hello everyone my name is Iman welcome back to my YouTube channel today we're going to be tackling chapter 11 for our MCAT General chemistry playlist this chapter is titled oxidation reduction reactions and we're going to focus on two main objectives here the first objective is titled oxidation reduction reactions we're going to begin by defining oxidation and reduction then we're going to learn how to assign oxidation numbers to various elements in chemical compounds and finally we're going to explore the methods for balancing oxidation reduction reactions then the second objective is titled net ionic equations we're going to start with an overview of net ionic equations and then following this we're going to discuss disproportionation reactions this is a special type of Redux reaction where a single substance underg goes both oxidation and reduction simultaneously lastly we're going to briefly cover oxidation reduction titrations this is a quantitative analytical technique that's used to determine the concentration of an unknown Solution by using a Redux reaction with that introduction let's go ahead and get started with with the first objective reactions that involve the transfer of electrons from one chemical species to another are classified as oxidation reduction reactions also known as Redux reactions at their core these reactions embody the principle of the conservation of charge which states that electrical charge can neither be created nor destroyed this fundamental law means that an electron lost by one species must be gained by another ensuring that oxidation and reduction occur simultaneously now oxidation is defined as the process of losing electrons while reduction is the process of gaining electrons these two processes are intrinsically linked oxidation cannot happen without reduction because the electron that's lost by one substance has to be gained by another now a useful pneumonic to remember these processes is Leo the Lion goes g Leo losing electrons is oxidation G gaining electrons is reduction another pneumonic is oil rig this is my favorite oil oxidation is loss of electrons rig reduction is gain of electrons now in every Redux reaction there are substances known as oxidizing agents and reducing agents the oxidizing agent is the substance that causes another substance to lose electrons and be oxidized in doing so the oxidizing agent itself self is reduced as it gains those electrons on the other hand the reducing agent is the substance that causes another to gain electrons and be reduced and so the reducing agent itself is oxidized as it loses electrons so for example let's consider the reaction between magnesium and hydrochloric acid so this is the balanced chemical equation you have magnesium plus two hydrochloric acid molecules they react to give us magnesium chloride and hydrogen gas in this reaction magnesium loses electrons and is oxidized all right we go from magnesium to magnesium 2+ plus 2 electrons at the same time the hydrogen ions from hydrochloric acid they're going to gain electrons and they become reduced so we go from 2 H+ plus those two electrons they gain to form hydrogen gas so here magnesium acts as a reducing agent because it donates electrons and it undergoes oxidation on the other hand the hydrogen ions they act as oxidizing agents because they accept electrons and they undergo reduction now being familiar being familiar with common oxidizing and reducing agents can significantly streamline your problem solving process especially in organic chemistry reactions so on exams like the MCAT recognizing these agents quickly can be a timesaver so let's delve into the specific provided in this table and understand the characteristics and roles of these agents oxidizing agents are substances that gain electrons in a chemical reaction and in the process they cause another substance to be oxidized these agents are often characterized by the presence of oxygen or another strong electronegative element such as a hallogen here are some common oxidizing agents here we have oxygen hydrogen peroxide the halogens halogens these are highly electronegative elements that readily accept electrons we also have sulfuric acid nitric acid and sodium hypochlorite sodium hypoc chlorate is commonly found actually in household bleach we also have potassium per mangy this is a very versatile oxidizing agent that's used in both the laboratory setting and in Industry you're going to encounter this a lot in organic chemistry we also have chromium trioxide and sodium D chromate these are powerful oxidizing agents used in organic synthesis we also have PCC this is peridinium chlorochromate this is used specifically for the oxidation of alcohols to alahh and ketones and last we have nad+ and fadh these are co-enzymes involved in Redux reactions in biochemical processes they act as electron carriers now reducing agents on the other hand they donate electrons in chemical reactions and they cause another substance to be reduced these these agents typically contain metal ions or hydride here are some common reducing agents here we have carbon monoxide and carbon carbon acts as a reducing agent in the reduction of metal oxides in a lot of high temperature reactions here we have di Boran we also have 102 ion and other pure Metals these are commonly used in Redux reactions we also have hydrazine zinc amalgam and lindlar's Catalyst in addition we have sodium Borah hydde this is widely used in the reduction of carbonal compounds also lithium aluminum hydride this is a very strong reducing agent that's used in the reduction of Esters carboxilic acids and other functional groups and then last we have nadh and fadh2 these are biochemical reducing agents that act as electron carriers in cellular respiration and other metabolic processes now as a fun fact biochemical Redux reagents such as NAD plus they play crucial roles in metabolic pathways by acting as both oxidizing and reducing a agents nad+ accepts electrons during catabolic processes becoming reduced to nadh which stores energy during anabolic processes nadh donates electrons and becomes oxidized back to nad+ thus releasing the stored energy to drive biosynthesis this cyclical transformation allows NAD plus to mediate energy transfer within cells and it facilitates both energy release from food molecules and energy utilization in building cellular components this dual role is essential for maintaining the flow of energy through metabolic pathways now the next thing that we want to talk about is oxidation numbers and assigning oxidation numbers an oxidation state is a bookkeeping technique for keeping track of electrons it treats all of an atoms bonds as if they were ionic assigning the electrons in each bond to the more electronegative element it is important of course to know which atom is oxidized in which atom is reduced oxidation numbers are assigned to atoms in order to keep track of the redistribution of electrons during chemical reactions based on the oxidation numbers of the reactants and products it is possible to determine how many electrons are gained or lost by each atom so understanding how to assign oxidation numbers is crucial for tracking the transfer of electrons in Redux reactions so we're going to go over a couple of rules that is going to help us assign oxidation numbers to atoms in compounds the first rule is that the oxidation state of an element in its neutral Elemental form is zero this rule applies because atoms in their Elemental form have not gained or lost electrons so for example if we look at Helium it's going to have an oxidation number of zero if we look at diatomic nitrogen each nitrogen atom is going to have an oxidation number of zero the second rule is that the oxidation state of a monatomic ion is the same as its ionic charge this means that the oxidation state directly corresponds to the ion's charge for example if you look at sodium ion it's going to have an oxidation number of + one since it's lost one electron if we look at a chloride ion it's going to have an oxidation number of minus one because it has gained an electron the third rule is that in binary compounds the element with the greater electr negativity is assigned the negative oxidation state that is equal to the charge it would have if it were a monatomic ion so if we look at something like sodium chloride we're going to look at the most electronegative element that's going to be chlorine here all right and we're going to assign it an oxidation number that's equal to the charge it would be if it were a monatomic ion if it were chloride ion so it will have an oxidation number of minus1 now there's a lot that goes in here that we can discuss about different kinds of groups or elements and what they're typical oxidation numbers will be in compounds we're going to discuss that shortly but let me just talk about the last and final rule which says that the sum of the oxidation states of every atom equals the overall charge of the substance so the sum of the oxidation numbers of all the atoms present in a neutral compound is zero the sum of the oxidation numbers of the atoms present in a polyon IC ion that's going to be equal to the charge of the ion so for example if you're looking at a sulfate ion the sum of the oxidation numbers has to sum up to be min-2 to match the charge of the ion now like I said for rule three and for rule four what's going to help us is understanding some general trends for oxidation numbers for certain elements in a compound so we're going to scroll down here and look at our periodic table and I'm going to go over a couple of things the first thing that I want to go over is that the oxidation number of alkali metals all right for alkali metals in a compound is always going to be + one this is because alkaline metals tend to lose one electron to achieve a stable electronic configuration so let's go ahead and look at our example for sodium chloride again we said we traditionally look at our most electronegative element first and give it the oxidation number that it would have if it were a monatomic ion so it gets a NE one charge now sodium sodium is an alkaline metal it's going to have an oxidation state of + one and we know that we did this correctly because the sum of the oxidation numbers of all the atoms present in a neutral compound is zero and + one minus one sums up to zero now also another example that we can look at is for lithium chloride same thing goes here all right what is going to be the most electronegative element it's going to be chlorine we're going to give it an oxidation number as if it were a monatomic ion as if it were the chloride ion so it gets an oxidation number of minus1 lithium is an alkal metal alkaline metals have an oxidation state of + one and again we know we did this correctly because the sum of the oxidation numbers of all the atoms present in a neutral compound should add up to zero and it does now a second thing to keep in mind is that the oxidation number for alkaline earth metals in a compound is always + two these Metals typically lose two electrons to form stable cations so for example in calcium chloride calcium has an oxidation state of +2 for chlorine all right this is going to be minus1 and there's two of them so that gives us a total of -2 + 2 - 2 gives us a total of zero and remember that the sum of the oxidation numbers of all atoms present in this neutral compound of calcium chloride should be zero and it is zero so that's how we know we assigned these oxidation numbers appropriately now the oxidation number of hallogen elements in a compound is usually minus1 except when they're combined with an element of higher electron negativity so halogens they typically gain one electron to form stable annion so for example if we're looking at hydrogen chloride the chlorine has an oxidation number of minus1 all right however if we look at something else like hypochlorous acid h o chlorine is bonded to oxygen oxygen is more electronegative so chlorine will actually here have an oxid number of + one okay so that's four hogen atoms keep that in mind something else hydrogen the oxidation number of hydrogen is usually + one all right however when hydrogen is bonded to Metals in metal hydrides its oxidation number is actually going to be minus one because it acts as the more electronegative atom what's an example of this you ask so sodium hydde is a great example here sodium is a alkaline metal alkaline metals always have an oxidation state of + one all right of + one so here hydrogen will actually actually have a charge of minus1 and the way that you know you did this correctly is again remember our last and final rule we covered the sum of the oxidation numbers of all the atoms present in a neutral compound sums up to zero what about oxygen oxygen is a fun thing to talk about the oxidation number of oxygen is min-2 because oxygen typically gains two electrons to complete its veence shell for example in water H2 this oxygen atom all right in water is going to have an oxidation state of min-2 hydrogen has a oxidation state of + one and there's two of them so that gives us a total of + 2 + 2 - 2 0 that's exactly what we want for this neutral compound now there are exceptions for oxygen that is really important for us to talk about for example in peroxides H22 each oxygen atom actually has an oxidation number of minus1 this is because in this molecule we have an oxygen oxygen single Bond all right and in this oxygen oxygen single Bond the electrons are shared equally another exception to keep in mind in compounds Like Oxygen di floride o F2 oxygen is actually bonded to Florine which is more electronegative so here Florine is going to have the negative oxidation state of minus1 and there's two of them so that gives us min-2 here Oxygen's oxidation state is actually going to be plus two all right so those are the rules that we need to keep in mind for the MCAT again I want to repeat that last and final rule because it really helps us with a lot of the problems we're going to be asked in the MCAT and a lot of the problems we're going to tackle in our problem set that rule is that the sum of the oxidation numbers of all the atoms present in a neutral compound is zero and the sum of the oxidation numbers of the atoms present in a polyatomic ion is equal to the charge of the ion and I know I repeated that a lot of times but repetition is key that is a really important rule I want you to keep in mind now what we want to do next is tackle an example problem this example says aign oxidation number to the atoms in the following reaction to determine the oxidizing and reducing agents now just looking at this reaction all of these species are neutral so the oxidation numbers of each compound must add up to zero here we have 102 chloride reacts with lead 4 chloride to form 104 chloride and Lead 2 chloride let's start off with our first reactant 102 chloride now looking at this binary compound we need to remember rule three in binary compounds the element with the greater electr negativity is assigned a negative oxidation state that's equal to the charge it would have if it were a monatomic ion so here chlorine is obviously more electr negative so we're going to assign it a negative oxidation State that's equal to the charge it would have if it were chloride ion so we're going to assign it A1 charge now there's two of them so we have to multiply this by two we get a total of min-2 charge for cl2 now this is a neutral compound so tin is going to have an oxidation state of plus two to ensure that when you add up the oxidation a numbers here they add up to zero because this is a neutral compound again in this tin 2 chloride 10 has an oxidation number of plus two because there are two chlorines present and each chlorine atom has an oxidation number of minus1 similarly we can look at 104 chloride here okay we have four chlorine atoms each has a minus1 oxidation number for all four of these this gives us a total of Min -4 that means 10 in this compound is going to have an oxidation number of + 4 wonderful let's go ahead and look at lead for chloride here again chlorine atoms are the most electronegative we're going to assign it a minus1 oxidation number there's four chlorine atoms for a total of-4 4 that means that lead here is going to have an oxidation number of plus4 to ensure that for this neutral compound the sum of the oxidation numbers adds up to zero similarly we can do this for lead 2 chloride again chlorine is the most electr negative atom here we're going to assign it a -1 charge there's two of them that gives this a total of min-2 that means that lead in this compound is going to have an oxidation number of + two okay now the oxidation number of 10 it goes from + 2 to+ 4 it loses electrons and thus it's oxidized it's oxidized which makes it the reducing agent 10 is our reducing ing agent now looking at lead because the oxidation number of lead goes from + 4 to + 2 it gains electrons and so it is reduced making it the oxidizing agent the sum of the CH of the charges on both sides of this reaction is equal to zero so we know that charge has been conserved and so at the end of this to answer this problem we want to assign our oxidizing and reducing agents tin is our reducing agent lead is our oxidizing agent and with that we have completed this example problem and we're actually going to end this first video for this chapter right here in the next video we're going to continue objective one and finish objective two let me know if you have any questions comments concerns down below other than that good luck happy studying and have a beautiful beautiful day future doctors