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. What are the Halogens in the Periodic Table?

- Quora Quora Daniel James Berger , PhD in organic chemistry The elements in Group 17 of the periodic table ( The Old fashioned name for Group 17 is Group VII ) are known collectively as The Halogens. They are the elements Fluorine, ( atomic number 9 ), Chlorine, ( atomic number 17 ), Bromine ( atomic number 35 ), Iodine ( atomic number 53 ), Astatine ( atomic number 85 ), and a newly discovered element not yet given a name, that has an atomic number of 117, ( Discovered in 2010). However both Astatine and Element No. 117 are too radioactive to be made in quantities that can be weighed. They are man made elements that have very short half lives; so decay into other elements before one has the chance to make any of their compounds. So in most laboratories Halogen chemistry is essentially about the study of Fluorine, Chlorine, Bromine & Iodine. The word Halogen was coined in 1842 by the Swedish Chemist Berzelius ( 1779 - 1848 ).

It is a made up Swedish word which literally means "Salt producer" in Swedish. A name in keeping with the reactive nature of these elements. Lots of Halogen salts are known. Eg. Sodium chloride. Potassium fluoride. Sodium Bromide. These elements are all members of the same group, because all of their atoms have the same number of electrons in their outermost shell or orbit. Group 17 elements all have 7 electrons in their outermost orbit. ​ ​The Periodic Table highlighting Group 17. Fluorine As mentioned previously fluorine has an atomic number of 9. This means that an atom of fluorine will have in addition to 9 protons in the nucleus of its atom, it will have a total of 9 electrons orbiting its nucleus. An individual proton always has a charge of +1 and an individual electron always has a charge of -1. We have 9 protons in an atom of fluorine each having a charge of +1, plus 9 electrons in that same atom orbiting the nucleus of that atom. ( +1 X 9 ) + (-1 X 9 ) = 0. So the atom is electrically neutral.

This is true for all neutral atoms of fluorine. The 9 electrons are arranged in 2 shells or orbit. 2 of the 9 electrons orbit in the 1st shell or orbit. The remaining 7 electrons orbit in a larger shell. The 7 electrons in the outermost orbit of all Halogens are the electrons that participate in chemical reactions. During chemical reactions the number of electrons in the outermost orbit can go up or down. Up to a max of 8, and down to 0 ( in theory ). Compare with Na+1 or with Li+1 ​ Written Feb 7, 2016 The halogens are the family of chemical elements that includes fluorine (atomic symbol F), chlorine (Cl), bromine (Br), iodine (I), and astatine (At). The halogens make up Group VIIA of the Periodic Table of the elements. Elemental halogens are diatomic molecules. However, due to their high reactivity, the halogens are never found in nature in native form. The family name means "salt-forming," from the Greek for salt, "hals", and for generating "genes".

The salinity of the oceans on Earth is due in large part to such halogen salts (halides) as sodium chloride (NaCl) and potassium iodide (KI). Halogens display physical and chemical properties typical of nonmetals. They have relatively low melting and boiling points that increase steadily down the group. Near room temperature, the halogens span all of the physical states: Fluorine and chlorine are gases, bromine is a liquid, and iodine is a solid. All of the elements are colored, with the color becoming more intense moving down the group. Fluorine gas is pale yellow, and chlorine gas is a yellowish green. Liquid bromine and its vapors are brownish red. Solid iodine appears as shiny, dark gray crystals, and the vapors are a deep purple. The halogens are poor thermal and electrical conductors in all phases, and as solids they are brittle and crumbly. The halogens have distinctive, unpleasant odors, will burn exposed flesh, and are toxic. The neutral atoms of the halogens possess seven outer electrons. An additional electron can be added to halogen atoms to form singly charged negative ions. These ions have a closed outer-shell configuration.

Electronegativity is a measure of the ability of an atom of one element to remove an electron from an atom of another element. As a group, the halogens are among the most electronegative elements. Fluorine has the highest electronegativity of all the elements. Halogens are so reactive that all the elements except helium and neon have been found to react with at least one of the halogens. Fluorine is always assigned a formal oxidation number of –1, whereas the other halogens can exhibit a range of oxidation numbers 526 Views · View Upvotes · Answer requested by Halogens - Chemistry Encyclopedia - uses, elements, gas, number, name, symbol, salt, atom Halogens - Chemistry Encyclopedia Halogens Photo by: The halogens are the family of chemical elements that includes fluorine (atomic symbol F), chlorine (Cl), bromine (Br), iodine (I), and astatine (At). The halogens make up Group VIIA of the Periodic Table of the elements. Elemental halogens are diatomic molecules. However, due to their high reactivity, the halogens are never found in nature in native form.

The family name means "salt-forming," from the Greek for salt, halos, and for generating genes. The salinity of the oceans on Earth is due in large part to such halogen salts (halides) as sodium chloride (NaCl) and potassium iodide (KI). Halogens display physical and chemical properties typical of nonmetals. They have relatively low melting and boiling points that increase steadily down the group. Near room temperature, the halogens span all of the physical states: Fluorine and chlorine are gases, bromine is a liquid, and iodine is a solid. All of the elements are colored, with the color becoming more intense moving down the group. Fluorine gas is pale yellow, and chlorine gas is a yellowish green. Liquid bromine and its vapors are brownish red. Solid iodine appears as shiny, dark gray crystals, and the vapors are a deep purple. The halogens are poor thermal and electrical conductors in all phases, and as solids they are brittle and crumbly. The halogens have distinctive, unpleasant odors, will burn exposed flesh, and are toxic. The neutral atoms of the halogens possess seven outer electrons.

An additional electron can be added to halogen atoms to form singly charged negative ions. These ions have a closed outer-shell configuration. Electronegativity is a measure of the ability of an atom of one element to remove an electron from an atom of another element. As a group, the halogens are among the most electronegative elements. Fluorine has the highest electronegativity of all the elements. Halogens are so reactive that all the elements except helium and neon have been found to react with at least one of the halogens. Fluorine is always assigned a formal oxidation number of –1, whereas the other halogens can exhibit a range of oxidation numbers. The ability of halogens to form chemical compounds with all metals and most nonmetals has led to a wide variety of uses for these elements. Chlorine is used as a bleach and a disinfectant. Iodine has been used as a topical microbicide. Iodine and bromine are added to halogen lamps to lengthen Chlorine gas belongs to the halogen chemical family. the life of the filament and prevent darkening of the bulb. Chloride and iodide are essential dietary minerals for humans.

Organic compounds containing halogens are used as fire-retardants (halons), as refrigerants (Freons), and in nonstick coatings (Teflon). Silver bromide and silver iodide have been used in photographic emulsions since the early days of photography. Many halogenated compounds are toxic. A well-known example is DDT (dichlorodiphenyltrichloroethane), once a widely applied pesticide that was banned in the United States after severe environmental effects were discovered. All known isotopes of astatine are radioactive, with the longest-lived isotope having a half-life of about eight hours. Relatively little is known of the physical and chemical properties of astatine. However, it is predicted to have properties similar to iodine. SEE ALSO John Michael Nicovich Bibliography Lide, David R., ed. (2003). The CRC Handbook of Chemistry and Physics , 84th edition. Boca Raton, FL: CRC Press. Internet Resources The Halogens (Group 17) About Watch and Favorite Watch Watching this resources will notify you when proposed changes or new versions are created so you can keep track of improvements that have been made.

Favorite Favoriting this resource allows you to save it in the “My Resources” tab of your account. There, you can easily access this resource later when you’re ready to customize it or assign it to your students. The Halogens (Group 17) The halogens are a series of highly reactive, nonmetal elements from Group 17 of the periodic table. Learning Objective Describe the general properties of the halogen elements. Key Points All halogens are electronegative: they gain electrons very quickly, making them the most reactive of all chemical elements . Halogens easily dissociate into atomic particles and can combine with surrounding elements to form compounds . Halogens are generally poisonous; each has a different level of toxicity. Halogens are poor conductors of heat and electricity. Halogens can exist in all three classical states of matter: solid , liquid , and gas . Terms Full Text Overview The halogens are nonmetallic elements in Group 17 of the periodic table. They include fluorine, chlorine, bromine, iodine, and astatine. Ununseptium, which is not a naturally occurring element, is also often considered a halogen.

As typical nonmetals, they have low melting and boiling points . A unique characteristic of halogens is that they are the only element group that includes elements that belong to all three classical states of matter: solid, liquid, and gas at room temperature . The halogens, as a group, are extremely reactive. In their elemental form, they are found as diatomic molecules . However, due to their extremely high reactivity (electronegativity), they are not found in their elemental form in any natural environment on Earth. Safety and Storage for Elemental Halogens Halogens are generally toxic , or even lethal, to biological organisms in their elemental forms. They tend to combine readily, and explosively, with metals , particularly with alkali metals . Fluorine is one of the most reactive elements in existence, attacking otherwise inert materials, such as glass , and forming compounds with the heavier noble gases . It is a corrosive and highly toxic gas. The reactivity of fluorine is such that if it is used or stored in laboratory glassware, it can react with glass in the presence of small amounts of water to form silicon tetrafluoride (SiF4).

Therefore, fluorine must be handled with substances such as Teflon (which is itself an organofluorine compound), extremely dry glass, or metals such as copper or steel that form a protective layer of fluoride on their surface. The high reactivity of fluorine also means that once it does react with something, it bonds with it so strongly that the resulting molecule is very inert and non-reactive to anything else. For example, Teflon is a polymer of fluorine bonded with carbon. The halogens are highly reactive, diatomic molecules. Reactivity of Halogens The halogens all form binary compounds with hydrogen, known as the hydrogen halides : hydrogen fluoride (HF), hydrogen chloride (HCl), hydrogen bromide (HBr), hydrogen iodide (HI), and hydrogen astatide (HAt). All of these are strong acids when dissolved in water, with the exception of HF. However, hydrofluoric acid is quite destructive to animal tissue. When in aqueous solution , the hydrogen halides are known as hydrohalic acids. The names of these acids are as follows: hydrofluoric acid, hydrochloric acid, hydrobromic acid, and hydroiodic acid.

All of these acids must be handled with great care because they are extremely dangerous. Some of these acids are also widely used in industry. The halogens can also react with each other to form interhalogen compounds. Reactivity of Halogens and Water Fluorine reacts vigorously with water to produce oxygen (O2) and hydrofluoric acid (HF): $2F_2 (g) + 2H_2O (l) \rightarrow O_2 (g) + 4HF (aq)$ Chlorine has a maximum solubility of about 7.1 g Cl2/kg H2O at ambient temperature (21 °C).

Dissolved chlorine reacts to form hydrochloric acid (HCl) and hypochlorous acid (HClO), a solution that can be used as a disinfectant or bleach: $Cl_2(g) + H_2O(l) \rightarrow HCl(aq) + HClO(aq)$ Bromine has a solubility of 3.41 g/100 g H2O, but it slowly reacts to form hydrogen bromide (HBr) and hypobromous acid (HBrO): $Br_2(g) + H_2O(l) \rightarrow HBr(aq) + HBrO(aq)$ Iodine, however, is minimally soluble in water (0.03 g/100 g H2O at 20 °C) and does not react with it. However, iodine will form an aqueous solution in the presence of iodide ions , such as with the addition of potassium iodide (KI), because the triiodide ion (I3-) is formed. BBC - GCSE Bitesize: Halogens The Group 7 elements are also known as the halogens.

They include fluorine, chlorine, bromine and iodine, which all have seven electrons in their outer shell. In a displacement reaction a less reactive element is removed and replaced by a more reactive element. Halogens The periodic table Group 7, also known as the halogens, includes fluorine, chlorine, bromine and iodine. Chlorine is used to sterilise the drinking water supply and in the manufacture of many important chemicals such as pesticides and plastics. Iodine can be used to sterilise wounds. An important compound of chlorine is sodium chloride or common salt. It's used to preserve food, as flavouring and to manufacture chlorine. The properties of the halogens The halogens have similar properties because they all have seven electrons in the outer shell. They all react very vigorously with alkali metals. Halogen Halogens | Passnownow Chemistry SS 2 Week 1 Topic: Halogens Introduction The halogens are the collective name given to the elements in group VII of the Periodic Table. There are five halogens; fluorine, chlorine, bromine, iodine and astatine. Astatine is very radioactive and cannot exist for more than a few microseconds before decaying.

We will thus be concerned with the chemistry of fluorine, chlorine, bromine and iodine. All these elements are most commonly found in the -1 oxidation state, as X– ions. These are known as halide ions. Halide is the name given to the ion of halogens. Table below shows the corresponding halide of the halogen Halogen Density of element increases (going down the group) – increase in atomic mass Chemical properties The halogens can act as oxidising agents by gaining electrons to form halide ions. X2 + 2e– → 2X– The oxidising ability decreases down the group with fluorine being the strongest oxidising agent. In simple terms the formation of X– occurs in two steps. A strong oxidising agent will readily form X–. It will therefore, have a low bond energy and a high electron affinity. Both bond energies and electron affinities decrease down the group as the nuclear pull on the outer electrons decrease. The change in electron affinities is most significant and so oxidising ability decreases down the group as the nuclear pull on the outer electrons decreases down the group. Fluorine has anomalous properties due to its small size.

It has a lower electron affinity than Chlorine due to electron repulsions in the overcrowded outer shell. However, it is a stronger oxidising agent than Chlorine due to its unusually low bond energy. This is due to the repulsions between outer electrons in the F2 molecule. Halide ions act as reducing agents by losing electrons. 2X– – 2e– → X2 The reducing ability increases from F– to I–. In all their reactions the halogens act as oxidising agents. Hence, their reactivity decreases down the group. Reaction with water F2 oxidises H2O to O2 gas in a very exothermic reaction. 2F2(g) + 2H2O(l) → O2(g) + 4HF(g) Cl2 dissolves in H2O and some hydrolysis occurs. A yellow solution of ‘chlorine water’ is formed which is a mixture of two acids. No O2 is evolved. Cl2(g) + H2O(l) → HCl(aq) + HOCl(aq) Br is only slightly soluble in H2O and there is less hydrolysis.

Br2(l) + H2O(l) → HBr(aq) + HOBr(aq) I2 is virtually insoluble in H2O. It is however soluble in KI solution due to the formation of the triiodide anion. I2(s) + I–(aq) → I3–(aq) Note: All halogens are more soluble in non-polar solvents such as CCl4. Cl2 gives a colourless solution. Br2 a red solution and I2 a violet one. Displacement reactions If a more reactive halogen is placed into a solution containing a less reactive halide a displacement reaction is seen. For example: Cl2(g) + KI(aq) → KCl(aq) + I2(s) F2 will displace Cl– Br– I– Cl2 will displace Br– I– Br2 will displace I– Reaction with alkali ( NaOH solution) This reaction depends on the conditions: Cold NaOH solution. 2OH–(aq) + Cl2(g) → Cl–(aq) + ClO–(aq) + H2O(l) Hot NaOH solution.

6OH–(aq) + Cl2(g) → 5Cl–(aq) + ClO3– + 3H2O(l) In hot solution the ClO– ions are disproportionating 3ClO–(aq) → 2Cl–(aq) + ClO3–(aq) Other Reactions of the Halogens Reaction with hydrogen H2 Halogens readily combine with hydrogen to form the hydrogen halides which are colourless gaseous covalent molecules e.g. hydrogen + chlorine ==> hydrogen chloride H2(g) + Cl2(g) ==> 2HCl(g) The hydrogen halides dissolve in water to form very strong acids with solutions of pH1 e.g. hydrogen chloride forms hydrochloric acid in water HCl(aq) or H+Cl–(aq) because they are fully ionised in aqueous solution even though the original hydrogen halides were covalent! An acid is a substance that forms H+ ions in water. Bromine forms hydrogen bromide gas HBr(g), which dissolved in water forms hydrobromic acid HBr(aq).

Iodine forms hydrogen iodide gas HI(g), which dissolved in water forms hydriodic acid HI(aq). Note the group formula pattern. Reaction with Group 1 Alkali Metals Li Na K etc. Alkali metals burn very exothermically and vigorously when heated in chlorine to form colourless crystalline ionic salts e.g. NaCl or Na+Cl–. This is a very expensive way to make salt! Its much cheaper to produce it by evaporating sea water! e.g. sodium + chlorine ==> sodium chloride 2Na(s) + Cl2(g) ==> 2NaCl(s) The sodium chloride is soluble in water to give a neutral solution pH 7, universal indicator is green. The salt is a typical ionic compound i.e. a brittle solid with a high melting point. Similarly potassium and bromine form potassium bromide KBr, or lithium and iodine form lithium iodide LiI. Again note the group formula pattern. Reaction with other metals If aluminium or iron is heated strongly in a stream of chlorine (or plunge the hot metal into a gas jar of chlorine carefully in a fume cupboard) the solid chloride is formed.

aluminium + chlorine ==> aluminium chloride(white solid) 2Al(s) + 3Cl2(g) ==> 2AlCl3(s) iron + chlorine ==> iron(III) chloride(brown solid) 2Fe(s) + 3Cl2(g) ==> 2FeCl3(s) If the iron is repeated with bromine the reaction is less vigorous, with iodine there is little reaction, these also illustrate the halogen reactivity series. Halogens as Oxidizing Agents - Chemistry LibreTexts Halogens as Oxidizing Agents This page examines the trend in oxidizing ability of the Group 7 elements (the halogens): fluorine, chlorine, bromine and iodine. It considers the ability of one halogen to oxidize the ions of another, and how this changes down the group. Basic facts Consider a situation in which one halogen (chlorine, for example) is reacted with the ions of another (iodide, perhaps) from a salt solution. In the chlorine and iodide ion case, the reaction is as follows: \[Cl_2 + 2I^- \rightarrow 2Cl^- + I_2\] The iodide ions lose electrons to form iodine molecules.

In other words, they are oxidized. The chlorine molecules gain electrons to form chloride ions— they are reduced. This is therefore a redox reaction in which chlorine acts as an oxidizing agent. Fluorine Fluorine must be excluded from this discussion because its oxidizing abilities are too strong. Fluorine oxidizes water to oxygen, as in the equation below, and so it is impossible to carry out reactions with it in aqueous solution. \[2F_2 + H_2O \rightarrow 4HF + O_2\] Chlorine, Bromine and Iodine In each case, a halogen higher in the group can oxidize the ions of one lower down. For example, chlorine can oxidize bromide ions to bromine: \[Cl_2 + 2Br^- \rightarrow 2Cl^- + Br_2\] The bromine forms an orange solution. As shown below, chlorine can also oxidize iodide ions to iodine: \[Cl_2 +2I^- \rightarrow 2Cl^- + I_2\] The iodine appears either as a red solution if little chlorine is used, or as a dark gray precipitate if the chlorine is in excess.

Bromine can only oxidize iodide ions, and is not a strong enough oxidizing agent to convert chloride ions into chlorine. A red solution of iodine is formed (see the note above) until the bromine is in excess. Then a dark gray precipitate is formed. \[Br_2 + 2I^- \rightarrow 2Br^- + I_2\] Iodine won't oxidize any of the other halide ions, except possibly the extremely radioactive and rare astatide ions. To summarize Oxidation is the loss of electrons. Each of the elements (for example, chlorine) could potentially take electrons from something else and are subsequently ionized (e.g. Cl-). This means that they are all potential oxidizing agents. Fluorine is such a powerful oxidizing agent that solution reactions are unfeasible. Chlorine has the ability to take electrons from both bromide ions and iodide ions. Bromine and iodine cannot reclaim those electrons from the chloride ions formed. This indicates that chlorine is a more powerful oxidizing agent than either bromine or iodine. Similarly, bromine is a more powerful oxidizing agent than iodine.

Bromine can remove electrons from iodide ions, producing iodine; iodine cannot reclaim those electrons from the resulting bromide ions. In short, oxidizing ability decreases down the group. Explaining the trend Whenever one of the halogens is involved in oxidizing a species in solution, the halogen ends is reduced to a halide ion associated with water molecules The following figure illustrates this process: Down the group, the ease with which these hydrated ions are formed decreases; the halogens become less effective as oxidizing agents, taking electrons from something else less readily. The reason that the hydrated ions form less readily down the group is due to several complicated factors. Unfortunately, this explanation is often over-simplified, giving a faulty and misleading explanation. The wrong explanation is dealt with here before a proper explanation is given. The Incorrect Explanation The following explanation is normally given for the trend in oxidizing ability of chlorine, bromine and iodine: The ease of ionization depends on how strongly the new electrons are attracted. As the atoms get larger, the new electrons are further from the nucleus and increasingly shielded by the inner electrons (offsetting the effect of the greater nuclear charge). The larger atoms are therefore less effective at attracting new electrons and forming ions.

This is equivalent to saying electron affinity decreases down the group. Electron affinity is described in detail on another page. The problem with this argument is that it does not include fluorine. Fluorine's tendency to form a hydrated ion is much higher than that of chlorine. However, fluorine's electron affinity is less than that of chlorine. This contradicts the above argument. This problem stems from examining a single part of a very complicated process. The argument about atoms accepting electrons applies only to isolated atoms in the gas state picking up electrons to form isolated ions, also in the gas state. The argument must be generalized. In reality: The halogen starts as a diatomic molecule, X2. This may be a gas, liquid or solid at room temperature, depending on the halogen. The diatomic molecule must split into individual atoms (atomization) Each atom gains an electron (electron affinity; this is the element of the process of interest in the faulty explanation.) The isolated ions are surrounded by water molecules; hydrated ions are formed (hydration).

The Correct Explanation The table below shows the energy involved in each of these changes for atomization energy, electron affinity , and hydration enthalpy (hydration energy): Reactions of Main Group Elements with Halogens - Chemistry LibreTexts Reactions of Main Group Elements with Halogens Last updated Answers This section describes the chemistry of halogens with the main group elements such as the alkali metals , alkaline earth metals , and Groups 13 and 14 . The word halogen itself means "salt former" in Greek. Halogens such as chlorine, bromine and iodine have properties that enable them to react with other elements to form important salts such as sodium chloride, also known as table salt. Properties of Halogens Elements such as fluorine, chlorine, bromine, iodine, and astatine belong to Group 17, the halogen group. At room temperature fluorine is a yellow gas, chlorine is a pale green gas, bromine is a red liquid, and iodine is a purple solid. Astatine is a radioactive element, and exists in nature only in small amounts. All the halogens exist as diatomic molecules. They have high ionization energies and form the most electronegative group of elements.

Their electron configuration, ns2np5, allows them to easily react with Group 1 and 2 metals; each halogen tends to pick up one electron, and the Group 1 and Group 2 elements each tend to lose one or two electrons, respectively. Halogens therefore react most vigorously with Group 1 and Group 2 metals of all main group elements. Reaction with Water From a standard reduction potential table, it is determined that iodine and bromine cannot oxidize water to oxygen because they have smaller reduction potentials than oxygen. Thus, iodine and bromine do not react with water. However, fluorine and chlorine have larger reduction potentials, and can oxidize water. Fluorine reacts with water vapor to form oxygen and ozone: \[2F_2(g) + 2H_2O(g) \rightarrow 4HF(g)+ O_2(g) \] \[3F_2(g) + 2H_2O(g) \rightarrow 6HF(g) + O_3(g)\] The reaction of water with chlorine, shown below, proceeds very slowly.

\[Cl_2 + H_2O \rightarrow H^+ + Cl^- + HClO\] Chlorine and bromine are moderately soluble in water. These solutions form solid hydrates within an ice lattice. These solutions are good oxidizing agents. Chlorine reacts reversibly with water to produce acids as in the following example, in which chloric acid and hydrochloric acid are formed: \[Cl_2 + H_2O \rightleftharpoons HClO + HCl\] Iodine is slightly soluble in water. It has the lowest standard reduction potential of the halogens, and is therefore the least powerful oxidizing agent. Air and other reagents can oxidize acidified solution of iodide ions. Reaction with Hydrogen All the halogens react directly with hydrogen, forming covalent bonds and—at sufficient levels of purity—colorless gases at room temperature. Hydrogen reacts with fluorine, chlorine, bromine, and iodine, forming HF, HCl, HBr, and HI, respectively. The bond strength of these molecules decreases down the group: \(HF > HCl > HBr > HI\).

Iodine and hydrogen react non-spontaneously to produce hydrogen iodide: \[H_2 + I_2 \rightarrow 2HI\] All the hydrogen halides are soluble in water, in which they form strong acids (with the exception of \(HF\)). The general equation of hydrogen halide for the acid reaction is given below: \[HX + H_2O \rightarrow H_3O^+ + X^-\] Group 1: The Alkali Metals All the alkali metals react vigorously with halogens to produce salts, the most industrially important of which are NaCl and KCl. \[ 2Na(s) + Cl_2(g) \rightarrow 2NaCl(s)\] Sodium Chloride is used as a preservative for meat and to melt the ice on the roads (via freezing point depression ). KCl is important for plant fertilizers because of the positive impact of potassium on plant growth. These metal halides form white ionic crystalline solids and are all soluble in water except LiF, because of its high lattice enthalpy attributed to strong electrostatic attraction between Li+ and F- ions.

Group 2: The Alkaline Earth Metals The alkaline earth metals react to form hydrated halides. These halides are ionic except for those involving beryllium (the least metallic of the group). Because alkaline earth metals tend to lose electrons and halogen atoms tend to gain electrons ( Table P2 ), the chemical reaction between these groups is the following: \[M + X_2 \rightarrow MX_2\]

Question: What name is given to the group of elements that includes fluorine, chlorine, bromine and iodine?
Answer: Halogens