why water can't react as anode , with metals , in a redox reaction? while studying electrochemistry, i passed through a table with certain metals , all having voltage (potential energy) less than water, which can react with water in redox reaction. in this case, the water react as oxidation agent. my question, is why water cant react as reducing agent, with metals having potential energy bigger that it?

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    $\begingroup$ Water surely can and does react, given the right conditions. Could you please be more specific? $\endgroup$ Nov 16 '17 at 15:07
  • $\begingroup$ i will state from my lecture slideshow: metals will react spontaneously with water if their E is less than -0.83V my question is why ? why a metal which has E>-0.83 V can't react with water? $\endgroup$ Nov 16 '17 at 15:19

Water can react as a reducing agent. However, you would have to look at the oxidation equation for water, which has an electrode potential of 1.23 V. If you wanted water to act as a reducing agent, you would have to react it with a metal whose E > 1.23 V. The equation that you are referring to (where E = -0.83 V) is the reduction of water. Therefore, you can only use it when you want water to act as an oxidising agent.

A note about spontaneity; it only refers to whether a reaction is likely to occur without any addition or removal of energy. In other words, the reaction would occur if the chemicals were brought in contact with each other. It does not mean that the reaction is impossible, because if you provided sufficient energy, the chemicals would react. This is usually supplied via a change in temperature. The reason metals whose E < -0.83 V react spontaneously with water is because the nature of the Gibbs free energy change (ΔG). For redox reactions, ΔG = -n x F x E cell where n is the moles of electrons transferred in the reaction, F is Faraday's constant and E cell is the cell potential. Reactions are spontaneous if ΔG is negative, meaning that E cell must be positive for the reaction to be spontaneous. When you add the half cell equations for metals whose E < -0.83 V and that of water, you get a positive value for E cell. Additionally, redox reactions are based on the idea that more reactive metals displace less reactive ones from compounds. Metals above water are reactive enough to displace hydrogen from water. Metals below water are not reactive enough to displace hydrogen and so cannot react spontaneously.


Water cannot act as a reducing agent to metals, simply because of the fact that metals do not react as oxidants, however metal-ions do. Below the values of the data of the standard electrode potentials (and the source site -- Wikipedia) is shown. Water is highlighted. Standard Electrode Potentials
Data of standard electrode potentials
The table is sorted so that the strongest oxidants are placed highest in the table. As you can see, above water as a reductant, there are no metals reacting in half-reactions as oxidants. There are some metal-ions that would react as an oxidant. The reason behind why metals don't act as an oxidant, is probably because of their low electronegativity. The high electronegativity of some elements prevents them from becoming positively charged, and in a similar fashion, the low electronegativity of a metal prevents it from becoming negatively charged.


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