduced. Oxidation happens at the anode (positive charged one, it wants electrons), thus the cathode is protected look into cathodic protection.

The long answer:

Metals tend to become positive ions thus they are easily oxidized not reduced. Oxidation happens at the anode, reduction at the cathode. During electrolysis the anode is connected to the positive terminal thus it wants electrons, though a lot of the electrons will come from the electrolyte Cl$^-$ and H$_2$O some electrons are also stripped from the metal electrode causing it to rust/corrode. We know that thermodynamically a reaction is spontaneous if its Gibbs free energy is negative. In electrochemistry the Gibbs free energy is given by the Nernst equation $$ \Delta G=-nfE $$ so in electrolysis we need to supply a voltage (E) just as high as those given in the standard reduction potential table to cause electrolysis/decomposition. We can tell which reactions are occurring using the standard reduction potential table and comparing them to the supplied voltage. Though we must consider kinetics as well not just thermodynamics for the real solution.

Metals tend to be oxidized, not reduced. Oxidation happens at the anode (positive charged one during electrolysis, it wants electrons), thus the cathode is protected look into cathodic protection.

The long answer:

Metals tend to become positive ions thus they are easily oxidized not reduced. Oxidation happens at the anode, reduction at the cathode. During electrolysis the anode is connected to the positive terminal thus it wants electrons, though a lot of the electrons will come from the electrolyte Cl$^-$ and H$_2$O some electrons are also stripped from the metal electrode causing it to rust/corrode. We know that thermodynamically a reaction is spontaneous if its Gibbs free energy is negative. In electrochemistry the Gibbs free energy is given by the Nernst equation $$ \Delta G=-nfE $$ so in electrolysis we need to supply a voltage (E) just as high as those given in the standard reduction potential table to cause electrolysis/decomposition. We can tell which reactions are occurring using the standard reduction potential table and comparing them to the supplied voltage. Though we must consider kinetics as well not just thermodynamics for the real solution.

metals tend to be oxidised not reduced. oxidation happens at the anode (positive charged one, it wants electrons) thus the cathode is protected look into cathodic protection

the long answer: Metals tend to become positive ions thus they are easily oxidised not reduced. Oxidation happens at the anode, reduction at the cathode. During electrolysis the anode is connected to the positive terminal thus it wants electrons, though alot of the electrons will come from the electrolyte Cl- and H2O some electrons are also stripped from the metal electrode causing it to rust/corrode. We know that thermodynamically a reaction is spontaneous if its gibbs free energy is negative. In electrochemistry the Gibbs free energy is given by the Nernst equation $$ \Delta G=-nfE $$ so in electrolysis we need to supply a voltage (E) just as high as those given in the standard reduction potential table to cause electrolysis/decomposition. We can tell which reactions are occurring using the standard reduction potential table and comparing them to the supplied voltage. Though we must consider kinetics as well not just thermodynamics for the real solution

Metals tend to be oxidized, not reduced. Oxidation happens at the anode (positive charged one, it wants electrons), thus the cathode is protected look into cathodic protection.

The long answer:

Metals tend to become positive ions thus they are easily oxidized not reduced. Oxidation happens at the anode, reduction at the cathode. During electrolysis the anode is connected to the positive terminal thus it wants electrons, though a lot of the electrons will come from the electrolyte Cl$^-$ and H$_2$O some electrons are also stripped from the metal electrode causing it to rust/corrode. We know that thermodynamically a reaction is spontaneous if its Gibbs free energy is negative. In electrochemistry the Gibbs free energy is given by the Nernst equation $$ \Delta G=-nfE $$ so in electrolysis we need to supply a voltage (E) just as high as those given in the standard reduction potential table to cause electrolysis/decomposition. We can tell which reactions are occurring using the standard reduction potential table and comparing them to the supplied voltage. Though we must consider kinetics as well not just thermodynamics for the real solution.