The equation below shows an obvious example of oxygen transfer in a simple redox reaction:. Copper II oxide and magnesium oxide are both ionic compounds. If the above is written as an ionic equation, it becomes apparent that the oxide ions are spectator ions. Omitting them gives:. In the above reaction, magnesium reduces the copper II ion by transferring electrons to the ion and neutralizing its charge.
Therefore, magnesium is a reducing agent. Another way of putting this is that the copper II ion is removing electrons from the magnesium to create a magnesium ion. The copper II ion is acting as an oxidizing agent.
Confusion can result from trying to learn both the definitions of oxidation and reduction in terms of electron transfer and the definitions of oxidizing and reducing agents in the same terms. The following thought pattern can be helpful:. Jim Clark Chemguide. Reaction with hydrogen also came to be regarded as reduction.
Later, a more general idea of oxidation and reduction was developed in which oxidation was loss of electrons and reduction was gain of electrons. This wider definition covered the original one and also applies to reactions that do not involve oxygen.
However, it applies only to reactions in which electron transfer occurs — i. It can be extended to reactions between covalent compounds by using the concept of oxidation number or state. This is a measure of the electron control that an atom has in a compound compared to the atom in the pure element.
An oxidation number consists of two parts: 1 its sign, which indicates whether the control has increased negative or decreased positive ; 2 its value, which gives the number of electrons over which control has changed. The change of electron control may be complete in ionic compounds or partial in covalent compounds. Oxidation is a reaction involving an increase in oxidation number and reduction involves a decrease. I want to just assume hypothetically, what if these were ionic bonds?
And you say, well, if these had to be ionic bonds, then the oxygen would nab the electrons from these pairs. And so the oxygen would have a fully negative charge, a negative 2 charge. And the hydrogens would have a fully positive charge each. And so, if we were to write down the oxidation states for the atoms in the water molecule-- let's write that down, so H2O-- we would say that oxygen has an oxidation state of negative 2, and each hydrogen atom has an oxidation state of plus 1.
And notice, the whole molecule is neutral, and these things cancel out with each other. Positive 1, positive 1, that gets you to positive 2. Then you have negative 2. They cancel out. Now, the one thing, I keep saying this is negative 2, but I wrote the negative after it. If I wanted to write positive 1 as an oxidation state, I would actually write it as 1 positive, although you can assume that if someone just writes the positive.
And this is just the convention, to write the sign after the number when we are writing actually ionic charges or oxidation states, because an oxidation state is nothing but a hypothetical ionic charge. If you really had to-- if you were forced to assume these aren't covalent bonds, but these are ionic bonds.
Once again, I want to stress. This is the reality. These are partial charges, the oxidation state, intellectual tool, that's forcing us to pretend like these are ionic bonds. And you might say well, this kind of makes sense right over here. This involved oxygen in some way.
That's why it's called oxidation states. And that's how I initially conceptualized it when I first learned about this. You say, well, look, each of these hydrogens lost an electron to oxygen.
So it makes sense that we say that each hydrogen got oxidized, so hydrogen oxidized by oxygen. It makes sense that oxygen would oxidize something else. This got done to it. The charge was taken away by oxygen, so it got oxidized. Now, the other term on the other side of oxidized is reduced. And the word "reduced" really comes from the idea that oxygen's charge has been reduced. So we could say, O, or we could say oxygen has been reduced by the hydrogens.
And so there there's a temptation here to say, well, OK, this must always involve oxygen in some kind, because it seems to begin with the same words. Well, that is not the case. Let's take, for example, if this is an aqueous solution, hydrofluoric acid right over here. You have a hydrogen covalently bonded to a fluorine. Now, just like we saw in water, fluorine is one of the most electronegative elements.
It's going to hog the electrons in this covalent bond. Combination reactions are among the simplest redox reactions and, as the name suggests, involves "combining" elements to form a chemical compound.
As usual, oxidation and reduction occur together. The general equation for a combination reaction is given below:. In this reaction both H 2 and O 2 are free elements; following Rule 1 , their oxidation states are 0.
The product is H 2 O, which has a total oxidation state of 0. According to Rule 6 , the oxidation state of oxygen is usually A decomposition reaction is the reverse of a combination reaction, the breakdown of a chemical compound into individual elements:. This follows the definition of the decomposition reaction, where water is "decomposed" into hydrogen and oxygen. Note that the autoionization reaction of water is not a redox nor decomposition reaction since the oxidation states do not change for any element:.
A single replacement reaction involves the "replacing" of an element in the reactants with another element in the products:. A double replacement reaction is similar to a single replacement reaction, but involves "replacing" two elements in the reactants, with two in the products:.
An example of a double replacement reaction is the reaction of magnesium sulfate with sodium oxalate. Combustion is the formal terms for "burning" and typically involves a substance reacts with oxygen to transfer energy to the surroundings as light and heat.
Hence, combustion reactions are almost always exothermic.
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