Oxidizing and Reducing Agents — Other Ionic Reactions

by Adam Le Gresley, PhD

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    00:00 The species that is reduced itself and causes another species to be oxidized is therefore known as an oxidizing agent. So in our previous example, we could say that in this particular case, ion is reducing copper and therefore could be considered a reducing agent since if itself is oxidized and therefore similarly the species that is oxidized and causes another to be reduced is known as the reducing agent. Oxidation number. Now, if you recall back in the previous module, we discussed the idea of group 1 and group 2 metals losing electrons to form a stable outer shell. We also talked about group 7 elements gaining electrons to form a full outer shell.

    00:46 And I indicated at the time that for the case of group 1 and group 2 and group 3, now only 2 possible oxidation states. That is to say either it hadn't given up any electrons or it had given up either 1 or 2 or 3, but that will be it. However, when it comes to the transition metal or deep block elements, you see that the possible number of oxidation numbers is wide and varied. The concept of oxidation numbers is therefore very important way of keeping track of electrons in a reaction. When it's possible for a single element to have multiple oxidation states. The oxidation number, ON, or oxidation states of an atom in a substance is the actual charge of the atom if it's existed in a monoatomic form. Alternatively, it is the hypothetical charge aside to the atom in a substance using simple rules. So let's look at some of those rules. Rule 1, the oxidation number of any atom in an element is 0. So if we were to look at, for example, a single atom of fluorine or a single atom of sodium, the oxidation number of that will be 0. ___ able to look at a single atom of copper, the oxidation state would also be 0. When you're looking at monoatomic ions, the oxidation of an atom in a monoatomic ion equals the charge of the ion. In the case of oxygen, the oxidation number of oxygen is typically -2. The only exception to this being in peroxides is ___ when you can have exceptionally a -1 oxidation state. But in organic chemistry, which is what we're focusing on to this lecture, the only principal charge of oxygen is -2. In the case of hydrogen, the oxidation number of hydrogen is +1 and most of its compounds, the exception being in hydrites such as in the case of calcium hydrite where the electrons are transferred from the calcium formally to the hydrogen thus filling the outer shell of the hydrogen and forming a complete outer shell in the case of calcium. In the case of the halogens, typically you will encounter them in ionic chemistry as being -1. This would be for chlorine, bromine, or iodine thus forming the chloride, bromide, and iodide ions. There are some exceptions to this and one of them you've come across I'm I'm sure, no doubt, which is where you have reaction of those with oxygen. You may say you haven't come across it but if you've ever used bleach which is a solution of sodium hypochloride or chloride I which is where oxygen reacts with chlorine and actually deprives it of electrons thus actually creating a chlorine and its +1 oxidation state. Compounds in irons which is rule #6 is the sum of the oxidation numbers of the atom in a compound must be zero, i.e. the electrons must go somewhere and that the sum of the oxidation numbers of the atom and a polyatomic ion equals the charge on the ion itself. It's possible to predict the upper and lower limits of main group elements. The upper limit is equal to the group number so if for example we were to take sodium, the upper limit must be 1. If we were to take magnesium which is in group 2, the upper limit must be 2 and so on and so forth. The lower limit, I would agree to which can be reduced, is the group number minus 8. Obviously because if you're talking about the shell, that is the maximum that's allowable within the shell #2. Oxygen will never have an oxidation number of +6 and fluorine will never have an oxidation number of +7. That would imply to removing all of its electrons. So let's go back to our friend ion and copper. In this case, the ion moves its 2 electrons which is oxidized and in the same way it is acting as a reducing agent causing the copper to gain 2 electrons. Right. So how do we write redox reactions? These are probably in terms of knowing where the electrons are the more complex types of reactions you will encounter within the ionic sphere of chemistry. There are 2 ways to deal with redox reactions from a bouncing perspective. Either you treat them as any other reaction or you can write them in terms of 2 half reactions as we did for the copper and for the iron. The driving force for these reactions is in exchange of electrons, i.e. where one element has a greater or ion has a greater affinity than another element. A half reaction is one of two parts of an oxidation reduction or redox reaction and it involves the oxidation or loss of electrons and the other step involves the gate of them as we saw before. So if we look at the half reaction from our previous reaction of ion with copper sulfate, we'll see that ion as a solid is converted to ion 2+ in a clear solution and 2 electrons denoted to a minus a lost in the process. Those 2 electrons are then picked up by the Cu2+, which is in solution, forming our elemental copper as a solid and it is indeed the half reactions which are the basis of power generation in batteries where you have 1 element of one end which has a greater affinity for electrons than another element of another end. The movement of charge through a wire is obviously what is therefore the current that you pick up in your battery.

    About the Lecture

    The lecture Oxidizing and Reducing Agents — Other Ionic Reactions by Adam Le Gresley, PhD is from the course Ionic Chemistry.

    Included Quiz Questions

    1. 1.81 g
    2. 0.181 g
    3. 5.02 g
    4. 5.28 g
    5. 3.62 g
    1. During a chemical reaction, an oxidant gets oxidized while the reductant gets reduced under the influence of metal catalysts.
    2. An oxidizing agent (oxidant) gets reduced during a reaction and simultaneously causes oxidation of another compound.
    3. A reducing agent (reductant) gets oxidized and leads to the reduction of another compound participating in the chemical reaction.
    4. A reductant loses its electrons to the oxidant during a redox reaction.
    5. An oxidant acts as an electron acceptor because it gains electrons from an electron donor, or reducing agent, during a chemical reaction.
    1. Fe
    2. O2
    3. F2
    4. Br2
    5. O3
    1. The Daniell cell is an unusual type of electrochemical cell because it does not involve a redox reaction.
    2. Halogens act as good oxidizing agents because they try to obtain noble gas-like configurations by gaining electrons.
    3. H2, CO, Zn, Fe, Na, and Li are good reducing agents.
    4. Alkali and alkaline earth metals behave as reductants by losing electrons to achieve noble gas-like configurations.
    5. Fundamental biological processes, like metabolism and photosynthesis, involve both reducing and oxidizing agents to harvest energy.
    1. …of an oxidizing agent decreases while that of a reducing agent increases.
    2. …of an oxidizing agent increases while that of a reducing agent decreases.
    3. …of an oxidizing agent remains unchanged while that of a reducing agent can either decrease or increase.
    4. …of a reducing agent remains unchanged while that of an oxidizing agent can either decrease or increase.
    5. …of both the reducing agent and the oxidizing agent remain unaffected.
    1. …due to the presence of stable states for d-orbitals.
    2. …due to the presence of unstable states in the d-shell.
    3. …due to a more stable 4s-subshell.
    4. …due to high ionization energies.
    5. …because of the presence of paired electrons in all d-orbitals of the d-shell.
    1. In a water molecule, the algebraic sum of oxidation numbers of all the atoms can be +1.
    2. A free element has an oxidation number of zero.
    3. In the case of a monoatomic ion, the oxidation number is equal to the net charge on the ion.
    4. In a neutral molecule, the algebraic sum of oxidation numbers of all the atoms must be zero.
    5. Hydrogen, fluorine, and oxygen usually exhibit oxidation numbers +1, -1, and -2, respectively in most of the compounds.
    1. In peroxides, the oxygen molecule has an oxidation state of -3.
    2. Hydrogen exhibits an oxidation state of +1 with non-metals, but in metal hydrides of Group 1 metals like NaH or KH, it exists in a -1 oxidation state.
    3. In the compound F2O, the oxygen has an oxidation state of +2 due to the more electronegative character of fluorine.
    4. Chlorine in compounds with fluorine or oxygen exhibits variable oxidation states ranging from +1 to +7.
    5. The oxidation state of sulfur in the sulfate ion (SO42-) is +6, whereas in sulfite (SO32-) the oxidation state is +4.
    1. The upper limit is equal to the group number, whereas the lower limit is given by the formula: Lower limit = (Group number - 8).
    2. The upper limit of the oxidation state can be calculated by the formula: Upper limit = (Group number – 8).
    3. The lower limit of the oxidation state is equal to the group number.
    4. It is not easy to predict the upper limit of the oxidation state of main group elements.
    5. It is not easy to predict the lower limit of the oxidation state of main group elements
    1. …+6 and +7, respectively.
    2. …-2 and 0, respectively.
    3. …-1 and 0, respectively.
    4. …-2 and -1, respectively.
    5. …-1 and -1, respectively.
    1. …small and highly electronegative elements like oxygen and fluorine.
    2. …small and highly electropositive elements like sodium and lithium.
    3. …large and highly electropositive elements like cesium and francium.
    4. …helium gas at 80 °C.
    5. …d-block elements.
    1. …the exchange of electrons.
    2. …the exchange of neutrons.
    3. …the exchange of protons.
    4. …the exchange of photons.
    5. …the exchange of alpha particles.

    Author of lecture Oxidizing and Reducing Agents — Other Ionic Reactions

     Adam Le Gresley, PhD

    Adam Le Gresley, PhD

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