GROUP 2 & 7 CRAM

Added: 2026-05-31

[00:00:00]Hey all, hope your revision's going well.

[00:00:02]Uh you've probably seen I've done a cram video [music] for paper one, but somebody's messaged me um and said I've missed I've missed a couple of things out, which are group two and the halogens. So, this video is just [music] a mashup of the two quick revision videos I've already made for those topics. Um so, hence the little introduction [music] to explain why I'm doing them. All right, hope you find them useful. Cheers, bye. Quick revision video on group two.

[00:00:31]So, we'll start with some general information first. All group two metals have the outer shell S2 electron configuration.

[00:00:41]They lose these outer two electrons and form two plus ions when they react.

[00:00:47]So, because they lose electrons, they are classed as reducing agents, electron donors.

[00:00:54]And they only lose two electrons because the third ionization energy is too large. Because to get the third electron out, you're breaking into a new shell.

[00:01:08]So, the reactivity of group two metals now.

[00:01:11]So, the trend is they get more reactive as you go down the group because of the atomic radius increases.

[00:01:19]So, the outer electrons, those two outer electrons that they are losing, they're getting further from the nucleus.

[00:01:26]The amount of shells between the nucleus and the outer electrons increases, so shielding increases.

[00:01:34]So, because of that, the attraction between the nucleus and the outer electrons decreases.

[00:01:41]And as a result of all of that, the electrons are therefore lost more easily. So, their first and second ionization energies decrease as you go down the group, but remember the third ionization energy is very large.

[00:01:58]Reactions of the elements now.

[00:02:00]So, when they react with oxygen, they form oxides.

[00:02:04]So, there's an example there using magnesium.

[00:02:09]When they react with water, they form hydroxides and hydrogen.

[00:02:15]So, there's an example with calcium.

[00:02:21]And when they react with dilute acids, they form a salt and hydrogen.

[00:02:27]So, I've used barium and hydrochloric acid to illustrate that one.

[00:02:34]Reactions of the compounds now. So, it's just one reaction here. It's the action of water on the oxides. So, I'm using beryllium oxide to illustrate this one.

[00:02:44]So, we get beryllium hydroxide.

[00:02:47]The pH of these hydroxides is greater than seven because they are alkaline.

[00:02:53]And as you go down group two, the hydroxides get more soluble.

[00:02:58]And so, the hydroxides will become more alkaline because you're going to have a higher concentration of hydroxide ions.

[00:03:09]So, we'll finish with some uses of group two compounds.

[00:03:12]So, calcium hydroxide is used to neutralize acidic soils. So, there's an equation that illustrates that.

[00:03:22]Magnesium hydroxide is used as what's called an antacid. So, that's found in indigestion remedies. So, it's going to neutralize excess stomach acid.

[00:03:32]So, there's the equation that goes with that.

[00:03:36]And the final use I've got is calcium carbonate. Again, being used as an antacid. So, it's reacting with hydrochloric acid in the stomach.

[00:03:46]And that's the equation there.

[00:03:51]Quick revision video for the halogens topic. So, we'll start off with the physical properties of the halogens. So, halogens occur as diatomic molecules, so F2, Cl2, etc. In the solid state, they form simple covalent lattices.

[00:04:07]Their boiling and melting points increase down the group. So, there's a summary table of all the halogen molecules with their total number of electrons in the diatomic molecule. You can see the boiling points are increasing.

[00:04:20]And the appearance at room temperature and pressure. So, fluorine's a pale yellow gas.

[00:04:25]Uh chlorine's a pale green gas.

[00:04:28]Bromine's a red-brown liquid.

[00:04:30]And iodine is a shiny gray-black solid.

[00:04:35]So, we'll just quickly explain that. So, we think about the lattice, the So, we'll be talking about the melting point here. So, to break the um lattice down, we've got to overcome the intermolecular forces, the induced dipole-dipole forces between the molecules.

[00:04:51]And as the number of electrons in the molecules increases, these intermolecular forces get stronger. Just remember that we're not breaking the covalent bond between the atoms in the molecule. It's the induced dipole-dipole forces or London forces, you could call them.

[00:05:09]So, if we move on to the chemical properties of the halogens now, they all have seven valence electrons. So, their configuration is S2 P5 in the outer shell.

[00:05:21]They all accept one electron to become a one minus halide ion. So, fluorine becomes fluoride, F minus, and so on.

[00:05:30]Since the halogens are electron acceptors, they are classed as oxidizing agents. So, they cause other substances to lose electrons.

[00:05:40]The oxidizing power, their ability to accept an electron, decreases down the group.

[00:05:46]And that's due to the increasing atomic radius of the atom, the increased amount of shielding, they've got more shells, and therefore the attraction from the nucleus to the electron that's being attracted in gets weaker.

[00:06:02]And this trend in oxidizing power is demonstrated by the halogen halide displacement reactions, which we'll look at on the next slide.

[00:06:10]So, the procedure for these reactions is as follows. A solution of a halogen is added to a solution of a halide ion.

[00:06:18]If the halogen's more reactive, it's going to displace the less reactive halide, and the color of the test tube at the end of the reaction indicates which halogen is present.

[00:06:29]Now, the colors of the halogens aren't very sort of distinguishable from each other in aqueous conditions, and so what we do is we add a small amount of organic solvent to make the halogen much more easy to identify.

[00:06:45]So, I'll show you an example of that now.

[00:06:48]So, chlorine with KBR, so chlorine's more reactive, so it will displace the bromine. So, that's the full equation for that reaction.

[00:06:57]There's the ionic equation for the same reaction, and it's a redox reaction. Let's quickly explain that. Chlorine starts out its oxidation state zero as the element, goes to -1, so that's a reduction process, and the bromine starts out with oxidation number -1 in the R minus, and it goes up, it increases to zero. So, that's an oxidation process. So, we've got a reduction and an oxidation process happening in the same reaction, hence redox.

[00:07:26]Got this little cartoon here just to illustrate what's happening. So, there's your more reactive halogen, so chlorine in this case is two chlorine atoms covalently bonded together.

[00:07:36]There's your two separate bromide ions, and that ice cream is meant to represent the electron.

[00:07:41]So, because chlorine is more reactive, it's got smaller atomic radius, less shielding, greater attraction for the electron, it's going to pull the electron off the bromide ion.

[00:07:51]And the chlorine atoms will become chloride ions. And you're left with two separate bromine atoms, which then just covalently bond and form the diatomic bromine molecule.

[00:08:03]So, what would that look like in a test tube? So, you've got the bromine, the displaced bromine in that aqueous layer at the bottom of the test tube there.

[00:08:12]Um it's sort of a very pale yellow. But if you add some cyclohexane, an organic solvent, it will um the bromine dissolves much better in that layer, and you see its color more intensely. So, you get that um orange layer there for the bromine.

[00:08:29]So, we'll just look at the colors of all of the halogens in water and the organic solvent. So, there's a photo of all three. So, I'll just go through from left to right. So, chlorine, this this test tube here, it's very, very pale green in both uh layers. Aqueous is the lower layer. Organic solvent is the upper layer. It often looks colorless cuz it's so pale green.

[00:08:52]We've just seen that one in that example. So, that's your bromine, which is yellow in water and orange in the organic solvent.

[00:09:01]And this one here is the iodine, which is sort of a pale brown color in aqueous or water conditions, and violet in the organic solvent.

[00:09:12]So, on this slide, we'll look at the uses of chlorine. So, the first use is its um addition to drinking water to kill harmful bacteria. E.g., the bacteria that would cause cholera. So, the equation for that reaction looks like this: Cl2 + H2O gives HCl and HClO.

[00:09:29]The bacteria is killed by the chlorite one ion, ClO- ion here in this This is called chloric one acid, and that is the iron that kills the bacteria.

[00:09:43]The other reaction is the production of bleach. So, that's when chlorine is reacted with cold dilute sodium hydroxide.

[00:09:50]So, there's the equation there. So, you can see it's quite similar. We get instead of HCl, we get NaCl.

[00:09:57]Instead of HClO, we get NaClO. And we also get water.

[00:10:03]The bleach is this chemical here, sodium chlorate one.

[00:10:07]And you'll see it also contains that chlorate one ion.

[00:10:13]These are both examples of disproportionation reactions because the chlorine is oxidized and reduced. So, we'll just use this one to illustrate that.

[00:10:23]Chlorine starts out at zero in the element. It goes to -1 in HCl. So, that's a reduction process. It's a drop in the oxidation number.

[00:10:33]In here, it's +1.

[00:10:36]The oxygen's -2. Hydrogen's +1. So, the chlorine needs to be +1 to keep this overall neutral.

[00:10:45]Just look at some arguments for and against the chlorination of water now.

[00:10:49]So, arguments for, things like, well, it kills harmful bacteria. So, obviously that's going to reduce the risk of waterborne diseases such as cholera. So, it's got massive public health benefits.

[00:11:00]It prevents the formation of algae in water.

[00:11:05]It eliminates bad tastes and bad smells in water. And finally, it persists. The chlorine persists in the water, so it has a longer-lasting effect compared to alternatives such as the addition of ozone or UV to the water.

[00:11:22]Some arguments against chlorination of water now. So, chlorine is toxic. It's also a respiratory irritant. Chlorine reacts with organic matter that's in the water, and that can form chlorinated hydrocarbons, and the problem with those is they are carcinogens or cancer-causing chemicals. And another argument against is just that some people are against government intervention.

[00:11:45]So, the final aspect we're looking at now is the test for the aqueous halide ions. So, I'll run through the procedure first and then show you the observations, put in photos of the results.

[00:11:57]So, when silver nitrate is added to a solution that contains a halide ion, so it could be chloride ions, bromide ions, iodide ions, you get a silver halide precipitate.

[00:12:09]These silver halide precipitates have different colors.

[00:12:13]And they also have different solubilities in aqueous ammonia. So, we'll just run through the procedure now for the test for halide ion. I'm going to use chloride ions to illustrate this, but it's the same procedure for all of them.

[00:12:27]So, you take your test tube containing your aqueous halide ions or chloride ions in this case, and the first thing you do is add a small amount of dilute nitric acid, and the purpose of that is to remove any carbonate ions that might be present. So, if you've got carbonate ions present, they will react with the silver ions from the silver nitrate and give a precipitate of silver carbonate.

[00:12:49]So, you'd get a precipitate, and you would think it was a silver halide, but it's actually silver carbonate. It's It's giving you a false positive result.

[00:13:00]Once you've got rid of the carbonate ions, you would add a small amount of silver nitrate solution, and you'd observe the color of the silver halide precipitate. So, silver chloride in this case is white.

[00:13:13]You'd then add a small amount of dilute aqueous ammonia to the precipitate, and silver chloride actually fully dissolves in dilute aqueous ammonia. So, that white precipitate would disappear, and you'd just have a colorless solution left at the end of that process.

[00:13:30]So, the ionic equation for the reaction looks like that. So, the silver ions, the aqueous silver ions from the silver nitrate solution, combine with the aqueous chloride ions and give that solid silver chloride precipitate. So, because we're getting a precipitate in this reaction, it's called a precipitation reaction.

[00:13:50]So, I'll just do a summary of of all the halide ion tests now because obviously that previous slide just dealt with chloride. So, we've got this table here.

[00:13:59]So, this is what we've just observed.

[00:14:02]So, chloride ions, there's the ionic equation, there's the color of the precipitate, and there's its solubility in aqueous ammonia.

[00:14:11]So, if we move on to the bromide ion now, same looking um ionic equation, the precipitate is cream, and that's partially soluble in dilute aqueous ammonia, but it's fully soluble in concentrated aqueous ammonia.

[00:14:26]And then finally, the iodide ion equation looks like that. It's a yellow precipitate, silver iodide's yellow solid, and that's insoluble in both types of aqueous ammonia.

[00:14:39]So, I'll just show you this photo here.

[00:14:40]So, this test tube here is actually silver chloride.

[00:14:45]This one here is slightly creamier color, that's silver bromide. And this one here is silver iodide. So, the great thing about the solubility in the aqueous ammonia, it enables you to distinguish between certainly between those two cuz they are very very similar colors.

[00:15:03]So, if you let's say you had this test tube as your result, and you can't decide whether it's white or cream, if you added dilute aqueous ammonia to that and it dissolved, that's confirmation that it is um the chloride.

[00:15:18]If it only partially dissolved, then you could say well, it's not a chloride, so it's probably going to be a bromide cuz it's not very yellow color, but you could just then throw in some concentrated aqueous ammonia, and if it totally dissolved in the concentrated, you could say that it was bromide.

[00:15:37]Um if it didn't dissolve at all in the concentrated, then you would be dealing with an iodide.