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In this video, we are going to identify the oxidizing agent and the reducing agent in two different chemical equations.
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So basically, the oxidizing agent is the species that's reduced, and the reducing agent is the species that's oxidized.
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And to determine what's oxidizing reduced, we've got to look at the oxidation numbers.
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So we start with two elements in their natural state.
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That means that they have oxidation numbers of zero.
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And then we make a compound.
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Oxygen is going to be negative two.
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So this is a total of three times negative two is negative six.
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And then we have, so this whole thing has to be positive six.
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And we have two of them, so that's going to be plus three for the iron that's equal to its charge.
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So which one goes up in oxidation number? iron goes from zero to plus three.
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So that's what's going to be oxidized, meaning it's the reducing agent.
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So f .e is our reducing agent.
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And then o2 is going to be our oxidizing agent because it is reduced from zero.
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To negative 2.
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And we also need to say the change in oxidation number.
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So fe goes from 0 to plus 3, and o2 goes from 0 to negative 2.
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All right.
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Now we can do the second equation.
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So we want the oxidizing agent, the species that's reduced, and then the reducing agent, the species that is oxidized.
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Again, we're going to find the oxidation numbers.
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Okay, so let's see.
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Oxygen is negative 2 to get an overall charge of negative 1, this chlorine has to be plus 1.
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Then we have oxygen is negative 2, hydrogen is plus 1...