Group 7 (halogens)
I can describe the pattern in reactivity of Group 7 non-metals and name the products of reactions between metals and halogens.
Lesson details
Key learning points.
- The elements in Group 7 are all non-metals that have coloured vapours, most of which are toxic.
- The state of the elements at room temperature changes from gas to solid, down Group 7.
- Reactivity decreases down Group 7.
- Halogens react with Group 1 and 2 metals to form metal halides.
- Metal halides are soluble in water.
Common misconception
All halogens are gases at room temperature because they're non-metals; all halogens form toxic compounds.
Not all halogens are gases; bromine is a liquid and iodine is a solid at room temperature. While halogens are toxic, they also form essential compounds like table salt and disinfectants.
Halogen - Elements found in Group 7 of the periodic table.
Group - Group is a column on the periodic table that corresponds to the number of electrons in the outer shell of an atom.
Halide - Refers to a compound of a halogen with another element or group.
Salt - Is a substance that forms when a metal reacts with a non-metal. These reactions typically produce solid substances that can dissolve in water, like table salt.
Spotting tiles, aqueous samples of each halogen, sodium halide salts, dropping pipettes or bottles.
Content guidance
- Risk assessment required - equipment
Supervision
Adult supervision required
This content is © Oak National Academy Limited ( 2024 ), licensed on Open Government Licence version 3.0 except where otherwise stated. See Oak's terms & conditions (Collection 2).
Starter quiz
6 questions.
fluorine -
A pale yellow gas at room temperature.
chlorine -
A pale green gas at room temperature.
bromine -
A reddish–brown liquid at room temperature.
A purple–black solid at room temperature.
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Halogen reactions with iron wool
In association with Nuffield Foundation
- No comments
Use this demonstration to illustrate an exothermic redox reaction by heating iron wool in the presence of chlorine gas and the vapours of bromine and iodine
In this teacher demonstration, students observe what happens when iron wool is heated and chlorine gas and the vapours of bromine and iodine are passed over it. Students can watch the iron wool begin to glow as the expected exothermic redox reactions take place, forming iron(III) halides (FeX 3 ) as coloured solids.
The vigour of the reactions corresponds to the order chlorine > bromine > iodine, showing the trend of decreasing reactivity of the elements down Group 17.
Source: Royal Society of Chemistry
Watch as iron wool is heated in the presence of chlorine, bromine and iodine to produce a series of exothermic redox reactions.
These experiments must be carried out in a fume cupboard as both the reactants and products are hazardous. Teachers attempting this demonstration for the first time are strongly advised to carry out a trial run before doing it in front of a class.
The time allowed should be at least 20 minutes, depending on the amount of discussion and testing of the products between each experiment.
In addition to using this demonstration to show the relative reactivity of the halogens, the reaction of chlorine or bromine with iron could be used on its own to show the reaction between a reactive non-metallic element and a metal.
- Eye protection for teacher and students
- Protective gloves for teacher
- Access to a fume cupboard
- Apparatus to set up a chlorine generator
- Boiling tubes, x2
- Reduction tube (see note 3 below)
- Beakers, 100 cm 3 , x3
- Teat pipette (preferably glass, with a narrow tip)
- Test tubes, small, x3
- Test tube rack
- Bunsen burner
- Heat resistant mat
- Bosses, clamps and stands
- Iron wool, 3 tufts about 1 g mass each (see note 4 below)
- Hexane (HIGHLY FLAMMABLE, HARMFUL), 100 cm 3 (see note 5 below)
- Liquid bromine (VERY TOXIC, CORROSIVE), 0.5 cm 3
- Sodium thiosulfate solution, 1 M, 500 cm 3
- Iodine solid (HARMFUL), 0.5 g
- Silver nitrate solution, approximately 0.1 M, 10 cm 3
- Sodium chlorate(I) solution (sodium hypochlorite), 10–14% (w/v) (CORROSIVE), fresh
- Hydrochloric acid, 5 M (IRRITANT AT THIS DILUTION)
- Deionised or distilled water
Health, safety and technical notes
- Read our standard health and safety guidance.
- Wear suitable eye protection (goggles) throughout and work in a fume cupboard at each stage of this demonstration.
- The reduction tube should be fitted with a one-holed rubber stopper fitted with short length of glass tubing. Alternatively, an 8–10 cm length of wide-bore glass tubing with a stopper at each end fitted with a short length of glass tubing could be used. See diagram below.
- Iron wool, Fe(s) – see CLEAPSS Hazcard HC055A . The finest grade iron wool is best since it provides the maximum surface area. Iron wool is often sold with a thin layer of grease on its surface to stop it rusting. Working in a fume cupboard, the layer of grease can be removed by dipping the iron wool in hexane (or alternative) a few times. The solvent must be allowed to completely evaporate.
- Hexane, C 6 H 14 (l), (HIGHLY FLAMMABLE, HARMFUL, DANGEROUS FOR THE ENVIRONMENT) – see CLEAPSS Hazcard HC045a .
- Liquid bromine, Br 2 (l), (VERY TOXIC, CORROSIVE, DANGEROUS FOR THE ENVIRONMENT) – see CLEAPSS Hazcard HC015a . See our standard guidance on handling liquid bromine and preparing bromine water . Wear suitable protective gloves when handling bromine and have at least 500 cm 3 of 1M sodium thiosulfate readily available to treat any spillages.
- Sodium thiosulfate solution, Na 2 S 2 O 3 (aq) – see CLEAPSS Hazcard HC095A .
- Iodine solid, I 2 (s), (HARMFUL, DANGEROUS FOR THE ENVIRONMENT) – see CLEAPSS Hazcard HC054 .
- Silver nitrate solution, AgNO 3 (aq) – see CLEAPSS Hazcard HC087 . Although LOW HAZARD at this concentration, the solution can still stain skin, clothes and some bench materials.
- For generating chlorine gas, Cl 2 (g), (TOXIC, DANGEROUS FOR THE ENVIRONMENT), see CLEAPSS Hazcard HC022a and CLEAPSS Recipe Book RB024. See our guidance on standard techniques for generating, collecting, and testing gases .There are two methods given in the standard techniques for generating chlorine . The method that uses sodium chlorate(I) (sodium hypochlorite) is safer than the method that uses potassium manganate(VII), but will not work well if the sodium chlorate(I) (sodium hypochlorite) is an old sample. Note that sodium chlorate(I), NaOCl, is NOT the same as chlorate(V), NaClO 3 .
- Sodium chlorate(I) solution, 10–14% (w/v) NaOCl, fresh (CORROSIVE) – see CLEAPSS Hazcard HC089 .
- Hydrochloric acid 5 M, HCl(aq) – see CLEAPSS Hazcard HC047a .
Chlorine
- Place a 1 g tuft of cleaned iron wool in the reduction tube so that it is well spread out. Leave at least a 1 cm gap between the stopper and the iron wool.
- Connect the reduction tube to the chlorine generator with a short length of rubber tubing. Clamp it in position over a Bunsen burner.
The equipment required for heating iron wool as chlorine gas is passed over it
- Pass a slow stream of chlorine over the iron wool from the chlorine gas generator. Do this by allowing the hydrochloric acid to drip slowly on to the potassium manganate(VII). After a few seconds, it should be possible to see the greenish colour of the chlorine gas filling the reduction tube, as all the air is expelled.
- The iron wool may ignite without any heating. If not, gently heat at the end nearest to the chlorine generator until the wool does ignite (no further heating should be required).
- A vigorous reaction will occur and the glow will spread along the wool in the tube, producing clouds of brown iron(III) chloride. Some of this may emerge as a smoke from the end of the reduction tube.
- Continue passing chlorine over the iron wool until no further reaction occurs. Stop the chlorine supply and allow the tube to cool.
- When cool, disconnect the reduction tube and rinse a little of the product into a clean beaker with some distilled water. Pour some of this solution into a clean test tube and test with a few drops of silver nitrate solution. A white precipitate of silver chloride will form, confirming the presence of chloride ions.
- Wear suitable protective gloves and take care to avoid spillage when handling liquid bromine. It is CORROSIVE and VERY TOXIC. Transfer about 0.5 cm 3 liquid bromine into one of the boiling tubes, using the teat pipette. Care is needed to avoid spillage – the density and volatility of the bromine cause it to drip very easily from the pipette. Keep the bromine container and the mouth of the test tube close together. Replace the lid of the bromine container immediately.
- Using tongs or tweezers, place a 1 g tuft of cleaned iron wool into the boiling-tube so that it is well spread out and almost fills the boiling-tube. Leave a 2 cm gap between the iron wool and the surface of the liquid bromine.
- Clamp the test tube near the top and at an angle of about 45° – see diagram.
The equipment required for heating iron wool in the presence of vapours of bromine or iodine
- Heat the test tube, gently at first, with a yellow-tipped blue flame (air hole on Bunsen burner slightly closed). Do this by moving the flame slowly between the bottom half of the iron wool and the bromine. As the bromine vapour starts to rise up into the iron wool, heat the wool more strongly.
- Remove the heat when the wool starts to glow due to the heat of the reaction. Note the extra heating required to get this reaction started compared to the reaction involving chlorine. The iron will become coated with yellow-brown iron(III) bromide, and a brown ‘smoke’ may escape from the mouth of the test tube.
- When the reaction appears to be over, use tongs or tweezers to remove some of the remaining iron wool from the test tube.
- Rinse the iron wool in a few cubic centimetres of deionised/distilled water in a small beaker. Pour out some of the resulting solution into a clean test tube and test with a few drops of silver nitrate solution. Formation of a cream precipitate of silver bromide confirms that bromide ions are present.
- Transfer about 0.5 g of solid iodine (HARMFUL) into one of the boiling tubes. Place a 1 g tuft of cleaned iron wool in the test tube and clamp it as before.
- Working in a fume cupboard, heat the test tube with a yellow-tipped blue flame (air hole on Bunsen burner slightly closed). Heat gently at first by moving the flame slowly between the bottom half of the iron wool and the iodine.
- As the purple iodine vapour starts to rise up into the iron wool, heat the wool more strongly. Remove the heat when the reaction causes a dull glow – see Teaching notes below. Some red-brown iron(III) iodide should form.
- When the reaction appears to be over, remove some of the remaining iron wool from the test tube with tweezers and rinse it in a few cubic centimetres of deionised/distilled water in a small beaker.
- Pour some of the resulting solution into a clean test tube and test it with a few drops of silver nitrate solution. Formation of a yellow precipitate of silver iodide confirms that iodide ions are present.
Teaching notes
The order in which the experiments are done is a matter of choice, but it is probably best to leave the most reactive halogen (chlorine) to last, to end with a vigorous reaction – and confirm a class prediction?
The reaction with iodine is much less vigorous than that with bromine and it may be difficult to see a glow at all. A couple of trial experiments beforehand may be necessary to get the right balance between heating the iodine and getting the iron hot enough for a reaction to start. If the iron is heated too vigorously, it may start to glow from reaction with the oxygen in any air that may still be present in the test tube.
The general equation for the reactions involved is:
2Fe(s) + 3X 2 (g) → 2FeX 3 (s) (X = Cl, Br and I)
Additional information
This is a resource from the Practical Chemistry project , developed by the Nuffield Foundation and the Royal Society of Chemistry. This collection of over 200 practical activities demonstrates a wide range of chemical concepts and processes. Each activity contains comprehensive information for teachers and technicians, including full technical notes and step-by-step procedures. Practical Chemistry activities accompany Practical Physics and Practical Biology .
The experiment is also part of the Royal Society of Chemistry’s Continuing Professional Development course: Chemistry for non-specialists .
© Nuffield Foundation and the Royal Society of Chemistry
- 11-14 years
- 14-16 years
- 16-18 years
- Demonstrations
- Elements and the periodic table
- Redox chemistry
Specification
- (q) the reactions of halogens with alkali metals and with iron
- (p) the reactions of halogens with alkali metals and with iron
- Students should be able to describe the nature of the compounds formed when chlorine, bromine and iodine react with metals and non-metals.
- In Group 7, the reactivity of the elements decreases going down the group.
- Explain how the observed simple properties of Group 7 depend on the outer shell of electrons of the atoms and predict properties from given trends down the group.
- 6.9 Describe the reactions of the halogens, chlorine, bromine and iodine, with metals to form metal halides, and use this pattern to predict the reactions of other halogens
- C2.2.8 describe experiments to identify the reactivity pattern of Group 7 elements including displacement reactions
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| 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 before you go on. | ||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||
| In the chlorine and iodide ion case, the reaction would be: + 2I + IThe 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.
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. + 2H O
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: + 2Br + BrThe 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: + 2I + IThe 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. | ||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||
| The reason for the red solution is that iodine dissolves in potassium iodide (or other soluble iodides) by reacting to give a red ion, I . 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. | ||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||
| A red solution of iodine is formed (see the note above) until the bromine is in excess. Then you get a dark grey precipitate. + 2I + IIodine 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). ). That means that they are all potentially oxidising agents. That means that chlorine is a more powerful oxidising agent than either bromine or iodine. This all means that oxidising ability falls as you go down the Group.
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.
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. | ||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||
| If you haven't recently read about the , you ought to follow this link before you go on. Use the BACK button on your browser to return to this page. | ||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||
| 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: - which may be gas, liquid or solid, depending on the halogen. | ||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||
| For the next bit, if you aren't happy about enthalpy changes, you might want to explore the of Chemguide, or . | ||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||||
| The table below looks at how much energy is involved in each of these changes. To be sure that you understand the various terms:
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 X ). 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.
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: (g)
This is the energy released when 1 mole of gaseous ions dissolves in water to produce hydrated ions. (g) + (aq) (aq)
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:
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.
Where would you like to go now? To the Group 7 menu . . . To the Inorganic Chemistry menu . . . To Main Menu . . . © Jim Clark 2002 (modified February 2022)  Personalised lessons and regular feedback to ensure you ace your exams! Book a free consultation today 100+ Video Tutorials, Flashcards and Weekly Seminars Gain hands-on experience of how chemistry is used in different fields. Experience life as a uni student and boost your university application with our summer programme!
Group 7: Reactions & Displacement (GCSE Chemistry)Group 7: reactions & displacement. Table of Contents Halogens – Reactions with Non-MetalsHalogens also form covalent bonds with other non-metals to form a molecule. Lets work through an example, below is a generic formula that represents the halogens (X) reacting with hydrogen: H 2 (g) + X 2 (g) → 2 HX (g) When this reaction occurs a hydrogen halide is formed which is gaseous at room temperature. When dissolved in water this forms an acidic solution. Halogens – Reactions with MetalsThe halogens react with metals to form salts . When these salts are dissolved in water they form colourless solutions. Below is a generic formula that represents the halogens (X) reacting with sodium: 2 Na (s) + X 2 (g) → 2 NaX (s) When a halogen reacts with a metal, the metal atom loses an electron forming a positively charged ion. That lost electron is then taken up by the halogen atom, this forms a negatively charged ion. The oppositely charged ions are attracted to each other; this is known as ionic bonding .  The electron (blue) is then taken up by chlorine to form a negatively charged chloride ion. 2 Na (s) + Cl 2 (g) → 2 NaCl (s) Halogens – Displacement ReactionA displacement reaction is when more reactive halogen can displace a less reactive halogen from a solution of its salt. All these reactions occur in solution and all the salts when dissolved in water are colourless . Let’s work through some examples. 1. Chlorine solution reacting with potassium bromide2 KBr (aq) + Cl 2 (aq) → 2 KCl (aq) + Br 2 (aq) Chlorine is the more reactive halogen and it will displace bromine from potassium bromide. The aqueous bromine will turn the solution from colourless to yellow-orange. 2. Bromine solution reacting with potassium chloride2 KCl (aq) + Br 2 (aq) → NO REACTION Chlorine is the more reactive halogen and will not be displaced by bromine. Therefore no reaction will occur and the solution will remain colourless. Below is a summary of all the displacement reactions you need to know, make sure you learn how to write a balanced equation for all these reactions and observations.  Group 7 in GCSE chemistry refers to the halogens, a group of elements that include fluorine, chlorine, bromine, iodine, and astatine. Group 7 elements are important in GCSE chemistry because they demonstrate common properties and reactivity patterns that can be used to understand and predict chemical reactions. A displacement reaction is a type of chemical reaction in which one element is replaced by another element in a compound. This occurs when a more reactive element replaces a less reactive element. The reactivity series in GCSE chemistry is a list of elements ordered from most reactive to least reactive. It is used to predict the outcome of displacement reactions. Some common examples of displacement reactions in group 7 elements include the reaction between sodium chloride and silver nitrate to form silver chloride and sodium nitrate, and the reaction between copper and silver nitrate to form copper nitrate and silver. In a reaction between a metal and a halogen, the metal displaces the halogen to form a metal halide, which is a salt. The products of a displacement reaction can be determined by using the reactivity series and predicting which element will displace the other. Displacement reactions occur because elements in a compound have different electronegativities, and the more electronegative element will displace the less electronegative element. No, a displacement reaction can only occur between a metal and a non-metal. Still got a question? Leave a commentLeave a comment, cancel reply. Save my name, email, and website in this browser for the next time I comment. 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Boost your GCSE Chemistry PerformanceGet a 9 in GCSE Chemistry with our Trusted 1-1 Tutors. Enquire now. 100+ Video Tutorials, Flashcards and Weekly Seminars. 100% Money Back Guarantee Learn live with other students and gain expert tips and advice to boost your score.  Let's get acquainted ? What is your name?Nice to meet you, {{name}} what is your preferred e-mail address, nice to meet you, {{name}} what is your preferred phone number, what is your preferred phone number, just to check, what are you interested in, when should we call you. It would be great to have a 15m chat to discuss a personalised plan and answer any questions What time works best for you? (UK Time)Pick a time-slot that works best for you ? How many hours of 1-1 tutoring are you looking for?My whatsapp number is..., for our safeguarding policy, please confirm.... Please provide the mobile number of a guardian/parent Which online course are you interested in?What is your query, you can apply for a bursary by clicking this link, sure, what is your query, thank you for your response. we will aim to get back to you within 12-24 hours., lock in a 2 hour 1-1 tutoring lesson now. If you're ready and keen to get started click the button below to book your first 2 hour 1-1 tutoring lesson with us. Connect with a tutor from a university of your choice in minutes. (Use FAST5 to get 5% Off!) Group 7: The Halogens ( AQA GCSE Chemistry )Revision note.  Chemistry Lead Atomic structure of Group 7 elements
 The atoms of the elements of Group 7 all have 7 electrons in their outer shell You will only be expected to draw the electron configurations for fluorine and chlorine. This is because bromine, iodine and astatine are beyond the GCSE model and specification. Properties of Group 7 Elements
The properties of the Group 7 elements
Boiling points of the Group 7 elements This graph shows the melting and boiling points of the Group 7 elements Group 7 elements state at room temperature
The appearance and state of the Group 7 elements The physical states and colours of chlorine, bromine and iodine at room temperature Exam questions on this topic occur often so make sure you know and can explain the trends of the group 7 elements in detail, using their electron configurations. Reactivity of the Halogens
Displacement Reactions
Chlorine with Bromine & Iodine
chlorine + potassium bromide → potassium chloride + bromine Cl 2 + 2KBr → 2KCl + Br 2 chlorine + potassium iodide → potassium chloride + iodine Cl 2 + 2KI → 2KCl + I 2 Bromine with iodine
bromine + potassium iodide → potassium bromide + iodine Br 2 + 2KI → 2KBr + I 2 Summary table of the displacement reactions of the halogens: chlorine, bromine and iodine
Cl 2 + 2Br - → 2Cl - + Br 2 Cl 2 + 2I - → 2Cl - + I 2 Br 2 + 2I - → 2Br - + I 2 Reactions of the HalogensMetal halides.
Formation of sodium chloride Sodium donates its single outer electron to a chlorine atom and an ionic bond is formed between the positive sodium ion and the negative chloride ion Non-metal halides
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1.6.19 investigate the displacement reactions of Group 7 (VII) elements with solutions of other halides to establish the trend in reactivity within the group and make predictions based on this trend; ... Mandatory experiment 1.2 - Redox reactions of group VII elements - halogens as oxidising agents (reactions with bromides, iodides, Fe²⁺ and ...
Topic 4B: The elements of Group 7 (halogens) 11. understand the trend in reactivity of Group 7 elements in terms of the redox reactions of Cl₂, Br₂ and I₂ with halide ions in aqueous solution, followed by the addition of an organic solvent; OCR Chemistry A. Module 3: Periodic table and energy. 3.1 The periodic table. 3.1.3 The halogens
The reactions between hydrogen and halogens become less vigorous as you go down group 7: fluorine reacts explosively in the cold and dark. chlorine reacts when heated or in sunlight. bromine reacts in burning hydrogen. iodine reacts slowly during continuous heating.
produces the most reactions is the most reactive of the three. 1. Write the names of the halogens in order of reactivity, from the most reactive to the least. 2. Compare your reactivity series with group 7 in the periodic table - what do you notice? 3. What results might you get if fluorine water and potassium fluoride were in the experiment ...
A halogen displacement reaction occurs when a more reactive halogen displaces a less reactive halogen from an aqueous solution of its halide. The reactivity of Group 7 non-metals increases as you move up the group. Out of the three commonly used halogens, chlorine, bromine and iodine, chlorine is the most reactive and iodine is the least reactive.
Bookmark. The halogens, group 7, are an important group of elements both chemically and socially. They are unusual in that they are coloured and diatomic. They also have an order of reactivity that is the opposite of the metals in groups 1 and 2. In fact they are all very reactive, fluorine being the most reactive non-metal.
Key learning points. The elements in Group 7 are all non-metals that have coloured vapours, most of which are toxic. The state of the elements at room temperature changes from gas to solid, down Group 7. Reactivity decreases down Group 7. Halogens react with Group 1 and 2 metals to form metal halides. Metal halides are soluble in water.
hydrogenF2(g) + H2 (g) 2 HF (g) All the halogens react with hydrogen gas to produce hydrogen. halides. Br2(l) + H2(g) I2(s) + H2 (g) ⇌ 2 HBr (g) The reactions decrease in speed and vigour going down t. 2 HI (g)Thermal stability of halides: Hydrogen Iodide will decompose if a hot nichrome wire is plunged.
In Group 7, the reactivity of the elements decreases going down the group. AQA Combined science: Synergy. 4.5 Building blocks for understanding. 4.5.1 The periodic table. 4.5.1.5 Group 7. Explain how the observed simple properties of Group 7 depend on the outer shell of electrons of the atoms and predict properties from given trends down the group.
2F 2 + 2H 2 O 4HF + O 2. 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: Cl 2 + 2Br - 2Cl - + Br 2. The bromine appears as an orange solution.
The atoms of the elements of Group 7 all have 7 electrons in their outer shell. Reactivity of Group 7 non-metals decreases as you go down the group. As you go down Group 7, the number of shells of electrons increases, the same as with all other groups. However, halogen atoms form negative ions when they gain an electron to obtain a full outer shell. This means that the increased distance from ...
The reactivity of halogens is also shown by their displacement reactions with other halide ions in solutions; A more reactive halogen can displace a less reactive halogen from a halide solution of the less reactive halogen; Eg. The addition of chlorine water to a solution of bromine water: Cl 2 (aq) + 2NaBr (aq) → 2NaCl (aq) + Br 2 (aq). The chlorine has displaced the bromine from solution ...
When a halogen reacts with a metal, the metal atom loses an electron forming a positively charged ion. That lost electron is then taken up by the halogen atom, this forms a negatively charged ion. The oppositely charged ions are attracted to each other; this is known as ionic bonding. GCSE Chemistry - Group 7: Reactions & Displacement.
Electronic configuration of Group 7 elements. The atoms of the elements of Group 7 all have 7 electrons in their outer shell. Reactivity of Group 7 non-metals decreases as you go down the group. As you go down Group 7, the number of shells of electrons increases, the same as with all other groups. However, halogen atoms form negative ions when ...
The elements in Group 7 are known as the halogens. These are fluorine, chlorine, bromine, iodine and astatine. These elements are non-metals that are poisonous. All halogens have similar reactions as they each have seven electrons in their outermost shell. Halogens are diatomic, meaning they form molecules made of pairs of atoms sharing ...