Question: How many electrons are lost by sodium metal in a redox reaction?
About the article "Chemical reaction", section "Oxidation and reduction".

Candidate passages:
Redox reactions can be understood in terms of transfer of electrons from one involved species (reducing agent) to another (oxidizing agent).
In this process, the former species is "oxidized" and the latter is "reduced".
Though sufficient for many purposes, these descriptions are not precisely correct.
Oxidation is better defined as an increase in oxidation state, and reduction as a decrease in oxidation state.
In practice, the transfer of electrons will always change the oxidation state, but there are many reactions that are classed as "redox" even though no electron transfer occurs (such as those involving covalent bonds).
In the following redox reaction, hazardous sodium metal reacts with toxic chlorine gas to form the ionic compound sodium chloride, or common table salt:
formula_21
In the reaction, sodium metal goes from an oxidation state of 0 (as it is a pure element) to +1: in other words, the sodium lost one electron and is said to have been oxidized.
On the other hand, the chlorine gas goes from an oxidation of 0 (it is also a pure element) to -1: the chlorine gains one electron and is said to have been reduced.
Because the chlorine is the one reduced, it is considered the electron acceptor, or in other words, induces oxidation in the sodium - thus the chlorine gas is considered the oxidizing agent.
Conversely, the sodium is oxidized or is the electron donor, and thus induces reduction in the other species and is considered the "reducing agent".
Which of the involved reactants would be reducing or oxidizing agent can be predicted from the electronegativity of their elements.
Elements with low electronegativity, such as most metals, easily donate electrons and oxidize – they are reducing agents.
On the contrary, many ions with high oxidation numbers, such as , , , , can gain one or two extra electrons and are strong oxidizing agents.
The number of electrons donated or accepted in a redox reaction can be predicted from the electron configuration of the reactant element.
Elements try to reach the low-energy noble gas configuration, and therefore alkali metals and halogens will donate and accept one electron respectively.
Noble gases themselves are chemically inactive.
An important class of redox reactions are the electrochemical reactions, where electrons from the power supply are used as the reducing agent.
These reactions are particularly important for the production of chemical elements, such as chlorine or aluminium.
The reverse process in which electrons are released in redox reactions and can be used as electrical energy is possible and used in batteries.
In redox processes, the reductant transfers electrons to the oxidant.
Thus, in the reaction, the reductant or "reducing agent" loses electrons and is oxidized, and the oxidant or "oxidizing agent" gains electrons and is reduced.
The pair of an oxidizing and reducing agent that are involved in a particular reaction is called a redox pair.
A redox couple is a reducing species and its corresponding oxidized form, e.g., Fe2+/Fe3+. 
Redox (reduction-oxidation) reactions include all chemical reactions in which atoms have their oxidation state changed; in general, redox reactions involve the transfer of electrons between species.
This can be either a simple redox process, such as the oxidation of carbon to yield carbon dioxide () or the reduction of carbon by hydrogen to yield methane (CH4), or a complex process such as the oxidation of glucose (C6H12O6) in the human body through a series of complex electron transfer processes.
The term "redox" comes from two concepts involved with electron transfer: reduction and oxidation.
It can be explained in simple terms:  Oxidation is the "loss" of electrons or an "increase" in oxidation state by a molecule, atom, or ion.
Reduction is the "gain" of electrons or a "decrease" in oxidation state by a molecule, atom, or ion.
Although oxidation reactions are commonly associated with the formation of oxides from oxygen molecules, these are only specific examples of a more general concept of reactions involving electron transfer.
Redox reactions, or oxidation-reduction reactions, have a number of similarities to acid–base reactions.
Like acid–base reactions, redox reactions are a matched set, that is, there cannot be an oxidation reaction without a reduction reaction happening simultaneously.
The oxidation alone and the reduction alone are each called a "half-reaction", because two half-reactions always occur together to form a whole reaction.
When writing half-reactions, the gained or lost electrons are typically included explicitly in order that the half-reaction be balanced with respect to electric charge.
Though sufficient for many purposes, these descriptions are not precisely correct.
Oxidation and reduction properly refer to "a change in oxidation state" — the actual transfer of electrons may never occur.
Thus, oxidation is better defined as an "increase in oxidation state", and reduction as a "decrease in oxidation state".
In practice, the transfer of electrons will always cause a change in oxidation state, but there are many reactions that are classed as "redox" even though no electron transfer occurs (such as those involving covalent bonds). 
Sodium metal is a single-electron reducing agent, meaning the sodium metal will transfer electrons one at a time.
Four sodium atoms are required to fully reduce each ester to alcohols.
Ethanol serves as a proton source. 
In acidic media, ions and water are added to half-reactions to balance the overall reaction.
For instance, when manganese(II) reacts with sodium bismuthate:The reaction is balanced by scaling the two half-cell reactions to involve the same number of electrons (multiplying the oxidation reaction by the number of electrons in the reduction step and vice versa): 8 (l) + 2 (aq) → 2 (aq) + 16 (aq) + 10  10 + 30 + 5 (s) → 5 (aq) + 15 (l) Adding these two reactions eliminates the electrons terms and yields the balanced reaction: 14 (aq) + 2 (aq) + 5 (s) → 7 (l) + 2 (aq) + 5 (aq) + 5 (aq)