in this video we will investigate electron transfer or Redux reactions which is just one type of the many types of chemical reactions during a Redux reaction one or more electrons are transferred from one reactant to another the process of losing electrons is called oxidation in this example substance a loses an electron and is oxidized the process of gaining electrons is called reduction substance B gains an electron and is reduced note that if substance B is a positive ion it might be reduced to a neutral particle instead of a negative ion as shown here oxidation of one species is always accompanied by reduction of another species the two go hand in hand and occur at the same time since both processes occur simultaneously the term Redux describes these reactions red for reduction and Ox for oxidation we can use the handy acronym oil rig to remember that oxidation is the loss of electrons and reduction is the gain of electrons here is an example of a Redux reaction at the beginning of the reaction zinc metal is placed in a Sol solution containing copper 2+ ions during the reaction zinc loses two electrons and forms the zinc 2+ ion in solution zinc is oxidized as the reaction proceeds so the blue color of the copper 2+ solution would fade and copper would be deposited as a dark solid on the remaining zinc metal the electrons lost by zinc are gained by or trans transferred to the Copper 2+ ions copper is reduced to its Elemental form we can break up a Redux reaction into two half reactions or half equations one for oxidation and one for reduction we can write the oxidation half reaction like this we can see that on the right hand side the electrons have been separated or removed from zinc since it was oxidized we can write the reduction half reaction like this we can see on the left hand side that electrons combine with a copper 2+ ion since it is reduced overall electrons are transferred from zinc metal to Copper ions the oxidizing agent is the species that causes the oxidation of another species by taking electrons from it since the copper 2+ ion gained or received electrons from zinc causing zinc to be oxidized the copper 2+ ion is the oxidizing agent the reducing agent is the species that causes the reduction of another species by giving electrons to it since zinc metal transferred electrons to The Copper 2+ ions causing them to be reduced Z Inc is the reducing agent in our Redux reaction example we deduced that zinc lost electrons and was oxidized since it went from a neutral state to a positive charge we also deduced that copper gained electrons and was reduced since its charge went from a positive 2 plus state to a neutral state in some Redux reactions dedicating which species is oxidized and which is reduced is more challenging we must therefore identify the oxidation state or oxidation number of each element in each species and investigate how these values change during the reaction to identify which species is reduced and which is oxidized the oxidation state on an atom is a hypothetical charge if we treated all of its bonds as if they were ionic or if all the atoms were ions oxidation numbers can be positive negative or zero the magnitude and sign of the oxidation number helps us identify the degree to which an atom is oxidized here are the rules for assigning oxidation numbers an element in its Elemental state has an oxidation number of zero in compounds group one Metals always have A+ one oxidation number and group two metals always have A+ two oxidation number halogens are minus1 hydrogen at the top of group one is usually + one unless bonded to a metal where it is minus one oxygen is usually Min -2 except in peroxides where it is minus1 and transition metals and most main group non-metals have variable oxidation numbers depending on the compound they are in the remaining rules are the sum of the oxidation numbers in a neutral compound is always equal to zero the sum of the oxidation numbers in a polyatomic ion equals the charge on the ion note that since there are four oxygen atoms in this example each with an oxidation number of minus 2 together they contribute a total of minus8 the oxidation number increases during oxidation it becomes more positive or less negative and the oxidation number decreases during reduction becoming less positive or more negative let's use these rules to identify which species is oxidized and which is reduced in this reaction according to the oxidation number rule pottassium which is in group one must have A+ one oxidation number the iodide ion has a minus1 oxidation number hydrogen is always + one when bonded to non-metals and in hydrogen peroxide since there are two hydrogen atoms they together contribute plus two oxygen usually has an oxidation number of min-2 except in peroxides where it is min-1 and since there are two oxygen atoms in peroxide we write min-2 at the top elements in their Elemental State have an oxidation number equal to zero so we can write zero Above This datomic iodine potassium's oxidation number is unchanged hydrogen's oxidation number is unchanged iodine's oxidation number has increased it has become less negative or more positive and so potassium iodide is oxidized Oxygen's oxidation number has decreased it has become more negative so hydrogen peroxide is reduced potassium iodide causes the reduction of hydrogen peroxide and so potassium iodide is the reducing agent hydrogen peroxide causes the oxidation of potassium iodide and so hydrogen peroxide is the oxidizing agent in this example we can see the progressive oxidation of a transition metal resulting in varying possible oxidation states in the first step two electrons are lost and in the second oxidation one electron is lost we have seen that oxidation is the loss of electrons or the increase in oxidation number it can also be considered to be the gain of oxygen atoms these two reactions are examples of the gain of oxygen in the first equation sulfur in sulfur dioxide gains oxygen and is oxidized in the second equation sodium gains oxygen and is oxidized if we were to check the oxidation numbers in these equations we would find that oxygen is reduced in in both cases we have also seen that reduction is the gain of electrons or the decrease in oxidation number however it can also be considered to be the loss of oxygen or the gain of hydrogen these two reactions are examples of reduction of the non oxygen element in the first equation carbon dioxide loses oxygen and so carbon is reduced in the second equation chlorine gains hydrogen and so is reduced if we were to check the oxidation numbers we would find that in equation one oxygen is oxidized and in equation 2 hydrogen is oxidized you can pause the video at this point to examine these two slightly more complex examples of oxidation and reduction regarding oxygen or hydrogen loss or gain different elements have different Tendencies to be either oxidized or reduced some elements are more easily oxidized than other elements and some elements are more easily reduced metals tend to lose their outer electrons and be oxidized when they react so the reactivity of a metal is a measure of its tendency to lose outer electrons and form positive ions since non-metals gain the electrons lost by Metals metals are reducing agents non-metals tend to gain electrons and be reduced when they react except for the noble gases and Except for hydrogen which tends to lose an electron like the other elements in group one the reactivity of a non-metal is a measure of its tendency to gain outer El electrons and form negative ions since non-metals gain electrons by removing electrons from another element during a reaction some non-metals can act as oxidizing agents let's investigate the differing ease with which different metals react and are oxidized for example the oxidation of group one Metals the alkaline metals is a rapid spontaneous process when a clean surface of a group one metal is exposed to oxygen in the air it rapidly oxidizes to form a white layer of metal oxide that is why alkal Metals need to be stored under oil all group one metals react energetically they only need to lose one electron to make them stable relatively little energy is required to remove this one veence electron Going Down group one the reaction with oxygen gets more vigorous as there are more core shells shielding the veence electron from the nuclear pull and less energy is required to remove the veence electron and so we say the reactivity of the metals increases going down the group this reactivity trend is very noticeable when a piece of alkal metal is placed on the surface of water lithium pops and fizzes as it reacts while cesium produces Sparks and a small explosion is heard a similar trend is found going down other groups containing Metals in general going down a group metal reactivity or ease of oxidation increases however this trend differs from left to right across a period transision metals are less reactive than the metals in groups 1 and two for example rusting is a well-known oxidation reaction of the transition metal ion but this oxidation process occurs slowly going further to the right we find some transition metals which do not react these inert or unreactive Metals include gold and platinum The General enal Trend across a period is decreasing reactivity or ease of oxidation from left to right or increasing reactivity from right to left this is because going to the left across a period so the nuclear charge decreases due to decreasing protons and so less energy is required to remove veence electrons now let's investigate the differing ease with which different non-metals react excluding hydrogen and the noble gases the halogens in group 17 are readily reduced they need only one electron to fill their outer shell but the ease with which they gain this electron during a reaction differs going up the group there are fewer core shells and less shielding so the attractive force from the nucleus for an incoming electron in increases in turn reactivity increases for example the Vigor with which the halogens are reduced when they react with hydrogen increases going up the group Florine reacts explosively even in the dark and at very low temperatures while iodine near the bottom of the group needs high temperatures and a catalyst to react a similar trend is found going from left to right across a period the reactivity of non-metals increases this is because of the increasing number of protons going from left to right and increasing nuclear attraction for an incoming electron we can put the reactivity Trends together on one periodic table like this we need to know and understand these General reactivity Trends on the periodic table as this will help us deduce which element is oxidized in a Redux reaction and which is reduced here is an example in which Beaker will a reaction occur write the Redux equation that occurs from the reactivity Trends we know that metal reactivity or ease of oxidation tends to decrease from left to right across the periodic table since calcium is further to the left than copper it is more reactive than copper this means that calcium will preferentially be oxidized and therefore copper will be reduced so a Redux reaction would occur in Bea a and this is the Redux equation you might be wondering why in a previous example we looked at zinc is oxidized and copper is reduced in this Redux reaction instead of the other way around well zinc's outer electron configuration is 4s2 3d10 and Copper's is expected to be 4s2 3d9 however an atom is more stable if its D subshell is either completely empty half filled or completely filled and so one of Copper's 4S electrons is used to fill the 3D suev leaving copper with a stable outer electron configuration of 4s1 3d10 then when zinc reacts to form the zinc 2+ ion it loses both its outer electrons that is its two 4S electrons this leaves the zinc 2+ Ion with a full and stable D subshell since zinc achieves a stable State when it forms the 2+ ion zinc reacts readily however when copper loses two electron to form a copper 2+ ion it is left with an incomplete D subshell since the state is less stable than a completely filled D subshell copper reacts less readily than zinc so zinc is more reactive than copper despite being further to the right on the period rest assured that you will always be given enough information to deduce the relative reactivities of elements which do not follow the general Trends from knowledge of reactivity Trends or data given to us in a particular question we could construct an activity or reactivity series which is a list of elements in order of decreasing reactivity elements at the top of the list are readily oxidized and are strong reducing agents while elements at the bottom are difficult to oxidize and are weak reducing agents be aware that the non-metals hydrogen and carbon are often included in an activity series near the middle it's time to summarize the main points we have learned we learned that Redux reactions involve the transfer of electrons oxidation is the loss of electrons the gain of oxygen or the loss of hydrogen reduction is the gain of electrons the gain of Hy hydrogen or the loss of oxygen a Redux reaction can be written as two half equations an oxidation and a reduction half equation a reducing agent causes reduction of another species and an oxidizing agent causes oxidation of another species oxidation numbers increase during oxidation and decrease during reduction different elements display differing ease of oxidation or reduction and finally reactivity Trends occur across periods and downg groups of the periodic table