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Reducing agents are the element(s) that are oxidized (oxidation state increases) and oxidizing agents are the element(s) that are reduced oxidation state decreases). You MUST have BOTH for a redox reaction to occur! So, look for the two (or more) elements in an equation whose reaction change requires a CHANGE in their oxidation state. If no change is found, the reaction is not a redox reaction ...
You rank oxidizing agents according to their standard reduction potentials. > Here's a typical table of standard reduction potentials. (From wps.prenhall.com) The species at the top left have the greatest "potential" to be reduced, so they are the strongest oxidizing agents. The strongest oxidizing agent in the list is "F"_2, followed by "H"_2"O"_2, and so on down to the weakest oxidizing ...
Oxidation agent: Something that is "reduced" Reduction agent: Something that is "oxidized" Whenever you see "oxidation agent", think of reduction. An oxidation agent is reduced, and gains electrons in a reaction. The oxidation number will therefore go down. A reduction agent is oxidized in a reaction. It loses electrons and the oxidation number goes up.
The species whose oxidation number is increased is the reducing agent, and the species whose oxidation number is decreased in the oxidizing agent. Let's take a simple example of combustion, which is formally a redox reaction: C(s) + O_2(g) rarr CO_2(g) Carbon is oxidized from elemental carbon, the zerovalent, elemental state (oxidation state = 0), to C(IV), its maximum oxidation state ...
Balancing a redox reaction requires identifying the oxidation numbers in the net ionic equation, breaking the equation into half reactions, adding the electrons, balancing the charges with the addition of hydrogen or hydroxide ions, and then completing the equation.
What are the five strongest reducing and oxidizing agents? How do you identify reducing agents, oxidizing agents, or whether neither of these two apply? In the reaction: #2H_2(g) + O_2(g) -> 2H_2O(g)#, what is the oxidizing agent?
The oxidizing agent is "I"_2. A quick technique to use here would be to look at the fact that you're going from iodine, "I"_2, on the reactants' side to the iodide anion, "I"^(-), on the products' side. In this case, you're going from a neutral molecule to a negatively charged ion, so right from the start, you know that iodine is being reduced, i.e. it is taking in electrons. This can only ...
Look at the electronegativity values. Oxidation is the removal of electrons from an atom or polyatomic ion. The higher the electronegativity the greater the pull an oxidizing agent has for electrons. The higher the pull for electrons the stronger the oxidizing agent. So the element with the highest electronegativity is the strongest oxidizing agent.
An oxidizing agent is reduced by the oxidized ion. For instance, Pb(s)+H_2O(l)+O_2(g) to Pb(OH)_2(s) + H_2(g) I know, it's not balanced! Pb has an oxidation state of 0, H has one of +2 in water. After the reaction, lead is oxidized to a state of +2, losing electrons, while hydrogen turns into gas, gaining electrons and having a state of 0. So, since Pb is oxidized, H is the oxidizing agent ...
If the oxidation number increases upon reaction, the species is a reducing agent. And if the oxidation number decreases, the species is an oxidizing agent. Redox transfer is formalized on the basis of loss or gain or electrons. Recall the old mnemonic: "LEO SAYS GER", "loss of electrons = oxidation; gain of electrons = reduction." And thus if something has been oxidized, it is A SOURCE of ...