Halogens - Chlorine, Bromine, Iodine, Astatine, Fluorine, Unexplored Sources And Problems - Elements, Chemical, Compounds, and Salt - JRank Articles Halogens - Chlorine Chlorine was the first halogen to be separated and recognized as an element. It was named in 1811 by Humphry Davy from a Greek term for its greenish yellow color. Huge deposits of solid salt, mostly sodium chloride, and salts dissolved in the oceans are vast reservoirs of chloride compounds. "Salt" is a general term for a metal and nonmetal combination; there are many different salts… Halogens - Bromine The next heaviest element in the halogen family is bromine, named from a Greek word for stink, because of its strong and disagreeable odor. It was first isolated as an element in 1826. Bromine is a reddish brown liquid that vaporizes easily. The vapors are irritating to the eyes and throat. Elemental bromine is made by oxidation, removal of electrons from bromide ions in brine. Brines in Arkansas … Halogens - Iodine The heaviest stable halogen is iodine.

Iodine forms dark purple crystals, confirming its name, Greek for violet colored. It was first obtained in 1811 from the ashes of seaweed. Iodine is purified by heating the solid, which sublimes, or goes directly to the gas state. The pure solid is obtained by cooling the vapors. The vapors are irritating to eyes and mucous membranes. Iodine was obtained comm… Halogens - Astatine Astatine could be described as the most rare element on Earth. All isotopes, atoms with the same number of protons in the nucleus and different numbers of neutrons, are radioactive; even its name is Greek for "unstable." When an atom decays its nucleus breaks into smaller atoms, subatomic particles, and energy. Astatine occurs naturally as one of the atoms produced when the uranium 2… Halogens - Fluorine Fluorine was the most difficult halogen to isolate because it is so chemically reactive. H. Moissan first isolated elemental fluorine in 1886, more than seventy years after the first attempts. Moissan received the 1906 Nobel Prize for Chemistry for this work.

The technique that he developed, electrolysis of potassium fluoride in anhydrous liquid hydrogen fluoride, is still used today, with some mo… Halogens as oxidising agents THE OXIDISING ABILITY OF THE GROUP 7 ELEMENTS (THE HALOGENS) This page explores the trend in oxidising ability of the Group 7 elements (the halogens) - fluorine, chlorine, bromine and iodine. We are going to look at the ability of one halogen to oxidise the ions of another one, and how that changes as you go down the Group. Note: If you aren't comfortable with terms like oxidation and oxidising agent in terms of electron transfer, then you should explore the area of the site dealing with redox reactions before you go on. The facts We are going to look at the reactions between one halogen (chlorine, say) and the ions of another one (iodide ions, perhaps). The iodide ions will be in a solution of a salt like sodium or potassium iodide. The sodium or potassium ions will be spectator ions, and are completely irrelevant to the reaction.

In the chlorine and iodide ion case, the reaction would be: The iodide ions have lost electrons to form iodine molecules. They have been oxidised. The chlorine molecules have gained electrons to form chloride ions. They have been reduced. This is obviously a redox reaction in which chlorine is acting as an oxidising agent. Fluorine We'll have to exclude fluorine from this descriptive bit, because it is too strong an oxidising agent. Fluorine oxidises water to oxygen and so it is impossible to do simple solution reactions with it. Chlorine, bromine and iodine In each case, a halogen higher in the Group can oxidise the ions of one lower down. For example, chlorine can oxidise the bromide ions (in, for example, potassium bromide solution) to bromine: The bromine appears as an orange solution. As you have seen above, chlorine can also oxidise iodide ions (in, for example, potassium iodide solution) to iodine: The iodine appears either as a red solution if you are mean with the amount of chlorine you use, or as a dark grey precipitate if the chlorine is in excess.

Note: The reason for the red solution is that iodine dissolves in potassium iodide (or other soluble iodides) by reacting to give a red ion, I3-. If the chlorine is in excess, obviously there isn't anything left for the iodine to react with, and so it remains as a dark grey precipitate. Bromine can only oxidise iodide ions to iodine. It isn't a strong enough oxidising agent to convert chloride ions into chlorine. (You have just seen exactly the reverse of that happening.) A red solution of iodine is formed (see the note above) until the bromine is in excess. Then you get a dark grey precipitate. Iodine won't oxidise any of the other halide ions (unless you happened to have some extremely radioactive and amazingly rare astatide ions - astatine is at the bottom of this Group). To summarise Oxidation is loss of electrons. Each of the elements (for example, chlorine) could potentially take electrons from something else to make their ions (e.g. Cl-). That means that they are all potentially oxidising agents. Fluorine is such a powerful oxidising agent that you can't reasonably do solution reactions with it.

Chlorine has the ability to take electrons from both bromide ions and iodide ions. Bromine and iodine can't get those electrons back from the chloride ions formed. That means that chlorine is a more powerful oxidising agent than either bromine or iodine. Similarly bromine is a more powerful oxidising agent than iodine. Bromine can remove electrons from iodide ions to give iodine - and the iodine can't get them back from the bromide ions formed. This all means that oxidising ability falls as you go down the Group. Explaining the trend Whenever one of these halogens is involved in oxidising something in solution, the halogen ends up as halide ions with water molecules attached to them. Looking at all four of the common halogens: As you go down the Group, the ease with which these hydrated ions are formed falls, and so the halogens become less good as oxidising agents - less ready to take electrons from something else. The reason that the hydrated ions form less readily as you go down the Group is a fairly complicated mixture of several factors. Unfortunately, this is often over-simplified to give what is actually a faulty and misleading explanation. We'll deal with this first before giving a proper explanation.

The faulty explanation This is normally given for the trend in oxidising ability of chlorine, bromine and iodine, and goes like this: How easily the element forms its ions depends on how strongly the new electrons are attracted. As the atoms get bigger, the new electrons find themselves further from the nucleus, and more and more screened from it by the inner electrons (offsetting the effect of the greater nuclear charge). The bigger atoms are therefore less good at attracting new electrons and forming ions. That sounds reasonable! What's wrong with it? What we are describing is the trend in electron affinity as you go from chlorine to bromine to iodine. Electron affinity tends to fall as you go down the Group. This is described in detail on another page. Note: If you haven't recently read about the electron affinities of the halogens , you ought to follow this link before you go on. Use the BACK button on your browser to return to this page. The snag comes if you try to expand the argument to include fluorine. Fluorine has a much higher tendency to form its hydrated ion than chlorine does. BUT . . .

the tendency of the fluorine atom to gain an electron is less than that of chlorine - as measured by its electron affinity! That makes a nonsense of the whole argument. So, what is going wrong? The mistake is to look at only one part of a much more complicated process. The argument about atoms accepting electrons applies to isolated atoms in the gas state picking up electrons to make isolated ions - also in the gas state. That's not what we should be talking about. In reality: The halogen starts as diatomic molecules, X2 - which may be gas, liquid or solid, depending on the halogen. These have to be split apart to make individual atoms. Those atoms each gain an electron. (That's the stage of the process we've been concentrating on in the faulty explanation.) The isolated ions become wrapped in water molecules to form hydrated ions. Note: For the next bit, if you aren't happy about enthalpy changes, you might want to explore the energetics section of Chemguide, or my chemistry calculations book . The proper explanation The table below looks at how much energy is involved in each of these changes.

To be sure that you understand the various terms: Atomisation energy This is the energy needed to produce 1 mole of isolated gaseous atoms starting from an element in its standard state (gas for chlorine, and liquid for bromine, for example - both of them as X2). For a gas like chlorine, this is simply half of the bond enthalpy (because breaking a Cl-Cl bond produces 2 chlorine atoms, not 1). For a liquid like bromine or a solid like iodine, it also includes the energy that is needed to convert them into gases. Electron affinity The first electron affinity is the energy released when 1 mole of gaseous atoms each acquire an electron to form 1 mole of gaseous 1- ions. In symbol terms: Hydration enthalpy (hydration energy) This is the energy released when 1 mole of gaseous ions dissolves in water to produce hydrated ions. -293 -481 There's quite a lot of data here to look at. Concentrate first on the final column which shows the overall heat evolved when all the other processes happen. It is calculated by adding the figures in the previous 3 columns.

You can see that the amount of heat evolved falls quite dramatically from the top to the bottom of the Group, with the biggest fall from fluorine to chlorine. Fluorine produces a lot of heat when it forms its hydrated ion, chlorine less so, and so on down the Group. Note: Don't forget that we are only talking about half of a redox reaction in each case. There will be other energy terms involving whatever the halogen is oxidising. Those changes will be overall endothermic. For example, if chlorine oxidises iodide ions to iodine, that half of the total reaction would need +481 kJ mol-1, giving an enthalpy change of reaction of (-592 + 481) = -111 kJ per mole of I- oxidised. Why is fluorine a much stronger oxidising agent than chlorine? What produces the very negative value for the enthalpy change when fluorine turns into its hydrated ions? There are two main factors. The atomisation energy of fluorine is abnormally low. This reflects the low bond enthalpy of fluorine. Note: The reason for fluorine's low bond enthalpy is described on another page.

The main reason, though, is the very high hydration enthalpy of the fluoride ion. That is because the ion is very small. There is a very strong attraction between the fluoride ions and water molecules. The stronger the attraction, the more heat is evolved when the hydrated ions are formed. Why the fall in oxidising ability from chlorine to bromine to iodine? The fall in atomisation energy between these three elements is fairly slight, and would tend to make the overall change more negative as you go down the Group. The explanation doesn't lie there! It is helpful to look at the changes in electron affinity and hydration enthalpy as you go down the Group. Using the figures from the previous table: going from +29 +42 You can see that both of these effects matter, but that the more important one - the one that changes the most - is the change in the hydration enthalpy. As you go down the Group, the ions become less attractive to water molecules as they get bigger. Although the ease with which an atom attracts an electron matters, it isn't actually as important as the hydration enthalpy of the negative ion formed. The faulty explanation misses the mark even if you restrict it to chlorine, bromine and iodine!

Warning! You really need to find out what (if any) explanation your examiners expect you to give for this. If their mark schemes (or the way they phrase their questions) suggest that they want the faulty explanation, there isn't much you can do about it. Unfortunately, there are times in exams when you have to grit your teeth and give technically wrong answers because that's what your examiners want. It shouldn't happen like this, but it does! UK A' level students should search their syllabuses, past exam papers, mark schemes and any other support material available from their Exam Board. If you haven't got any of this, you can find your Exam Board's web address by following this link. Students elsewhere should find out the equivalent information from their own sources. Questions to test your understanding If this is the first set of questions you have done, please read the introductory page before you start. You will need to use the BACK BUTTON on your browser to come back here afterwards. The Halogen Battle - Bromine, Fluoride, Chlorine Fight Against Iodine - Miss Lizzy in Iodine Dr. D.C. Jarvis wrote that Iodine is necessary for the thyroid gland’s proper performance of its work.

1 In addition to the thyroid, Iodine is needs Iodine. 2 At one time Iodine was added to bread and table salt but that is no longer the case. Historically, people living in certain areas — such as the “Goiter Belt” – have less Iodine in the soil and higher rates of Hypothyroidism. Hypothyroidism is still a prevalent problem across the entire USA 3 and around the world. When thinking about Iodine and it’s importance in the body, we need to consider its relationship to other Halogens. There are four Halogens: Fluorine Bromine Iodine Here is the part that blows my mind. There is a well-known law of halogen displacement, whereby the halogen with a higher weight displaces those at a lower weight. This means, essentially that the higher weight halogen boots out the lower weight halogen. And here is the significance. Fluorine, Chlorine and Bromine all have a higher weight than Iodine. One might think this is not such a big deal. Until we consider the use of Fluorine (Fluoride), Chlorine and Bromine in our food, water and environment.

Up until the 1960s, Iodine was used in bread. In the 1970s this changed, Iodine was removed from bread, and Bromine was used instead! Since bromine and Iodine compete for the same receptors guess what happens to the Iodine? 4 . In many towns Fluorine (Fluoride) is added to the drinking water Chlorine can also be used to treat drinking water 5 and is used in swimming pools 6 In addition to bread and baked goods, bromine is used many OTC medicines, toothpaste, mouthwashes, flame retardants, and some cosmetics This all sounds scary, a bit. But it certainly explains why Iodine supplementation can be so important. With Iodine deficiency already present, supplemental Iodine may be needed to compete with the other halogens. Notes: Folk Medicine, Dr. D.C. Jarvis, p.138 ↩ ↩ The iodine Crisis, Lynne Farrow, p. 165 ↩ Folk Medicine, Dr. D.C.

Jarvis, p.141 ↩ #.U0X1s_0uinA ↩ Reactions of chlorine, bromine and iodine with aluminium- Learn Chemistry Home > Resources > Reactions of chlorine, bromine and iodine with aluminium Reactions of chlorine, bromine and iodine with aluminium Description In this demonstration experiment aluminium reacts with three halogens, producing spectacular exothermic redox reactions. In this demonstration experiment aluminium reacts with three halogens, producing spectacular exothermic redox reactions. Subjects : Credits : This is an experiment from the Practical Chemistry project, developed by the Nuffield Foundation and the Royal Society of Chemistry. Reactions of chlorine, bromine and iodine with aluminium Demonstration In this demonstration experiment, aluminium reacts with each of three halogens, chlorine , bromine and iodine . Spectacular exothermic redox reactions occur, accompanied by flames and coloured ‘smoke’, forming the solid aluminium halides: 2Al + 3X2 → 2AlX3 (X = Cl, Br and I). Lesson organisation This is a class demonstration that shows the spectacular reactivity of three non-metals from Group 17 with a metal.

These experiments must be done in a fume cupboard as both the reactants and the products are hazardous. Teachers attempting this demonstration for the first time are strongly advised to do a trial run before doing it in front of a class. Each experiment should take about 5 minutes. Apparatus Thick chemically-resistant gloves such as marigold industrial blue nitrile Access to a fume cupboard The teacher will require: Heat resistant mat, 30 x 30 cm Watch glasses, about 10 cm diameter, 2 Reduction tube (Note 1) Aluminium foil, a few cm2 Aluminium powder (HIGHLY FLAMMABLE, CONTACT WITH WATER MAY LIBERATE HYDROGEN), 0.1 g Liquid bromine (CORROSIVE, VERY TOXIC), 1 cm3 (Note 2) Solid iodine (HARMFUL), 0.4 g Silver nitrate solution, about 0.1M, about 10 cm3 A little deionised water in a wash bottle Hydrochloric acid , 5M (CORROSIVE), about 50 cm3 Refer to Health & Safety and Technical notes section below for additional information.

Wear goggles and heavy duty gloves (due to the corrosive nature of liquid bromine) Aluminium powder - see CLEAPSS Hazcard. Liquid bromine - see CLEAPSS Hazcard. Solid iodine - see CLEAPSS Hazcard. Silver nitrate solution - see CLEAPSS Hazcard. Chlorine - see CLEAPSS Hazcard and Recipe card. Hydrochloric acid - see CLEAPSS Hazcard. 1 The reduction tube should be fitted with a one-holed rubber stopper fitted with short length of glass tubing and attached to the chlorine generator. Alternatively an 8–10 cm length of wide bore glass tubing with a stopper fitted with a short length of glass tubing at each end could be used – see diagram. 2 Wear suitable protective gloves (thick, chemically resistant) when handling liquid bromine. Have 500 cm3 of 1M solution of sodium thiosulfate available to treat any spillages on the skin. NOTE: The quantities specified in the reaction of Al + I2 should not be exceeded unless the school has an Explosives Certificate issued by the police.

Procedure Reaction of Al + Cl2 a Set up the chlorine generator in a fume cupboard. Make sure it is securely clamped. b Loosely crumple a piece of aluminium foil (10 x 5 cm) so that it will just fit inside the reduction tube and push it into the tube. Attach the tube to the generator with a short length of rubber tubing and clamp it in position at the end nearest to the generator, so that the aluminium foil can easily be heated using a Bunsen burner – see diagram. c Generate a gentle stream of chlorine by dripping the hydrochloric acid slowly on to the sodium chlorate(I) solution, and allow it to pass over the aluminium. When the green colour of the chlorine gas fills the reduction tube, start heating the aluminium foil with a Bunsen burner at the end nearest to the chlorine supply. Once the aluminium is hot, a bright glow will suddenly appear where it starts to react with chlorine. d Remove the heat. The bright glow should spread along the aluminium. If not, heat again, and increase the flow of chlorine gas.

A lot of white ‘smoke’ – aluminium chloride – is produced, some of it condensing as a white powder on the walls of the reduction tube and the rest streaming out of the hole in the end of the tube. When the reaction is over, stop the chlorine supply and remove the heat. e When the reduction tube has cooled down, disconnect it and, still in the fume cupboard, scrape a little of the white powder into a test-tube. Add a little deionised/distilled water to the solid to dissolve it. Care: the reaction between anhydrous aluminium chloride and water can be quite vigorous – an audible hiss can often be heard - see Teaching Notes. f Test a drop of the solution with Universal indicator paper. It is strongly acidic. Test the remainder with a little silver nitrate solution. A white precipitate forms, showing the presence of chloride ions. Reaction of Al + Br2 a Tear or cut some aluminium foil into several small pieces about 2 x 2 mm in size. Carefully pour 1 cm3 of liquid bromine onto a watchglass on a heat resistant mat in a fume cupboard. Sprinkle a few pieces of aluminium foil on to the surface of the bromine.

Within a minute there are flashes of flame and a white ‘smoke’ of aluminium bromide is formed, together the orange vapour of bromine. Carefully hold another watchglass over the reaction to condense some of the ‘smoke’ on to its surface as a solid. b Wash any aluminium bromide collected in this way off the watchglass into a test-tube using a little deionised water (Care: see Teaching Notes below). Test the solution with indicator paper and silver nitrate solution as above. The solution is acidic and a cream precipitate of silver bromide is formed. Reaction of Al + I2 a Weigh out 0.4 g of iodine, which should be dry, and grind it to a powder in a fume cupboard, using a mortar and pestle. Place the powdered iodine on a filter paper on a dry heat resistant mat and add 0.1 g of aluminium powder to it. Mix the two solids together in the fume cupboard using a wooden splint – do not grind them together. When they are thoroughly mixed, pour the mixture into a heap on the heat resistant mat or in a metal tray, such as a tin lid, positioned in the middle of the fume cupboard.

b To start the reaction, use a teat pipette to place a few drops of water on the mixture. After a time lag, the water begins to steam and clouds of purple iodine vapour are given off, indicating that an exothermic reaction has started. After a few more seconds sparks are given off and the mixture bursts into flame. When the reaction subsides, a white residue of aluminium iodide remains. Scrape a little of this into a test-tube (Care: see Teaching Notes below), add some deionised water and filter if necessary. Test the solution with indicator paper and silver nitrate solution as above. The solution is acidic and a yellow precipitate indicates the presence of iodide ions. Teaching notes These reactions make quite spectacular demonstrations in themselves, the bromine + aluminium reaction even more so in a partly darkened room. Classroom management in semi-darkness (You need to plan carefully for such lessons. Ensure that students are clear about what they need to do during such activities). The demonstrations can be used to show the reaction between reactive non-metallic elements and a fairly reactive metal to form compounds, or as part of the study of the reactions of the Group 17 elements.

Here the apparent order of reactivity is not that predicted from their position in the Group (that is chlorine → bromine → iodine). This is because of the different physical states of the three halogens, and the different surface area of the aluminium as a powder or foil. This can be used to make an important point about ‘fair’ comparisons of reactivity. These reactions also serve to show that aluminium is in fact a more reactive metal than it appears in everyday use. The protective oxide layer of aluminium has to be penetrated by the halogens before the reactions can start, hence the delays, and the need for water to assist the two solid elements getting into contact, in the case of aluminium and iodine. The clouds of iodine vapour released when aluminium and iodine react can stain the inside of a fume cupboard. Teachers may prefer to demonstrate this reaction outdoors, if possible. The anhydrous aluminium halides are vigorously hydrolysed (sometimes violently if freshly prepared and hot, as here) by water, giving off fumes of a hydrogen halide and a forming an acidic solution of the aluminium salt.

To dispose of the solid residues, allow them to cool completely before adding in small amounts to 1 M sodium carbonate solution in a fume cupboard. Wait until the reaction has subsided before adding more solid. Dispose of the resultant slurry with plenty of water. Health & Safety checked, January 2016 Credits BBC - GCSE Bitesize: Physical properties of halogens Next The elements in Group 7 of the Periodic Table are called the halogens. They include chlorine, bromine and iodine. The halogens are diatomic - this means they exist as molecules, each with a pair of atoms. Chlorine molecules have the formula Cl2, bromine Br2 and iodine I2. Physical properties of halogens The halogens show trends in physical properties down the group. Melting point and boiling point The halogens have low melting points and boiling points. This is a typical property of non-metals. Fluorine has the lowest melting point and boiling point. The melting points and boiling points then increase as you go down the group. Melting and boiling points of Group 7 elements State at room temperature Room temperature is usually taken as being 25°C.

At this temperature, fluorine and chlorine are gases, bromine is a liquid, and iodine and astatine are solids. There is therefore a trend in state from gas to liquid to solid down the group. Colour The halogens become darker as you go down the group. Fluorine is very pale yellow, chlorine is yellow-green, and bromine is red-brown. Iodine crystals are shiny purple - but easily turn into a dark purple vapour when they are warmed up. Predictions When we can see a trend in the properties of some of the elements in a group, it is possible to predict the properties of other elements in that group. Astatine is below iodine in Group 7. The colour of these elements gets darker as you go down the group. Iodine is purple, and, as we would expect, astatine is black. Page: Atomic and physical properties of Periodic Table Group 7 (the halogens) ATOMIC AND PHYSICAL PROPERTIES OF THE GROUP 7 ELEMENTS (THE HALOGENS) This page explores the trends in some atomic and physical properties of the Group 7 elements (the halogens) - fluorine, chlorine, bromine and iodine.

You will find separate sections below covering the trends in atomic radius, electronegativity, electron affinity, melting and boiling points, and solubility. There is also a section on the bond enthalpies (strengths) of halogen-halogen bonds (for example, Cl-Cl) and of hydrogen-halogen bonds (e.g. H-Cl) Even if you aren't currently interested in all these things, it would probably pay you to read the whole page. The same ideas tend to recur throughout the atomic properties, and you may find that earlier explanations help to you understand later ones. Trends in Atomic Radius Note: You will find atomic radius covered in detail in another part of this site. If you choose to follow this link, use the BACK button on your browser to return quickly to this page. You can see that the atomic radius increases as you go down the Group. Explaining the increase in atomic radius The radius of an atom is governed by the number of layers of electrons around the nucleus the pull the outer electrons feel from the nucleus. Compare fluorine and chlorine: Cl 2,8,7 In each case, the outer electrons feel a net pull of 7+ from the nucleus.

The positive charge on the nucleus is cut down by the negativeness of the inner electrons. This is equally true for all the other atoms in Group 7. The outer electrons always feel a net pull of 7+ from the centre. The only factor which is going to affect the size of the atom is therefore the number of layers of inner electrons which have to be fitted in around the atom. Obviously, the more layers of electrons you have, the more space they will take up - electrons repel each other. That means that the atoms are bound to get bigger as you go down the Group. Trends in Electronegativity Electronegativity is a measure of the tendency of an atom to attract a bonding pair of electrons. It is usually measured on the Pauling scale, on which the most electronegative element (fluorine) is given an electronegativity of 4.0. Note: You will find electronegativity covered in detail in another part of this site. If you choose to follow this link, use the BACK button on your browser to return quickly to this page. Notice that electronegativity falls as you go down the Group. The atoms become less good at attracting bonding pairs of electrons.

Explaining the decrease in electronegativity This is easily shown using simple dots-and-crosses diagrams for hydrogen fluoride and hydrogen chloride. The bonding pair of electrons between the hydrogen and the halogen feels the same net pull of 7+ from both the fluorine and the chlorine. (This is exactly the same sort of argument as you have seen in the atomic radius section above.) However, in the chlorine case, the nucleus is further away from that bonding pair. That means that it won't be as strongly attracted as in the fluorine case. The larger pull from the closer fluorine nucleus is why fluorine is more electronegative than chlorine is. Summarising the trend down the Group As the halogen atoms get bigger, any bonding pair gets further and further away from the halogen nucleus, and so is less strongly attracted towards it. In other words, as you go down the Group, the elements become less electronegative. Trends in First Electron Affinity Defining first electron affinity The first electron affinity is the energy released when 1 mole of gaseous atoms each acquire an electron to form 1 mole of gaseous 1- ions. This is more easily seen in symbol terms.

It is the energy released (per mole of X) when this change happens. First electron affinities have negative values. For example, the first electron affinity of chlorine is -349 kJ mol-1. By convention, the negative sign shows a release of energy. The first electron affinities of the Group 7 elements Note: You will find electron affinity covered in detail in another part of this site. The current page duplicates much of that material, but you might like to read it again in different words. If you choose to follow this link, use the BACK button on your browser to return quickly to this page. Notice that the trend down the Group isn't tidy. The tendency is for the electron affinities to decrease (in the sense that less heat is given out), but the fluorine value is out of line. The electron affinity is a measure of the attraction between the incoming electron and the nucleus. The higher the attraction, the higher the electron affinity. In the bigger atom, the attraction from the more positive nucleus is offset by the additional screening electrons, so each incoming electron feels the effect of a net 7+ charges from the centre - exactly as when you are thinking about atomic radius or electronegativity.

As the atom gets bigger, the incoming electron is further from the nucleus and so feels less attraction. The electron affinity therefore falls as you go down the Group. But what about fluorine? That is a very small atom, with the incoming electron quite close to the nucleus. Why isn't its electron affinity bigger than chlorine's? There is another effect operating. When the new electron comes into the atom, it is entering a region of space already very negatively charged because of the existing electrons. There is bound to be some repulsion, offsetting some of the attraction from the nucleus. In the case of fluorine, because the atom is very small, the existing electron density is very high. That means that the extra repulsion is particularly great and lessens the attraction from the nucleus enough to lower the electron affinity below that of chlorine. Trends in Melting Point and Boiling Point You will see that both melting points and boiling points rise as you go down the Group. If you explore the graphs, you will find that fluorine and chlorine are gases at room temperature, bromine is a liquid and iodine a solid. Nothing very surprising there!

Explaining the trends in melting point and boiling point All of the halogens exist as diatomic molecules - F2, Cl2, and so on. The intermolecular attractions between one molecule and its neighbours are van der Waals dispersion forces. Note: If you aren't sure about van der Waals dispersion forces , you will find them covered in detail in another part of this site. You won't understand the next bit unless you are happy about dispersion forces and how they vary with the size of the molecule. Use the BACK button on your browser to return quickly to this page. As the molecules get bigger there are obviously more electrons which can move around and set up the temporary dipoles which create these attractions. The stronger intermolecular attractions as the molecules get bigger means that you have to supply more heat energy to turn them into either a liquid or a gas - and so their melting and boiling points rise. Solubilities Solubility in water Fluorine reacts violently with water to give hydrogen fluoride gas (or a solution of hydrofluoric acid) and a mixture of oxygen and ozone. So thinking about its solubility is pointless.

Chlorine, bromine and iodine all dissolve in water to some extent, but there is no pattern in this. The following table shows the solubility of the three elements in water at 25°C. chlorine 0.091 bromine0.21 iodine0.0013 Note: These figures come from page 476 of Advanced Inorganic Chemistry (third edition) by Cotton and Wilkinson. Chlorine solution in water is pale green. Bromine solution in water is anything from yellow to dark orange-red depending on how concentrated it is. Iodine solution in water is very pale brown. Chlorine reacts with water to some extent to give a mixture of hydrochloric acid and chloric(I) acid - also known as hypochlorous acid. The reaction is reversible, and at any one time only about a third of the chlorine molecules have actually reacted. You will sometimes find the chloric(I) acid written as HOCl. That represents the way the atoms are actually joined together. Bromine and iodine do something similar, but to a much lesser extent. In both cases, about 99.5% of the halogen remains as unreacted molecules.

The solubility of iodine in potassium iodide solution Although iodine is only faintly soluble in water, it does dissolve freely in potassium iodide solution to give a dark red-brown solution. There is a reversible reaction between iodine molecules and iodide ions to give I3- ions. These are responsible for the colour. In the lab, iodine is often produced by oxidation of a solution containing iodide ions, so this colour is actually quite familiar. As long as there are any excess iodide ions present, the iodine will react with them to make the I3- ions. Once the iodide ions have all reacted, the iodine is precipitated as a dark grey solid, because there isn't anything left for it to react with to keep it in solution. Solubility in hexane The halogens are much more soluble in organic solvents like hexane than they are in water. Both hexane and the halogens are non-polar molecules attracted to each other by van der Waals dispersion forces. That means that the attractions broken (between hexane molecules and between halogen molecules) are similar to the new attractions made when the two substances mix. The colours of the solutions formed are much what you would expect.

Solutions of iodine in organic solvents tend to be pinky-purple colour. Bond enthalpies (bond energies or bond strengths) Bond enthalpy is the heat needed to break one mole of a covalent bond to produce individual atoms, starting from the original substance in the gas state, and ending with gaseous atoms. So for chlorine, Cl2(g), it is the heat energy needed to carry out this change per mole of bond: For bromine, the reaction is still from gaseous bromine molecules to separate gaseous atoms. Bond enthalpy in the halogens, X2(g) A covalent bond works because the bonding pair is attracted to both the nuclei at either side of it. It is that attraction which holds the molecule together. The size of the attraction will depend, amongst other things, on the distance from the bonding pair to the two nuclei. As with all halogens, the bonding pair will feel a net pull of 7+ from both ends of the bond - the charge on the nucleus offset by the inner electrons. That will still be the same whatever the size of the halogen atoms.

As the atoms get bigger, the bonding pair gets further from the nuclei and so you would expect the strength of the bond to fall. So . . . are the actual bond enthalpies in line with this prediction? The bond enthalpies of the Cl-Cl, Br-Br and I-I bonds fall just as you would expect, but the F-F bond is way out of line! Because fluorine atoms are so small, you might expect a very strong bond - in fact, it is remarkably weak. There must be another factor at work as well. As well as the bonding pair of electrons between the two atoms, each atom has 3 non-bonding pairs of electrons in the outer level - lone pairs. Where the bond gets very short (as in F-F), the lone pairs on the two atoms get close enough together to set up a significant amount of repulsion. In the case of fluorine, this repulsion is great enough to counteract quite a lot of the attraction between the bonding pair and the two nuclei. This obviously weakens the bond. Bond enthalpies in the hydrogen halides, HX(g) Where the halogen atom is attached to a hydrogen atom, this effect doesn't happen.

There are no lone pairs on a hydrogen atom! As the halogen atom gets bigger, the bonding pair gets more and more distant from the nucleus. The attraction is less, and the bond gets weaker - exactly what is shown by the data. There is nothing complicated happening in this case. This is important in the thermal stability of the hydrogen halides - how easily they are broken up into hydrogen and the halogen on heating. Hydrogen fluoride and hydrogen chloride are very stable to heat. They don't split up into hydrogen and fluorine or chlorine again if heated to any normal lab temperature. Hydrogen bromide splits slightly into hydrogen and bromine on heating, and hydrogen iodide splits to an even greater extent. As the bonds get weaker, they are more easily broken. Note: Breaking the hydrogen-halogen bond is only one of the steps in the overall reaction, of course. You don't end up with hydrogen atoms and halogen atoms - you get diatomic molecules, H2 and X2. If you have done some energetics calculations, it would be a useful exercise to use bond enthalpies and atomisation enthalpies to calculate the overall enthalpy changes for the decomposition of the hydrogen halides.

(Remember that bond enthalpies only apply to substances in the gas state, and bromine and iodine would end up as liquid and solid respectively. Using atomisation enthalpies for the halogens avoids this problem. If you don't understand what I am talking about, you don't yet have enough knowledge to be able to do this.) Questions to test your understanding If this is the first set of questions you have done, please read the introductory page before you start. You will need to use the BACK BUTTON on your browser to come back here afterwards. Iodine Protects Against Fluoride Toxicity Iodine Protects Against Fluoride Toxicity Published on November 15, 2011 10:21am on November 15th, 2011 8:44pm on December 10th, 2016 share on Print this The toxicity of modern life is impacting iodine levels and in the countries that fluoridate their water this impact is maximized. It is well known that the toxic halides: fluoride and bromide, having structure similar to iodine, can competitively inhibit iodine absorption and binding in the body.

All the halogens use the same receptors in the body so fluoride’s danger for people is centered in great part on this fact. Americans and Brazilians, who are more exposed to fluoride than other populations, have a desperate need for more iodine. Taking iodine in its nascent form is not only the best way to increase iodine levels in the safest and most effective way possible for adults and children whose thyroids are already compromised, but it will also greatly aid in ridding the body of dangerous fluoride, bromide, chlorine, perchlorates and heavy metals. Is fluoride bad for your health ? In our age of increasing radioactivity and toxic poisoning specifically with fluoride, [1] chlorine, bromide, and even mercury, iodine is a necessary mineral. Learn a system of medicine that is safe, simple and affordable! Discover Dr Sircus Natural Allopathic Medicine » Iodine is extremely important since the cells need it to regulate their metabolism. Without it, people are known to suffer from swollen glands in the throat, thyroid diseases, increased fluoride toxicity, decreased fertility rates, increased infant mortality rates, and (with severe deficiency) mental retardation.

It has been theorized that iodine deficiency is a causal factor of ADHD in babies of iodine-deficient mothers. Iodine intake immediately increases the excretion of bromide, fluoride, and some heavy metals including mercury and lead. Bromide and fluoride are not removed by any other chelator or detoxifying technique. Dr. Kenezy Gyula Korhaz states that iodine chelates heavy metals such as mercury, lead, cadmium, aluminum, and halogens such as fluoride and bromide, thus decreasing their iodine-inhibiting effects, [2] especially of the halogens. Iodine has the highest atomic weight of all the common halogens (126.9). Iodine is the only option when it comes to removing these toxic haloids from the thyroid and even the pineal gland where fluoride concentrates, especially when there is a deficiency of iodine in the body. “The human pineal gland contains the highest concentration of fluoride in the body. Fluoride is associated with depressed pineal melatonin synthesis and this depression increases one’s chance of cancer.” Dr.

David Brownstein says that fluoride inhibits the ability of the thyroid gland to concentrate iodine and research has shown that fluoride is much more toxic to the body when there is iodine deficiency present. Brownstein says that after only one dose of iodine, the excretion of fluoride increases by 78%. [3] On January 7, 2011, the US Department of Health & Human Services (HHS) proposed lowering the recommended level used in the water fluoridation program to 0.7 ppm, because of the very high incidence of dental fluorosis among American children. An amazing 41% of ALL American children aged 12-15 are now impacted by this condition. Sodium fluoride is commonly used as a rat poison. Globalists and eugenicists have decided to add it to water supplies with the message to the public that it is good for teeth, despite warnings from the ADA stating that young children risk a disease called dental fluorosis.

After hailing water fluoridation as one of the 10 greatest health achievements of the 20th Century (CDC) , the government is calling for a reduction in the amount of fluoride it adds to public water supplies, citing its negative effect on teeth when promotion of healthy teeth is the basic reason given for adding fluoride to the water. An August 2006 Chinese study found that fluoride in drinking water damages children’s liver and kidney functions. One of the strongest physiological effects of fluorides in drinking water (e.g. hydrofluorosilicic acid) is on the kidney, a point to consider in light of increased rates of kidney failure during recent decades. [4] Kidney disease markedly increases an individual’s susceptibility to fluoride toxicity . In healthy adults, the kidneys are able to excrete approximately 50% of an ingested dose of fluoride. However, in adults with kidney disease, the kidneys may excrete as little as 10-20%, and young children may only excrete 15% of an ingested dose—thus increasing the body burden of fluoride and increasing an individual’s susceptibility to fluoride poisoning (e.g. renal osteodystrophy ).

Scientific evidence over the past 50 plus years has shown that sodium fluoride shortens our life span, promotes various cancers and mental disturbances, and most importantly, makes humans stupid, docile, and subservient, all in one neat little package. A Scientific American study “concluded that fluoride can subtly alter endocrine function, especially in the thyroid.” The National Research Council of the National Academies in a 2006 report on page 266 said, “In summary, evidence of several types indicates that fluoride affects normal endocrine function or response; the effects of the fluoride-induced changes vary in degree and kind in different individuals. Fluoride is therefore an endocrine disruptor in the broad sense of altering normal endocrine function.” Halogen Displacement The mechanism behind “halogen displacement” was probably best described by J. C. Jarvis, M.D. (Folk Medicine, Henry Holt & Co., 1958, HB, p. 136), who wrote: “The clinical activity of any one of these four halogens is in inverse proportion to its atomic weight.

This means that any one of the four can displace the element with a higher atomic weight, but cannot displace an element with a lower atomic weight. For example, flourine can displace chlorine, bromine, and iodine because fluorine has a lower atomic weight than the other three. Similarly, chlorine can displace bromine and iodine because they both have a higher atomic weight.” Likewise, bromine can displace iodine from the body because iodine has a higher atomic weight. A reverse order is not possible. European doctors used fluoride as a thyroid-suppressing medication for patients with HYPER-thyroidism (over-active thyroid). Fluoride was utilized because it was found to be effective at reducing the activity of the thyroid gland—even at doses as low as 2 mg/day. The regular use of iodine will go a very long way toward mitigating the damages done in our bodies by the fluoride that we are exposed to. The Nascent form is the easiest way of increasing iodine levels for adults and children whose thyroids are already compromised by fluoride. Nascent Iodine is the atomic form and it is special.

Nascent has the advantage of being in the I¹ atomic (as opposed to I² or I³ molecular forms) form meaning that there is no digestion, no breaking down of molecules of iodine. It is already broken down through an electromagnetic process. For heavy lifting when needing iodine in large quantities for transdermal application (painting the breasts daily with iodine for breast cancer is one example) I recommend the age old Lugol’s formula . Liquid forms of minerals always seem better than solid pill forms because the body has an easier time digesting and absorbing liquid forms with the nascent form ready to be utilized instantly in any way the body needs. Some tissues utilize iodide and others iodine and the thyroid always needs atomic iodine to make T3 and T4 metabolic thyroid hormones. [1] Fluoride is associated with cancer and it also accumulates in the thyroid as well as the pineal gland, an important hormone control center. Dr. Jennifer Luke found out that the pineal gland which produces serotonin and melatonin was also a calcifying tissue, like the teeth and the bones, so she hypothesized it would concentrate fluoride to very high levels.

Luke had 11 cadavers analyzed in the UK and found very high levels of fluoride in the calcium hydroxy apatite crystals produced by the gland. The average was 9000 ppm and went as high as 21,000 in one case. These levels are equal to or higher than fluoride levels in the bones of people suffering from skeletal fluorosis. Luke hypothesizes that one of the four enzymes needed to convert the amino acid tryptophan (from the diet) into melatonin is being inhibited by fluoride. Melatonin is responsible for regulating all kinds of activities including the onset of puberty. It is thought that the fall of melatonin levels acts like a biological clock and triggers the onset of puberty. In her gerbil study she found that the high fluoride treated animals were reaching puberty earlier than the low fluoride ones. Considering the seriousness of a possible interference by fluoride on a growing child’s pineal gland (and for that matter, elderly pineal glands) underlines the need for higher iodine intake to increase fluoride elimination. [2] Sticht, G., Käferstein, H., Bromine.

In Handbook on Toxicity of Inorganic Compounds – Seiler HG and Sigel, H Editors, Marcel Dekker Inc, 143-151, 1988. [3] David Brownstein, Iodine, Why You Need It, Why You Can’t Live Without It;

Question: What are chlorine, bromine and iodine?
Answer: Halogens