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This page examines the trend in oxidizing ability of the Group 17 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:
\[\ce{Cl_2 + 2I^- \rightarrow 2Cl^- + I_2}\]
- 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.
\[\ce{2F_2 + 2H_2O \rightarrow 4HF + O_2}\]
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:
\[\ce{Cl_2 + 2Br^- \rightarrow 2Cl^- + Br_2}\]
The bromine forms an orange solution. As shown below, chlorine can also oxidize iodide ions to iodine:
\[\ce{Cl_2 +2I^- \rightarrow 2Cl^- + I_2}\]
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.
\[\ce{Br_2 + 2I^- \rightarrow 2Br^- + I_2}\]
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 endis 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):
| atomization energy (kJ mol-1) | electron affinity (kJ mol-1) | hydration enthalpy (kJ mol-1) | overall (kJ mol-1) | |
|---|---|---|---|---|
| F | +79 | -328 | -506 | -755 |
| Cl | +121 | -349 | -364 | -592 |
| Br | +112 | -324 | -335 | -547 |
| I | +107 | -295 | -293 | -481 |
Consider first the fifth column, which shows the overall heat evolved, the sum of the energies in the previous three columns.
The amount of heat evolved decreases quite dramatically from the top to the bottom of the group, with the biggest decrease between fluorine and chlorine. Fluorine generates a large amount of heat when it forms its hydrated ion, chlorine a lesser amount, and so on down the group.
The first electron affinity is defined as the energy released when 1 mole of gaseous atoms each acquire an electron to form 1 mole of gaseous 1- ions, as in the following equation: In symbol terms:
\[ X(g) + e^- \rightarrow X^-(g)\]
The fifth column measures the energy released when 1 mole of gaseous ions dissolves in water to produce hydrated ions, as in the following equation, which is not equivalent to that above:
\[ X^-(g) \rightarrow X^- (aq)\]
Why is fluorine a stronger oxidizing agent than chlorine?
There are two main factors. First, the atomization energy of fluorine is abnormally low. This reflects the low bond enthalpy of fluorine.
The main reason, however, is the very high hydration enthalpy of the fluoride ion. That is because fluoride is very small. There is a very strong attraction between fluoride ions and water molecules. The stronger the attraction, the more heat is evolved when the hydrated ions are formed.
Why does oxidizing ability decrease from chlorine to bromine to iodine?
The decrease in atomization energy between these three elements is relatively small, and would tend to make the overall change more negative down the group. It is helpful to look at the changes in electron affinity and hydration enthalpy down the group. Using the figures from the previous table:
| change in electron affinity (kJ mol-1) | change in hydration enthalpy (kJ mol-1) | |
|---|---|---|
| Cl to Br | +25 | +29 |
| Br to I | +29 | +42 |
Both of these effects contribute, but the more important factor—the one that changes the most—is the change in the hydration enthalpy. Down the group, the ions become less attractive to water molecules as they get larger. Although the ease with which an atom attracts an electron matters, it is not as important as the hydration enthalpy of the negative ion formed.
The faulty explanation is incorrect even if restricted to chlorine, bromine and iodine:
- 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 the form 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.
Contributors and Attributions
Jim Clark (Chemguide.co.uk)
(Video) Strength of halogens as oxidizing agents, s and p block elements, Lecture # 60, urdu/hindi
FAQs
How do halogens act as oxidizing agents? ›
Halogens serve as a strong oxidizing agent due to their high electronegativity and electron affinity, allowing them to take electrons from other elements quickly and easily and oxidize them. As a result, it is an excellent oxidizing agent.
Which of the following statement is not true for halogens? ›Chlorine has the highest electron-gain enthalpy due to electron repulsion domination nature which is highest in chlorine and smaller in fluorine. Hence from all the statements it is clear that the statement A is the false statement.
Are all halogens oxidizers? ›All halogens gain electrons to make halide ions, so all the halogens are oxidizing agents.
Does oxidising power of halogens decrease down the group? ›As the halogens become less reactive down the group, their oxidising ability decreases. Remember that oxidation is a gain of electrons. Halogens become less oxidising as you move down the group as it is more difficult to gain an electron.
What halogen is oxidising agent? ›Fluorine is such a powerful oxidizing agent that solution reactions are unfeasible.
Why do halogens act as reducing agents? ›When a halide ion acts as a reducing agent, it transfers electrons to something else. That means that the halide ion itself loses electrons. The larger the halide ion, the farther the outer electrons are from the nucleus, and the more they are shielded by inner electrons.
Which statement about halogens are correct? ›Halogens are all diatomic and form univalent ions. All halogens can act as both oxidizing and reducing agents and are capable of exhibiting more than one stable oxidation state except F.
What is true about all halogens? ›The halogens all have seven electrons in their outer shells. The electron configuration in the outer shell is ns2np5. As the atomic number increases, the reactivity of the halogens decreases. Fluorine and chlorine exist as gases at room temperature, while bromine is a liquid, and iodine is a solid.
What is the problem with halogens? ›Some of the problems caused by halogenated compounds are contamination of water, soil and air, and they also destroy the ozone layer.
Which halogens are the strongest oxidizing? ›Florine is the most powerful oxidizing agent because it is the most electronegative element.
Which halogen is the strongest oxidizing agent? ›
We conclude that fluorine is the most powerful oxidant among the halogens.
What does the oxidizing power of halogens depend on? ›The halogens are good oxidizers The oxidizing power of halogens in aqueous medium depends on their electron affinity, bond dissociation energy and heat of hydration. Basic character of their anions depends on their size and electronegativity.
Which is the strongest oxidizing agent? ›Fluorine is the strongest oxidizing agent because it is the strongest oxidant among all the elements.
What is the strongest reducing agent in halogens? ›In solution the halide ions act as reducing agents, the strongest ability increases down the group. HI is the strongest reducing agent.
Why do halogens have different oxidation states? ›Because ionization energies decrease down the group, the heavier halogens form compounds in positive oxidation states (+1, +3, +5, and +7).
Which halogen oxidizing agent is the weakest? ›I2 is weakest oxidising agent due to its lowest reduction potential value of +0.54 V.
What are halogens in order of oxidizing strength? ›In general, a halogen of a lower atomic number oxidizes halide ions of higher atomic number. Hence the correct order of the oxidizing power of halogen is ${I_2} < B{r_2} < C{l_2} < {F_2}$.
Which halogen can form only oxidation state? ›Fluorine exhibits only −1 oxidation state where as other halogens exhibit +1,+3,+5 and +7 oxidation states also.
Which ion is the best oxidizing agent? ›Elemental fluorine, for example, is the strongest common oxidizing agent. F2 is such a good oxidizing agent that metals, quartz, asbestos, and even water burst into flame in its presence.
Are halogen ions reducing agents? ›As a reducing agent, a halide ion loses an outermost electron and the ability to lose this electron depends on the shielding effect and the ionic radius. The greater the reducing power of a halide, the more easily it can lose electrons.
What is meant by the term oxidizing agent? ›
An oxidizing agent, or oxidant, gains electrons and is reduced in a chemical reaction. Also known as the electron acceptor, the oxidizing agent is normally in one of its higher possible oxidation states because it will gain electrons and be reduced.
What is the most common oxidation state of the halogens? ›All the elements of the halogen family exhibit -1 oxidation state.
Do all halogens show positive oxidation state? ›Fluorine, the most electronegative element, has no positive oxidation states, but the other halogens commonly exhibit +1, +3, +5, and +7 states. Most compounds containing halogens in positive oxidation states are good oxidizing agents, however, reflecting the strong tendency of these elements to gain electrons.
What is one reason the halogens are important? ›What are some uses of halogen elements? Chlorine is used to purify water. In addition, chlorine is part of table salt, sodium chloride, which is one of the most widely used chemical compounds. Fluorine is used in fluorides, which are added to water supplies to prevent tooth decay.
What are three unique things about halogens? ›They all form acids when combined with hydrogen. They are all fairly toxic. They readily combine with metals to form salts. They have seven valence electrons in their outer shell.
What are 3 facts about halogens? ›- Halogen have very high electronegativities.
- They have seven valence electrons (one short of a stable octet)
- They are highly reactive, therefore toxics.
- The halogens are Fluorine (F), Chlorine (Cl), Bromine (Br), Iodine (I) and Astatine (At)
- Down the group, atom size increases.
Summary of Common Properties
They have very high electronegativities. They have seven valence electrons (one short of a stable octet). They are highly reactive, especially with alkali metals and alkaline earths. Halogens are the most reactive nonmetals.
- Halogens are highly reactive and as such can be harmful or lethal to biological organisms in sufficient quantities. This reactivity is due to the high electronegativity of the atoms due to their high effective nuclear charge.
What makes halogens so reactive? ›Halogens are highly reactive because they readily gain an electron to fill their outermost shell. Alkali metals are highly reactive because they readily lose the single electron in their outermost shell.
Why are halogens unstable? ›Because the halogen elements have seven valence electrons, they only require one additional electron to form a full octet. This characteristic makes them more reactive than other non-metal groups.
Why do the halogens show an oxidation state? ›
Because ionization energies decrease down the group, the heavier halogens form compounds in positive oxidation states (+1, +3, +5, and +7).
How do halogens act? ›Halogens can act as electrophiles to attack a double bond in alkene. Double bond represents a region of electron density and therefore functions as a nucleophile.
How do halogens react with oxygen? ›Halogens reacts with oxygen to form acidic oxides.
Why do halogens get darker down the group? ›As the ionization energy decreases, the electron tends to easily excite to the higher energy level and the increase in atomic radii leads to absorb more visible light. Thus, the Halogens get darker down the group.
What does oxidizing power of halogens depend on? ›The halogens are good oxidizers The oxidizing power of halogens in aqueous medium depends on their electron affinity, bond dissociation energy and heat of hydration. Basic character of their anions depends on their size and electronegativity.
What happens when halogens react? ›Halogens react readily with all sorts of metals, including groups 1, 2, 3 and transition metals. They also react with hydrogen. When reacting with metals, halogens form salts with a giant ionic structure, and when reacting with hydrogen, they form hydrogen halides.
What is the function of halogens? ›Halogens are used in the chemical, water and sanitation, plastics, pharmaceutical, pulp and paper, textile, military and oil industries. Bromine, chlorine, fluorine and iodine are chemical intermediates, bleaching agents and disinfectants.
What do halogens react with the most? ›Halogens therefore react most vigorously with Group 1 and Group 2 metals of all main group elements.
What happens to halogens in water? ›Halogens react to a small extent with water, forming acidic solutions with bleaching properties. They also undergo redox reactions with metal halides in solution, displacing less reactive halogens from their compounds.
What happens when halogens react with water? ›Reactions with hydrogen
They dissolve in water to produce acidic solutions . Hydrogen chloride dissolves in water to produce hydrochloric acid, HCl(aq).
Do halogens increase or decrease? ›
The reactivities of the halogens decrease down the group ( At < I < Br < Cl < F). This is due to the fact that the atomic radius increases in size with an increase of electronic energy levels. This lessens the attraction for valence electrons of other atoms, decreasing reactivity.
Which halogen is the strongest reducing agent? ›In solution the halide ions act as reducing agents, the strongest ability increases down the group. HI is the strongest reducing agent.