Hello and welcome to this video about oxidation-reduction reactions (respredox reactions, abbreviations).
These reactions include theelectron transferamong chemical species and are ubiquitous in nature and synthetic chemistry. When you burn propane in your grill, that combustion reaction is a redox reaction. The annoying rust on your car is also caused by a redox reaction. until yourcellsGenerating energy through redox reactions. So you are literally surrounded by them.
Let's start by understanding the terminology. Redox reactions are so named because a chemical species is formed during the reaction.rustywhile another is at the same timereduced🇧🇷 But what does it mean?
if there is somethingrusty, Terlost electronsand when something isreduced, gained electrons. An easy-to-remember mnemonic isOIL-RIG: You lose oxidized, you gain reduced.
Oxidation and reduction reactions are always paired. There is no simple oxidation reaction. When a chemical species is oxidized, it loses electrons and those electrons have to go somewhere, so they reduce another chemical species. Therefore, the oxidized species is also called a reducing agent, as it provides the electrons for reduction. Likewise, the reduced species is called an oxidizing agent because it accepts electrons for oxidation. This terminology can often be confusing, so familiarize yourself with it.
Consider a redox reaction involving a known product, sodium chloride, which can be formed from sodium metal and chlorine gas.
We know that sodium chloride is aionic compound, which means that sodium transferred an electron to chlorine to form more Na and less Cl. Given what we just learned about redox reactions, we can now say that sodium is oxidized because it loses an electron, and since chlorine gains an electron, it is reduced.
In a simple reaction like this, it's easy to follow the transfer of 1 electron. However, for unknown and more complex reactions, it is difficult to follow the flow of electrons. Thus, chemists attribute aoxidation number(known as "ON" or "oxidation state") to any atom, usually reflecting how rich or poor in electrons that atom is. We simplify the process by assuming that all bonds, even covalent ones, are ionic, meaning 100% electron transfer. Because of this assumption, oxidation numbers do not represent the actual charge on an atom. But by tracking the change in oxidation numbers, we can track the flow of electrons in a chemical reaction.
To assign oxidation numbers, we need to learn a set of rules. Let's start by learning what it takes to assign oxidation numbers for the formation of sodium chloride.
Rule 1:The oxidation of each free element is 0. So, in our example, our sodium chlorine metallic gas has an oxidation number of 0.
Rule 2:The oxidation number of a Group 1A element in a compound is +1. This rule applies to sodium in sodium chloride, so we can assign it an oxidation number of +1.
Rule 3:The oxidation number of a Group 7A(17) element in a compound is -1, unless paired with an atom of a higherelectro-negativity,in this case takes a value ofBalancethe compound load. In our case, chlorine is paired with sodium, which has much lower electronegativity, so chlorine has an oxidation number of -1.
Now that we've assigned an oxidation number to each element in our reaction, let's perform a final check to ensure the assignments are correct using the following rule:
Rule 4:The sum of the oxidation numbers of the atoms within a compound must equal the compound's charge. In our example, sodium chloride is the only compound and it has a total charge of 0, which is confirmed by the fact that the oxidation numbers of Na, +1 and Cl, -1 actually add up to 0.
Assigning oxidation numbers to each species, we can see that the sodium atoms were oxidized because their oxidation number was increased, while the chlorine atoms were reduced because their oxidation number was decreased. We can also say that sodium acts as a reducing agent and chlorine acts as a reducing agent.oxidizing agentin this reaction.
Now let's consider a slightly more complex reaction. What happens when saline solution is added to silver nitrate solution? Sodium chloride and silver nitrate are soluble in water, which means they are broken down into solvated ions.
When the two solutions are mixed, the silver and chloride ions collide to form silver chloride, a water-insoluble solid. The overall reaction can be written as sodium chloride, which is soluble in water, plus silver nitrate, which is also soluble in water, react to form sodium nitrate, which is soluble in water, plus silver chloride, which is a solid.
Is this a redox reaction? To answer this question, we need to know whether electrons were transferred between species. To do this, let's assign oxidation numbers.
We can start with sodium chloride since we assigned these oxidation numbers. Sodium is +1 and chlorine is -1. To assign silver nitrate, we need to know this rule:
Rule 5:The oxidation number of oxygen is -2 unless it is bonded to fluorine or other oxygen.
In the nitrate ion, all three oxygens are bonded to nitrogen, so each oxygen has an oxidation number of -2. Chemically, this means that although electrons are shared in the bond between nitrogen and oxygen, they spend more time with oxygen, so we consider them "belonging" to oxygen.
To assign the oxidation number of nitrogen, we apply the fourth rule we discussed earlier. Knowing that the total charge on a nitrate ion is -1 and that the sum of the oxidation numbers must equal the charge on the compound, we can calculate the oxidation state of nitrogen: oxygen plus the oxidation number of nitrogen. Negative one is equal to three times negative 2 plus the oxidation number of nitrogen. The positive five corresponds to the oxidation number of nitrogen.
We now know that the oxidation number of nitrogen is +5.
We then assign silver nitrate to the oxidation number of silver, which brings us to our sixth rule.
Rule 6:The oxidation number of a monatomic ion is the charge on the ion. When silver nitrate dissolves in water, it splits into a silver ion and a nitrate ion. We know that the silver ion has a +1 charge, so silver must also be in the +1 oxidation state.
We've assigned the oxidation number to each element on the educt side, now let's quickly move to the right side. Sodium nitrate is soluble in water, so it will exist as sodium and nitrate ions. Sodium is still in a +1 oxidation state and the nitrogen and oxygen atoms are still at +5 and -2, respectively, in the nitrate ion.
And finally, by Rule 3, the chlorine in silver chlorine is -1, and since the general compound is neutral, silver must be +1.
Note that the oxidation state has not changed for any of the elements. Therefore, no electrons were transferred between the species, so it's not a redox reaction!
This was an example ofdouble displacement reaction, or metathesis reactions that are not redox reactions. In addition to precipitation reactions, this category includes acid-base reactions and takes the general form AB plus CD reacting with AD plus CB.
Therefore, when faced with a reaction in this form, you can quickly identify it as NOT being a redox reaction without assigning oxidation states.
Finally, an example: the decomposition of potassium chlorate. If you like, pause the video and take a second to assign the oxidation states yourself.
Okay, let's try it together now.
According to Rule 1, the oxygen atoms on the product side have an oxidation number of 0.
Using rules 2 and 3, we can easily assign oxidation numbers +1 and -1 to potassium and chlorine in potassium chloride on the product side.
Since there are no exceptions to Rule 2, we can also assign an oxidation state of +1 to potassium in potassium chlorate.
Chlorine is bonded to oxygen, which is the most electronegative atom, which means that the oxidation number of Cl is not -1, and instead we have to solve for it. Knowing that the oxidation state of oxygen is -2, we can establish the following equation: The total charge on potassium chlorate is equal to the oxidation number of potassium plus the oxidation number of chlorine plus three times the oxidation number of oxygen. Zero equals a positive one plus the oxidation number of chlorine plus three times negative 2. A positive 5 equals the oxidation number of chlorine.
In potassium chlorate, chlorine has an oxidation number of +5.
Now let's consider the general motion of electrons.
The oxidation number of chlorine dropped from +5 to -1, so chlorine gained electrons and was thus reduced. The oxidation number of oxygen increased from -2 to 0, so it lost electrons and became oxidized. Note that in a redox decomposition reaction, the reactant is both the oxidizing and reducing agent.
Potassium chloride is described as a strong oxidizing agent, which means it can be easily reduced. We can understand this by evaluating the oxidation numbers. Chlorine, a relatively electronegative element, has an oxidation state of +5 because it is bonded to 3 oxygen atoms. Consequently, potassium chlorate readily accepts electrons to reduce chlorine.
Elements on the left side of the periodic table, such as metals in Groups 1 and 2, are generally easily oxidized because they have low electronegativity. In contrast, the elements on the right, such as the oxygen family and the halogens, are easily reduced because they have high electronegativity.
But, as you've probably noticed, there are often exceptions to the rules.
For example, hydrogen in compounds usually has a +1 oxidation state. In water, H2O, hydrochloric acid, HCl and methane, CH4, and basically all other organic compounds, hydrogen has an oxidation number of +1. This is because, in these cases, the hydrogen is bonded to a more electronegative atom. However, hydrogen can form compounds with less electronegative atoms, in which case its oxidation state is -1! For example, in sodium hydride, NaH, electrons spend more time near hydrogen, so we assign hydrogen an oxidation state of -1 and sodium +1. Although these compounds are rare, they do show up in synthetic chemistry, so it's good to know about them.
In summary, the best approach to solving redox reactions is to learn the rules for assigning oxidation states, but always stop and consider whether the assignments make chemical sense. Remember that the oxidation state generally should reflect whether the atom is electron-rich or electron-poor; so check each chemical bond and consider where the electrons would spend most of their time. This is to help you avoid mistakes when working with oxidation states and redox reactions.
Thanks for looking and happy learning!