Question: Element that easily donate electrons and oxidize have what property?
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.
Conjugated systems not only have low energy excitations in the visible spectral region but they also accept or donate electrons easily.
Phthalocyanines, which, like Phthalo Blue and Phthalo Green, often contain a transition metal ion, exchange an electron with the complexed transition metal ion that easily changes its oxidation state.
Pigments and dyes like these are charge-transfer complexes. 
Electronegativity, symbol χ, is a chemical property that describes the tendency of an atom or a functional group to attract electrons (or electron density) towards itself.
An atom's electronegativity is affected by both its atomic number and the distance that its valence electrons reside from the charged nucleus.
The higher the associated electronegativity number, the more an element or compound attracts electrons towards it.
First proposed by Linus Pauling in 1932 as a development of valence bond theory, it has been shown to correlate with a number of other chemical properties.
Electronegativity cannot be directly measured and must be calculated from other atomic or molecular properties.
Several methods of calculation have been proposed, and although there may be small differences in the numerical values of the electronegativity, all methods show the same periodic trends between elements.
The most commonly used method of calculation is that originally proposed by Linus Pauling.
This gives a dimensionless quantity, commonly referred to as the Pauling scale, on a relative scale running from around 0.7 to 3.98 (hydrogen = 2.20).
When other methods of calculation are used, it is conventional (although not obligatory) to quote the results on a scale that covers the same range of numerical values: this is known as an electronegativity in Pauling units.
Electronegativity, as it is usually calculated, is not strictly a property of an atom, but rather a property of an atom in a molecule.
Properties of a free atom include ionization energy and electron affinity.
It is to be expected that the electronegativity of an element will vary with its chemical environment, but it is usually considered to be a transferable property, that is to say that similar values will be valid in a variety of situations.
On the most basic level, electronegativity is determined by factors like the nuclear charge (the more protons an atom has, the more "pull" it will have on negative electrons) and the number/location of other electrons present in the atomic shells (the more electrons an atom has, the farther from the nucleus the valence electrons will be, and as a result the less positive charge they will experience—both because of their increased distance from the nucleus, and because the other electrons in the lower energy core orbitals will act to shield the valence electrons from the positively charged nucleus).
The opposite of electronegativity is electropositivity: a measure of an element's ability to donate electrons. 
The key terms involved in redox are often confusing to students.
For example, an element that is oxidized loses electrons; however, that element is referred to as the reducing agent.
Likewise, an element that is reduced gains electrons and is referred to as the oxidizing agent.
Acronyms or mnemonics are commonly used to help remember what is happening:  "OIL RIG"—Oxidation Is Loss of electrons, Reduction Is Gain of electrons.
"LEO the lion says GER" — Loss of Electrons is Oxidation, Gain of Electrons is Reduction.
"LEORA says GEROA" — Loss of Electrons is Oxidation (Reducing Agent) and Gain of Electrons is Reduced (Oxidizing Agent).
"RED CAT" and "AN OX" — Reduction occurs at the Cathode and the Anode is for Oxidation.  
In energy terms, natural photosynthesis can be divided in three steps:  Light-harvesting complexes in bacteria and plants capture photons and transduce them into electrons, injecting them into the photosynthetic chain.
Proton-coupled electron transfer along several cofactors of the photosynthetic chain, causing local, spatial charge separation.
Redox catalysis, which uses the aforementioned transferred electrons to oxidize water to dioxygen and protons; these protons can in some species be utilized for dihydrogen production.
Using biomimetic approaches, artificial photosynthesis tries to construct systems doing the same type of processes.
Ideally, a triad assembly could oxidize water with one catalyst, reduce protons with another and have a photosensitizer molecule to power the whole system.
One of the simplest designs is where the photosensitizer is linked in tandem between a water oxidation catalyst and a hydrogen evolving catalyst:  The photosensitizer transfers electrons to the hydrogen catalyst when hit by light, becoming oxidized in the process.
This drives the water splitting catalyst to donate electrons to the photosensitizer.
In a triad assembly, such a catalyst is often referred to as a donor.
The oxidized donor is able to perform water oxidation.
The state of the triad with one catalyst oxidized on one end and the second one reduced on the other end of the triad is referred to as a charge separation, and is a driving force for further electron transfer, and consequently catalysis, to occur.
The different components may be assembled in diverse ways, such as supramolecular complexes, compartmentalized cells, or linearly, covalently linked molecules.
Research into finding catalysts that can convert water, carbon dioxide, and sunlight to carbohydrates or hydrogen is a current, active field.
By studying the natural oxygen-evolving complex, researchers have developed catalysts such as the "blue dimer" to mimic its function.
Photoelectrochemical cells that reduce carbon dioxide into carbon monoxide (CO), formic acid (HCOOH) and methanol (CH3OH) are under development.
However, these catalysts are still very inefficient.