 Let's now talk about the group 17 elements, the hyperreactive supercharged halogens. Now as you can see here we are almost at the end of the periodic table, which means now we are looking at these elements that have the smallest size and highest electronegativity in their respective periods. As they are just one electron short of the stable octet configuration, all of these halogens are extremely reactive. They react with metals and non-metals to form an entire class of compounds called halides. However, within this group we can see that the reactivity decreases as we go down the group from fluorine to iodine. We won't discuss much about acetone because it is an unstable radioactive element and almost all of its isotopes are very short-lived. So that means for all our discussions henceforth we will mostly be talking about fluorine, chlorine, bromine and iodine. So in this video we will briefly look at the atomic properties of the group 17 elements especially their negative electron gain enthalpy. We will also discuss their oxidation states and look at scenarios where halogens can exhibit positive oxidation states. And lastly we will look at the oxidizing power or oxidizing nature of halogens and see how it varies across the group. Now if you look at their atomic properties, halogens have the small less size, highest ionization enthalpy and highest electronegativity in their respective periods. And as they have very little tendency to lose electrons, they also have maximum negative electron gain enthalpy. Not surprising at all because they are just one electron short of the stable octet configuration right? And because it is small size and high effective nuclear charge, the incoming electron is actually quite close to the nucleus and this kind of helps stabilize it. Although we just have to deal with one exception. Just as we observed in the case of oxygen versus sulfur, the negative electron gain enthalpy of fluorine is lesser than that of chlorine. Just as we observed in the case of group 16 where the negative electron gain enthalpy of oxygen is again lesser than that of sulfur. Now this is again because of the super small size of fluorine. Now because it is small size of fluorine, the incoming electron experiences stronger electron-electron repulsion in the small 2p orbitals. On the other hand, chlorine has relatively larger 3p orbitals and as a result the incoming electron experiences much less electron-electron repulsion. And this is why less energy is released when you add an electron to fluorine as compared to chlorine. Now coming to oxidation state, all halogens exhibit minus 1 oxidation state. However, except fluorine, all other halogens also exhibit positive oxidation states. So here again the logic is that if these halogens being electronegative, if they have to exhibit a positive oxidation state that means they have to give away their electron right? And that means they have to combine with an element that is more electronegative than themselves. And this is why chlorine, bromine and iodine usually exhibit higher oxidation states or positive oxidation states when they combine with smaller highly electronegative elements like oxygen or fluorine. For example, positive oxidation states are usually seen in interhalogen compounds like IBr5, ICl5 and so on. And here the oxidation state of iodine is plus 5 and as you can see, iodine combines with more electronegative elements than itself. Similarly, in oxides like ClO2, BrO2 and oxoacids like HClO4 or HBrO4, halogens predominantly show plus 4 and plus 7 oxidation states. But fluorine on the other hand can show only minus 1 oxidation state because it is the most electronegative element. And this readiness or easiness to accept an electron makes fluorine a very strong oxidizing agent. In fact, fluorine is such a good oxidizing agent that it can literally oxidize other halide ions from their solutions. That is, it can displace all other halide ions like chlorine, bromine and iodine from their solutions in a reaction like this. In fact, we can extend this logic and say that any halogen will be able to oxidize the halide ions that are higher in their atomic number. For example, when chlorine acts as an oxidizing agent, it can displace halogens that are higher in number than chlorine, right? Which means it can displace bromide ions and iodine ions from their solutions. Similarly, when bromine acts as an oxidizing agent, the only halogen that is higher in number than bromine is iron. That means it can only displace iodine ions from their solutions. So clearly we can see that this oxidizing nature or the strength of oxidizing power decreases as we go down the group. We can also understand the decreasing oxidizing ability by looking at the reaction of these halogens with water. For instance, in the reaction of fluorine with water, fluorine oxidizes water to oxygen gas. And if you replace fluorine with chlorine or bromine, you can see that they react with water to form corresponding hydrohalic and hypohalus acids. On the other extreme, the reaction of iron with water is completely non-spontaneous. But again you can see how the oxidizing power decreases as we go down the group. So to quickly summarize, we learned that halogens have the smallest size, highest ionization enthalpy and electronegativity in the respective periods. And they also have the maximum negative electron gain enthalpy. The only exception was that of fluorine with chlorine. We then looked at the various oxidation states of halogens. We saw in which cases halogens show positive oxidation states. And we also discussed the oxidizing power of the halogens by taking the example of the reaction with water.