1.6 THE PERIODIC TABLE- HALOGENS AND PRACTICAL WORK

IF YOU WOULD LIKE A COPY OF PAST PAPER QUESTIONS FOR THIS TOPIC OR ANY OTHER TOPIC PLEASE EMAIL ME (space is provided on the main page of my blog)

IF YOU WOULD LIKE TO GO THROUGH THE QUESTIONS WITH ME, PLEASE VISIT MY YOUTUBE CHANNEL

Hi people! I hope my blog helps you with AS-level chemistry. My goal for the next school year is to prepare past paper questions and posts with all A-level content. Everything that you'll find here is written in the most basic form you can get but at the same time, this is all you need for your exams. Of course, you could get many other websites written by professors and they will definitely be more detailed. This is an explanation of the topics from a student who just finished A-levels.

Ok, so this post focusses on halogens and we'll firstly talk about the group of halogens as a whole.

1. Halogens are in group 7 (on the right-hand side of the periodic table) and this group consists of non-metals.

2. Halogens are diatomic molecules. This means they always go in pairs, I2, Cl2, Br2, etc. Two molecules (in this diatomic molecule) are bonded covalently to each other (very strong bonds). However, these diatomic molecules are bonded together by Van der Waals (very weak forces).

3. As you go down the group the boiling point of halogens increases. As the halogens increase, the number of forces increases. The reason for this is the increased number of electrons, which means more electrons contribute to the induced-dipole, induced-dipole interactions.

4. The electronegativity decreases as you go down the halogen group owing to the size of the molecules. As you go down the group the halogens get larger and therefore, we have a weaker nuclear charge (the more electrons, the less they are held by the nucleus).

5. As we mentioned, halogens are in group 7 of the periodic table, so they need one electron in order to be stable and become an anion. In the photo above, you can see an equation related to this point.

6. When you react halogens with metals, you form halides. This is called a displacement reaction:

Let's start with a definition:

Displacement reactions - part of one reactant is displaced by another reactant.

We talked about fluorine being the most electronegative element in the periodic table (electronegativity increases across the period and up the group) and the order of electronegativity goes like this:

fluorine > chlorine > bromine> iodine

This is also the order of how halogens displace each other. So, if we had a chlorine and potassium bromide, this would change to potassium chloride and bromine. In this example, chlorine displaces bromine.

--> Fluorine - displaces any halogen because it is the most electronegative;

--> Chlorine - will displace bromine, and iodine but won't displace fluorine;

--> Bromine- will displace iodine but won't displace bromine, chlorine and fluorine;

--> Iodine- won't displace any of the above;

On the photo above, we have a table with all the possible combinations. Some elements will displace the others and with some nothing will happen at all. Of course, if we have NaCl and add chlorine, chlorine can't displace itself from the sodium chloride, therefore nothing will happen.

Lastly, there are a couple of equations that I included in the photo above. The 2 first ones are the overall equations and the two bottom ones are the ionic equations (the more detailed ones).

1. Chlorine with sodium bromide- chlorine displaces bromide, leaving bromine on its own.

2. Bromine with sodium iodide - bromine displaces iodide, leaving iodine on its own.

Then, we have ionic equations too which are more detailed. I included them just in case you would be asked about them in the exam.

All the aqueous solutions dissociate into ions. Bromine is liquid, so it won't. We can cancel out all the same ions that are on both sides of the equation, these are called spectator ions. The ones that aren't spectator ions are going to represent the overall ionic equation.

Exactly the same thing happens with the second equation. So, just go through the equation and I am sure you will follow it.

Displacement reactions are known as redox equations. Underneath the equations, I wrote the oxidation numbers of each of them. The oxidation number of chlorine molecule is 0 and that of the ions is -1.

• Chlorine is an oxidising agent, it oxidises bromide ions to bromine. It gets reduced (gains electrons to become an ion) as the oxidation number goes from 0 to -1.

• Bromine is oxidised (lost electrons) as the oxidation number goes from -1 to 0.

• As I included in the photo above, the bromide ions come from NaBr, however, sodium ions are spectator ions because they are in aqueous solution on both sides, so we don't have to include them.

Silver nitrate test:

A silver nitrate test is a test for halogens. This reaction is something we already went through in previous posts. What you do is add sulfuric acid and then add silver nitrate.

Above, I included an equation for this reaction. We have silver ions and the halogen ions and then a solid is formed, this is a precipitate. The sodium ions are spectator ions. Different colours of the precipitates relate to the different halogens.

--> Chlorine - gives white precipitate;

--> Bromine- gives cream precipitate;

--> Iodine- gives yellow precipitate;

There is one further test that can be done if you can't recognise the colour easily and it is with ammonia.

--> Chlorine- dissolves in dilute ammonia;

--> Bromine- dissolves in concentrated ammonia;

--> Iodine - doesn't dissolve in ammonia at all;

Chlorine and fluorine:

Chlorine:

• Used in water treatment and makes it safe to drink;

• It kills the bacteria in water;

Fluorine:

• It reduces tooth decay, used in toothpaste;

• It strengthens the bones, used in water;

There are also some disadvantages in using chlorine and fluorine such as:

• Too much can cause health problems;

• Can leave a nasty taste;

However, the advantages outweigh the disadvantages, therefore we just need to control the amount that is added.

---

Soluble salt formation:

There is a practical work that you need to recognise and it is a soluble salt formation.

Here are a couple of steps that have to be followed to form a soluble salt.

1. Base and acid are added together and you keep adding the base to the acid until no more base dissolves.

2. The excess solid (base) is filtered out (at this point all the acid has been used up).

3. Then you heat the solution to evaporate some of the water (but only some, otherwise you will form a powder and not crystals).

4. Leave to cool.

That's the process of the reactions, they may ask you in the exam to write the process or just some points from it, such as why you only evaporate some water and not all, etc.

Ok, everyone, this is all for today's topic. Be sure to do some past paper questions which you get from my youtube channel (all links at the top of this page). Remember, that doing past paper questions helps. The questions seem to be slightly repetitive and they just change the numbers or context, but the answer is similar. Furthermore, by working through past papers, you can see what the exam board is expecting you to say in different types of questions.

PS. Please remember, I am only a student, and as anyone, I can make mistakes. If you think you can see one, don't hesitate and comment (either here on on my youtube channel) Thank you!