The Group 17 elements in the periodic table, known as halogens, include fluorine (F), chlorine (Cl), bromine (Br), iodine (I), and astatine (At). Positioned in the second-to-last column of the periodic table, just before the noble gases, these elements exhibit interesting properties and possess great significance in chemistry.
Chemistry of Halogens
The term “halogen” originates from the Greek words “halos,” meaning “salt,” and “genes,” meaning “producer,” which aptly describes their ability to form salts. The halogens’ reactivity is a defining characteristic, with fluorine being the most reactive, followed by chlorine, bromine, iodine, and astatine.
The chemistry of Group 17 elements plays a vital role in various fields. Fluorine, as the most reactive halogen, is used in the production of numerous essential compounds, such as fluoropolymers and refrigerants. Its compounds are also utilized in dental care for strengthening tooth enamel.
Chlorine is extensively used in water treatment to eliminate harmful microorganisms and ensure safe drinking water. It is also applied plastics, solvents, and pharmaceuticals productions
Bromine which a liquid state at room temperature, is widely used as a flame retardant in textiles, electronics, and furniture. It is also employed in the production of dyes and pharmaceuticals.
Iodine, an essential element for thyroid function, is medically applied as an antiseptic for disinfecting wounds. It is also employed in the production of contrast agents for medical imaging.
Astatine, the rarest naturally occurring element, has limited applications due to its scarcity and high radioactivity. However, it has potential uses in cancer treatment.
Let’s explore the properties and behavior of these fascinating elements in detail.
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Occurrence and Extraction of Group 17 Elements (Halogens)
Natural Sources of Halogens
Fluorine (F)
Fluorine is not found freely in nature due to its high reactivity. However, it is present in various minerals such as fluorite (CaF₂), cryolite (Na₃AlF₆), and fluorspar (CaF₂). Fluorides are also found in the Earth’s crust, soil, and water sources.
Chlorine (Cl)
Chlorine is abundant in nature and is primarily found in the form of chloride salts. Common sources of chlorine include rock salt (NaCl), halite (NaCl), and seawater, where it exists predominantly as sodium chloride.
Bromine (Br)
Bromine occurs naturally in small quantities in seawater, brine wells, and underground brine pools. It is obtained as bromide ions (Br⁻) and can be extracted from brine sources by chemical or electrolytic methods.
Iodine (I)
Iodine is found in small amounts in seawater, seaweeds, and some mineral deposits. Seaweeds, such as kelp, are particularly rich sources of iodine. Iodine can be extracted from these sources through a process involving oxidation, filtration, and precipitation.
Extraction Methods and Industrial Production of Group 17 Elements
Fluorine (F)
The extraction of elemental fluorine is a challenging process due to its high reactivity. It is primarily produced industrially by the electrolysis of anhydrous hydrogen fluoride (HF) using a specialized apparatus called a fluorine cell.
Chlorine (Cl)
Chlorine is primarily produced through the electrolysis of brine (a concentrated solution of sodium chloride, NaCl) in a process known as the chloralkali process. The electrolysis of brine generates chlorine gas at the anode and hydrogen gas at the cathode.
2Cl⁻(aq) → Cl₂(g) + 2e⁻ (at the anode)
Bromine (Br)
Bromine is obtained as a byproduct during the extraction of salt from brine sources. The brine is first treated with chlorine gas, resulting in the oxidation of bromide ions to bromine. The liberated bromine is then separated and purified.
2Br⁻(aq) + Cl₂(g) → Br₂(aq) + 2Cl⁻(aq)
Iodine (I)
Iodine is mainly extracted from natural sources such as seaweed or brine wells. The process involves the extraction of iodide ions (I⁻) from the source material, followed by oxidation to yield iodine. The iodine is then isolated and purified.
Oxidation reaction:
2I⁻(aq) + Cl₂(aq) → I₂(aq) + 2Cl⁻(aq)
Purification:
I₂(aq) + 2NaOH(aq) → NaI(aq) + NaIO₃(aq) + H₂O(l)
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General Characteristics of Group 17 Elements (Halogens)
Electronic Configurations of Halogens
1. Fluorine (F): 1s² 2s² 2p⁵
2. Chlorine (Cl): 1s² 2s² 2p⁶ 3s² 3p⁵
3. Bromine (Br): 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁵
4. Iodine (I): 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁶ 5s² 4d¹⁰ 5p⁵
Atomic Properties of Halogens
Atomic Radius
The atomic radius of halogens increases down the group. This is due to the addition of new energy levels (n) as the atomic number increases. The increase in the number of energy levels leads to a larger atomic size because the electrons are farther from the nucleus and experience weaker attractive forces.
Atomic Mass
The atomic mass of halogens generally increases down the group. This is because as you move down the group, the number of protons and neutrons in the nucleus increases, resulting in a higher total mass for the atom.
Electron Configuration
Halogens have seven valence electrons, represented by ns²np⁵ configuration. The “n” represents the principal energy level, and the “s” and “p” orbitals indicate the types of sublevels where the electrons are found. The ns² portion of the configuration indicates the filling of the s sublevel, while the np⁵ portion indicates the filling of the p sublevel with five electrons.
The electron configuration of halogens follows the octet rule, where atoms tend to gain or share electrons to achieve a stable electron configuration with eight valence electrons (except for hydrogen and helium). Halogens have a tendency to gain one electron to achieve a stable configuration, resulting in a negative charge and forming anions.
Physical Properties
State and Color
Fluorine (F) and chlorine (Cl) are gases at room temperature, bromine (Br) is a liquid, and iodine (I) is a solid. Fluorine is a pale-yellow gas, chlorine is a greenish-yellow gas, bromine is a reddish-brown liquid, and iodine is a dark purple solid.
Melting and Boiling Points
The melting and boiling points of halogens generally increase down the group. Fluorine has the lowest melting and boiling points, while iodine has the highest melting and boiling points among the halogens.
Density and Solubility
The density of halogens also increases down the group. Chlorine is denser than air, while bromine and iodine are denser than water. Halogens are generally soluble in organic solvents but have limited solubility in water.
Variation in Atomic Radii and Electronegativity of Group 17 Element
Fluorine (F)
Fluorine has the smallest atomic radius and the highest electronegativity among the halogens due to its strong nuclear attraction.
Chlorine (Cl)
Chlorine has a larger atomic radius compared to fluorine but is still relatively small. It also exhibits a high electronegativity.
Bromine (Br)
Bromine has a significantly larger atomic radius compared to fluorine and chlorine. Its atomic radius continues to increase down the group. Bromine also has a lower electronegativity compared to fluorine and chlorine.
Iodine (I)
Iodine has the largest atomic radius among the halogens, much larger than that of fluorine, chlorine, and bromine. Iodine also has a lower electronegativity compared to the previous halogens.
Ionization Energy and Electron Affinity Trends of Group 17 Elements
The ionization energy, which refers to the energy required to remove an electron from an atom, generally decreases down Group 17. This is due to the increase in atomic size and the shielding effect of additional energy levels.
The electron affinity, which refers to the energy change when an atom gains an electron, generally increases down Group 17. This is because the effective nuclear charge increases, resulting in stronger attraction for additional electrons.
Oxidation States of Group 17 Elements (Halogens)
Common oxidation states of halogens
Fluorine (F)
Fluorine exhibits an oxidation state of -1 in almost all compounds. It is a powerful oxidizing agent and rarely shows positive oxidation states.
Chlorine (Cl)
Chlorine commonly exhibits an oxidation state of -1, similar to fluorine. However, it can also exhibit positive oxidation states such as +1, +3, +5, and +7 in certain compounds.
Bromine (Br)
Bromine commonly exhibits oxidation states of -1, +1, +3, +5, and +7. The most stable oxidation state for bromine is -1, similar to the other halogens.
Iodine (I)
Iodine exhibits oxidation states of -1, +1, +3, +5, and +7. However, the most common oxidation state for iodine is -1.
Some Exceptions for Group 17 elements
Fluorine (F)
Fluorine exhibits a stable oxidation state of -1 in compounds due to its high electronegativity and small atomic size. Its tendency to attract electrons and complete its octet leads to a strong reducing nature and high reactivity.
Chlorine (Cl)
Chlorine, similar to fluorine, primarily exhibits an oxidation state of -1. However, it can show positive oxidation states in certain compounds. Chlorine’s ability to expand its valence shell and accommodate more electrons in higher energy levels allows it to exhibit positive oxidation states, such as +1 in compounds like hypochlorous acid (HClO) and +7 in perchloric acid (HClO₄).
Bromine (Br)
Bromine exhibits similar oxidation states to chlorine but with greater ease due to its larger atomic size. It commonly exhibits an oxidation state of -1, like other halogens. However, it can also show positive oxidation states, such as +1 in compounds like hydrogen bromide (HBr) and +5 in bromic acid (HBrO₃).
Iodine (I)
Iodine exhibits oxidation states similar to the other halogens. However, it tends to exhibit a wider range of positive oxidation states, especially +3, +5, and +7. For example, iodine can form compounds such as iodine trichloride (ICl₃) and iodic acid (HIO₃) where it exhibits positive oxidation states.
| Element | Oxidation States | Examples |
|---|---|---|
| Fluorine (F) | -1 | Fluorine exhibits an oxidation state of -1 in almost all compounds e.g. Hydrogen flouride (HF). |
| Chlorine (Cl) | -1, +1, +3, +5, +7 |
|
| Bromine (Br) | -1, +1, +3, +5, +7 |
|
| Iodine (I) | -1, +1, +3, +5, +7 |
|
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Reactivity Trends within Group 17 Elements
The reactivity of halogens increases as you move down Group 17 from fluorine (F) to iodine (I). This trend is primarily attributed to the decreasing electronegativity and increasing atomic size down the group.
Fluorine (F)
Florine reacts vigorously with metals, nonmetals, and even noble gases and it is exothermic.
Reaction with Metals
Flourine reacts with metal to form ionic compounds
a. Reaction with sodium
F₂ + 2Na → 2NaF
Sodium fluoride formed is an ionic compound, with sodium losing one electron to form a sodium cation (Na+) and fluorine accepting that electron to form a fluoride anion (F–).
b. Reaction with magnesium: The same as magnesium fluoride (MgF₂) formed from the following reaction:
F₂ + Mg → MgF₂
Reaction with Non-metals
Flourine reacts with non-metal to form covalent compounds
a. Reaction with hydrogen:
F₂ + H₂ → 2HF
The combination of fluorine with hydrogen to form hydrogen fluoride (HF). Fluorine, being more electronegative than hydrogen, attracts the shared electrons in the H-H bond towards itself, resulting in the formation of positively charged hydrogen ions (H+) and negatively charged fluoride ions (F–). Hydrogen fluoride exists as a covalent compound and is a highly acidic gas.
b. Reaction with sulfur: The same also with sulfur forming a stable, non-flammable gas, Sulfur hexafluoride:
3 F₂ + S → SF₆
Reaction with Noble Gases
Noble gases rarely react except with flourine
a. Reaction with xenon:
F₂ + Xe → XeF₂
The above reaction shows the formation of xenon difluoride (XeF₂) when fluorine reacts with xenon (Xe).
b. Reaction with krypton:
F₂ + Kr → KrF₂
Likewise, the formation of krypton difluoride (KrF₂) when fluorine reacts with krypton (Kr).
Chlorine (Cl)
Reaction with Metals
– Chlorine reacts with metals to form metal chlorides through a process called oxidation. Chlorine gains electrons from the metal atoms, causing the metal to undergo oxidation, while chlorine itself undergoes reduction.
Example: 2Na + Cl₂ → 2NaCl
In this reaction, chlorine gains one electron from each sodium atom, resulting in the formation of sodium chloride (NaCl). Sodium loses one electron and is oxidized to Na+ ions, while chlorine is reduced to Cl- ions.
Reaction with Non-metals
– Chlorine readily reacts with nonmetals, often forming covalent compounds. It can act as both an oxidizing agent and a halogenating agent.
Example: H₂ + Cl₂ → 2HCl
In this reaction, chlorine combines with hydrogen to form hydrogen chloride (HCl). Chlorine gains one electron from hydrogen, resulting in the formation of chloride ions (Cl-). The reaction is highly exothermic and produces hydrogen chloride gas.
– Disproportionation reaction:
3Cl₂ + 6OH⁻ → 5Cl⁻ + ClO₃⁻ + 3H₂O
This reaction occurs in an alkaline solution where chlorine reacts with hydroxide ions (OH-). Chlorine undergoes disproportionation, meaning it simultaneously undergoes oxidation and reduction. It gains electrons from hydroxide ions, forming chloride ions (Cl-) and chlorate ions (ClO₃⁻). Water (H₂O) is also produced as a byproduct.
Reaction with Organic Compounds:
Chlorine reacts with organic compounds, particularly hydrocarbons, through a process called halogenation. Chlorine substitutes one or more hydrogen atoms in the organic compound, resulting in the formation of alkyl chlorides.
Example: CH₄ + Cl₂ → CH₃Cl + HCl
In this reaction, chlorine substitutes one hydrogen atom in methane (CH₄), resulting in the formation of chloromethane (CH₃Cl) and hydrogen chloride (HCl). The reaction is initiated by the presence of light or heat.
R-CH₃ + Cl₂ → R-C(=O)-Cl + HCl
Chlorine can also act as an oxidizing agent in reactions with organic compounds. It replaces a hydrogen atom in the organic compound (R-CH₃), leading to the formation of an acyl chloride (R-C(=O)-Cl) and hydrogen chloride (HCl).
Redox reactions involving Chlorine
– Chlorine can undergo redox reactions, where it acts as an oxidizing agent by accepting electrons from other substances. For example:
2Fe + 3Cl₂ → 2FeCl₃
In this reaction, chlorine accepts electrons from iron atoms, causing iron to undergo oxidation. Iron loses electrons and is oxidized to form iron(III) chloride (FeCl₃), while chlorine is reduced.
Another example,
H₂S + Cl₂ → 2HCl + S
Chlorine acts as an oxidizing agent in this reaction, accepting electrons from hydrogen sulfide (H₂S). This leads to the formation of hydrogen chloride (HCl) and elemental sulfur (S).
Bromine
Reaction with metals
Bromine reacts with metals to form metal bromides through a process called halogenation. Bromine replaces other halogens or nonmetals in compounds, resulting in the formation of metal bromides.
Br₂ + 2Na → 2NaBr
In this reaction, bromine replaces the chlorine in sodium chloride (NaCl), resulting in the formation of sodium bromide (NaBr). Sodium is oxidized from its elemental state to Na+ ions, while bromine is reduced.
Reaction with Non-metals
– Bromine can react with nonmetals, forming covalent compounds. The reactivity of bromine with nonmetals is generally lower compared to fluorine and chlorine.
Reaction with Hydrogen
Br₂ + H₂ → 2HBr
In this reaction, bromine combines with hydrogen to form hydrogen bromide (HBr). Bromine accepts one electron from hydrogen, leading to the formation of bromide ions (Br–). The reaction is exothermic.
Displacement reactions
Bromine can displace less reactive halogens from their compounds through a process called displacement. Bromine’s reactivity allows it to replace other halogens in compounds.
Displacement reaction of bromine with sodium iodide:
Br₂ + 2NaI → 2NaBr + I₂
In this reaction, bromine displaces iodine from sodium iodide (NaI), resulting in the formation of sodium bromide (NaBr) and elemental iodine (I₂). Bromine is more reactive than iodine, allowing it to displace iodine from the compound.
Reaction with Organic Compounds
– Bromine reacts with organic compounds, particularly hydrocarbons, through a process called halogenation. Bromine substitutes one or more hydrogen atoms in the organic compound, leading to the formation of alkyl bromides.
CH₄ + Br₂ → CH₃Br + HBr
In this reaction called halogenation of methane, bromine substitutes one hydrogen atom in methane (CH₄), resulting in the formation of bromomethane (CH₃Br) and hydrogen bromide (HBr). The reaction is initiated by the presence of light or heat.
Iodine (I)
Reaction with Metals
Iodine reacts with metals to form metal iodides through a process called iodination. Iodine replaces other halogens or nonmetals in compounds, resulting in the formation of metal iodides.
For example, the reaction with sodium:
I₂ + 2Na → 2NaI
In this reaction, iodine reacts with sodium to form sodium iodide (NaI). Sodium is oxidized from its elemental state to Na+ ions, while iodine is reduced.
Reaction with Nonmetals
– Iodine can react with nonmetals, forming covalent compounds. The reactivity of iodine with nonmetals is generally lower compared to fluorine, chlorine, and bromine. The reaction with sulfur by iodine to form sulfur diiodide (IS₂) has iodine accepts two electrons from sulfur, resulting in the formation of iodide ions (I-). The reaction is exothermic.
I₂ + S → IS₂
Reaction with compounds containing Iodine
– Iodine can undergo reactions with compounds that already contain iodine. These reactions can involve substitution, oxidation, or reduction processes.
I₂ + H₂O₂ → 2HI + O₂
In this reaction, iodine reacts with hydrogen peroxide (H₂O₂) to form hydrogen iodide (HI) and oxygen gas (O₂). Iodine is reduced from its elemental state to iodide ions, while hydrogen peroxide acts as an oxidizing agent.
Reaction with Organic Compounds
Iodine reacts with organic compounds, particularly hydrocarbons, through a process called iodination. Iodine substitutes one or more hydrogen atoms in the organic compound, leading to the formation of alkyl iodides.
CH₄ + I₂ → CH₃I + HI
In this iodination, iodine substitutes one hydrogen atom in methane (CH₄), resulting in the formation of iodomethane (CH₃I) and hydrogen iodide (HI). The reaction is initiated by the presence of light or heat.
Astatine (At)
Astatine, the heaviest halogen, is radioactive and has limited availability for experimental study.
Due to its short half-life and rarity, the chemical properties and reactions of astatine are not as well-documented as the other halogens.
Compounds of Group 17 Elements (Halogens)
Binary Compounds with Other Elements
Halides
Halogens form binary compounds known as halides when they react with metals or nonmetals.
Metal halides are formed when halogens react with metals, while nonmetal halides are formed when halogens react with nonmetals.
For example: Sodium chloride (NaCl): Formed by the reaction of sodium (metal) with chlorine.
2Na + Cl₂ → 2NaCl
Hydrogen chloride (HCl): Formed by the reaction of hydrogen (nonmetal) with chlorine.
H₂ + Cl₂ → 2HCl
Ternary Compounds with Other Elements
Oxyhalides
Oxyhalides are compounds containing halogens and oxygen.
They are formed when halogens react with compounds containing oxygen, such as oxides or hydroxides.
For examples: Chlorine dioxide (ClO₂): Formed by the reaction of chlorine with sodium chlorite.
Cl₂ + 2NaClO₂ → 2NaCl + ClO₂
Acids and Their Properties of Group 17 Elements
Hydrogen Fluoride (HF)
Hydrogen fluoride is a weak acid that dissociates partially in water.
It is used in various industrial processes and as a reagent in organic chemistry.
Example: Dissociation of hydrogen fluoride:
HF ⇌ H⁺ + F⁻
Hydrogen Chloride (HCl)
Hydrogen chloride is a strong acid that completely dissociates in water.
It is widely used in chemical laboratories and industries.
For example: Dissociation of hydrogen chloride:
HCl → H⁺ + Cl⁻
Hydrogen Bromide (HBr)
Hydrogen bromide is a strong acid similar to hydrogen chloride.
It is used in various applications, including as a catalyst and in organic synthesis.
For example: Dissociation of hydrogen bromide
HBr → H⁺ + Br⁻
Hydrogen Iodide (HI)
Hydrogen iodide is a strong acid that readily dissociates in water.
It is used in organic synthesis and as a reducing agent.
Example: Dissociation of hydrogen iodide:
HI → H⁺ + I⁻
Interhalogen Compounds and Their Characteristics
Interhalogen compounds are compounds formed by the combination of two different halogens.
They have distinct properties and reactivity patterns.
For examples: Chlorine monofluoride (ClF): Formed by the reaction of chlorine with fluorine.
Cl₂ + F₂ → 2ClF
Bromine trifluoride (BrF₃): Formed by the reaction of bromine with fluorine.
Br₂ + 3F₂ → 2BrF₃
