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Group 17: The halogens

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Intros
Lessons
  1. Halogens: A summary
  2. Introduction to halogens.
  3. Properties of the halogens.
  4. How do the properties change going down the group? (Trends in properties)
  5. Explaining the trends in properties.
  6. Reaction of halogens with group 1 and 2 metals.
  7. Halogen displacement reactions.
  8. Halogens as oxidising agents.
  9. Halogens and disproportionation reactions.
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Examples
Lessons
  1. Recall the properties of the halogens.
    Sort the halogens according to the properties listed below:
    1. Reactivity; lowest to highest.
    2. State at room temperature; solids, then liquids, then gases.
  2. Explaining the trend in properties of the halogens.
    A student in a school laboratory has a solution of sodium bromide (NaBr) in a beaker. To this, she adds an equal amount of chlorine water (chlorine dissolved in water).
    1. What observations would she make from the reaction?
    2. Write a chemical equation to describe this reaction, and explain why this observation and reaction occurred.
    3. Explain why no observed change would be seen if sodium bromide (NaBr) was reacted with iodine (I2) in solution.
Topic Notes
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Introduction to Group 17: The Halogens

Welcome to our exploration of the halogens, a fascinating group of elements in the periodic table! These reactive elements, found in Group 17, are some of the most intriguing in chemistry. The halogens include fluorine, chlorine, bromine, iodine, and astatine. They're known for their high reactivity and tendency to form salts when combined with metals. In the periodic table, you'll find them on the right side, just one column away from the noble gases. Their unique properties make them essential in various applications, from water purification to medical imaging. Before we dive deeper, I encourage you to watch our introductory video. It provides a visual overview of these elements and their behavior, which will greatly enhance your understanding of the topic. As we continue, we'll explore each halogen's characteristics and their roles in our daily lives. Ready to uncover the secrets of these powerful elements?

Properties and Characteristics of Halogens

Introduction to Halogens

Halogens are a group of highly reactive elements found in Group 17 of the periodic table. These elements, including fluorine, chlorine, bromine, iodine, and astatine, share unique properties that make them essential in various chemical processes and applications. Understanding the key properties of halogens is crucial for anyone studying chemistry or working with these elements.

Physical States and Colors of Halogens

One of the fascinating aspects of halogens is their diverse physical states at room temperature. Fluorine and chlorine exist as gases, bromine is a liquid, while iodine and astatine are solids. This progression in physical states correlates with their increasing atomic numbers. Each halogen also has a distinctive color: fluorine is pale yellow, chlorine is yellowish-green, bromine is reddish-brown, and iodine appears as shiny, purple-black crystals that produce a violet vapor when heated.

Electron Configuration and Chemical Behavior

The electron configuration of halogens plays a crucial role in their chemical behavior. All halogens have seven valence electrons in their outermost shell, giving them an ns²np configuration. This unique arrangement makes halogens highly reactive, as they eagerly seek to gain one electron to achieve a stable octet configuration. The tendency to gain electrons makes halogens strong oxidizing agents and explains their high electronegativity values.

Reactivity and Electronegativity

Halogens are among the most reactive elements in the periodic table. Their high reactivity stems from their strong tendency to form ionic or covalent bonds with other elements. The electronegativity of halogens decreases as we move down the group, with fluorine being the most electronegative element in the entire periodic table. This trend in electronegativity directly influences their reactivity, with fluorine being the most reactive and iodine the least among the commonly encountered halogens.

Comparison of Reactivity Among Halogens

The reactivity of halogens follows a clear trend: fluorine > chlorine > bromine > iodine > astatine. This order is evident in their ability to displace each other from compounds. For example, chlorine can displace bromine from bromide salts, and bromine can displace iodine from iodide salts. This displacement reaction is a practical demonstration of the relative reactivity of halogens and is often used in laboratory experiments to identify halide ions.

Formation of Diatomic Molecules

A key characteristic of halogens is their tendency to form diatomic molecules (X) in their elemental state. This property is due to the covalent bonding between two halogen atoms, sharing their unpaired electrons to achieve a stable electron configuration. The strength of these diatomic bonds decreases down the group, which correlates with the increasing size of the atoms and the decreasing electronegativity.

Solubility and Reactions with Water

Halogens exhibit varying degrees of solubility in water. Fluorine reacts vigorously with water, while chlorine, bromine, and iodine are moderately soluble, forming acidic solutions. The reaction of halogens with water produces hydrohalic acids and hypohalous acids, which have important applications in water treatment and disinfection processes.

Applications of Halogens

The unique properties of halogens make them invaluable in numerous applications. Chlorine is widely used in water purification and as a disinfectant. Fluorine compounds are essential in the production of non-stick cookware and pharmaceuticals. Bromine is used in flame retardants and pharmaceutical manufacturing. Iodine plays a crucial role in medicine, particularly in thyroid health and as an antiseptic.

Conclusion

The halogens, with their distinctive properties ranging from physical states to high reactivity, play a significant role in chemistry and everyday life.

Reactivity and Chemical Behavior of Halogens

Halogens, comprising fluorine, chlorine, bromine, iodine, and astatine, are known for their high reactivity in chemical reactions. This group of elements, located in Group 17 of the periodic table, exhibits unique properties that make them essential in various chemical processes. Among the halogens, fluorine stands out as the most reactive, a characteristic that stems from its atomic structure and electronegativity.

The exceptional reactivity of fluorine can be attributed to several factors. Firstly, fluorine has the smallest atomic radius among halogens, allowing it to form stronger bonds with other elements. Additionally, its high electronegativity means it has a strong tendency to attract electrons, making it highly reactive in chemical reactions. This fluorine reactivity is so intense that it can even react with noble gases under certain conditions, a feat unmatched by other elements.

Halogens participate in diverse chemical reactions, forming both ionic and covalent bonds depending on the elements they interact with. When halogens react with metals, they typically form ionic bonds, resulting in the creation of metal halides. For instance, sodium chloride (NaCl), commonly known as table salt, is formed when sodium reacts with chlorine. In this reaction, the sodium atom loses an electron to the chlorine atom, forming Na+ and Cl- ions, which are then held together by electrostatic forces.

On the other hand, when halogens react with non-metals, they tend to form covalent bonds. A prime example is the formation of hydrogen chloride (HCl) when hydrogen reacts with chlorine. In this case, the atoms share electrons to form a stable molecule. These covalent halogen compounds are often gases or liquids at room temperature and play crucial roles in various industrial and laboratory processes.

The concept of oxidation numbers is fundamental in understanding halogen reactions and their compounds. Oxidation numbers represent the degree of oxidation of an atom in a chemical compound, indicating the hypothetical charge an atom would have if all bonds were ionic. In halogen compounds, halogens typically have an oxidation number of -1 when they form halide ions. For example, in sodium chloride (NaCl), chlorine has an oxidation number of -1, while sodium has +1.

However, halogens can exhibit various oxidation states in different compounds. For instance, chlorine can have oxidation numbers ranging from -1 to +7, as seen in compounds like hypochlorite (ClO-, +1), chlorite (ClO2-, +3), chlorate (ClO3-, +5), and perchlorate (ClO4-, +7). These varying oxidation states contribute to the diverse chemistry of halogens and their compounds.

The formation of halide ions is a crucial aspect of halogen chemistry. When halogens gain an electron, they form negatively charged halide ions (F-, Cl-, Br-, I-). These ions play essential roles in biological systems, industrial processes, and environmental chemistry. For example, chloride ions are vital for maintaining proper fluid balance in living organisms, while fluoride ions are added to drinking water and toothpaste to prevent tooth decay.

In conclusion, the high reactivity of halogens, particularly fluorine, makes them fascinating subjects in chemistry. Their ability to form both ionic and covalent bonds, coupled with their variable oxidation states, contributes to their versatility in chemical reactions. Understanding these concepts is crucial for comprehending the behavior of halogens in various chemical processes and their applications in everyday life and industry. From the salt on our tables to the chlorine in swimming pools, halogens and their compounds continue to play significant roles in our world, demonstrating the importance of these highly reactive elements in chemistry and beyond.

Trends in Halogen Properties

The halogen group, consisting of fluorine, chlorine, bromine, iodine, and astatine, exhibits fascinating trends in properties as we move down the periodic table. These trends in halogen properties, particularly in atomic size, electronegativity, and reactivity, are crucial for understanding their behavior in chemical reactions and their role in various applications.

One of the most noticeable halogen trends is the increase in atomic size as we descend the group. This trend is primarily due to the addition of new electron shells with each successive element. For instance, fluorine, the smallest halogen, has only two electron shells, while iodine has five. This increase in atomic size has significant implications for the other properties of halogens.

The electronegativity trend in halogens is particularly interesting. Unlike the atomic size, electronegativity decreases as we move down the group. Fluorine, at the top of the group, is the most electronegative element in the entire periodic table. This high electronegativity is due to its small atomic size and the strong attraction between its nucleus and valence electrons. As we move down to chlorine, bromine, and iodine, the electronegativity gradually decreases.

The decrease in electronegativity is closely related to the concept of electron shielding. As the atomic size increases, the outer electrons are further from the nucleus and experience greater shielding from the inner electron shells. This shielding effect reduces the attraction between the nucleus and the valence electrons, resulting in lower electronegativity for the larger halogens.

The reactivity trend in halogens is closely tied to their electronegativity and atomic size. Generally, the reactivity of halogens decreases as we move down the group. Fluorine, being the most electronegative and smallest, is also the most reactive halogen. It can react vigorously with almost all elements, including some noble gases under specific conditions. In contrast, iodine, at the bottom of the group, is the least reactive halogen.

To illustrate these trends, let's compare fluorine and iodine. Fluorine has a much smaller atomic radius than iodine, higher electronegativity, and greater reactivity. These differences are evident in their physical and chemical properties. For example, fluorine exists as a pale yellow gas at room temperature, while iodine is a solid with a characteristic violet color. In chemical reactions, fluorine forms stronger bonds and reacts more readily than iodine.

The atomic structure plays a crucial role in these halogen trends. As we move down the group, each element adds a new electron shell, increasing the distance between the nucleus and the valence electrons. This increased distance, combined with the shielding effect of inner electrons, weakens the nuclear attraction on the outermost electrons. Consequently, it becomes easier to remove an electron from a larger halogen atom, affecting its chemical behavior.

These trends have significant implications in various fields. In organic chemistry, the reactivity trend influences the strength of carbon-halogen bonds, with C-F bonds being the strongest and C-I bonds the weakest. This property is utilized in the synthesis of various organic compounds. In materials science, the varying properties of halogens are exploited in the development of new materials, from fluoropolymers like Teflon to iodine-based liquid crystals used in displays.

Understanding these halogen trends is also crucial in environmental science. The reactivity of chlorine, for instance, plays a significant role in the depletion of the ozone layer. On the other hand, iodine's lower reactivity makes it an essential nutrient in biological systems, particularly in thyroid function.

In conclusion, the trends in halogen properties - increasing atomic size, decreasing electronegativity, and decreasing reactivity as we move down the group - are fundamental concepts in chemistry. These trends, rooted in the atomic structure and electron shielding effects, provide a framework for predicting and understanding the behavior of halogens in various chemical and physical processes. By comparing different halogens, we can appreciate the gradual change in properties and their impact on the elements' roles in nature and technology.

Halogen Compounds and Their Applications

Introduction to Halogen Compounds

Halogen compounds are a diverse group of chemical substances that contain one or more halogen elements: fluorine, chlorine, bromine, iodine, or astatine. These compounds play crucial roles in our daily lives, industries, and medical applications. From common table salt to complex organic molecules, halogens are ubiquitous and indispensable in modern society.

Common Halide Salts

Halide salts are among the most prevalent halogen compounds. Sodium chloride, or common table salt, is perhaps the most well-known example. This essential compound not only enhances food flavor but also plays a vital role in maintaining proper bodily functions. Other important halide salts include potassium chloride, used as a salt substitute and in fertilizers, and calcium chloride, which finds applications in road de-icing and food preservation.

Fluoride Compounds

Fluoride compounds are renowned for their dental health benefits. Sodium fluoride and stannous fluoride are commonly added to toothpaste and drinking water to prevent tooth decay. In industry, hydrofluoric acid is used for etching glass and in the production of fluoropolymers like Teflon. These materials are prized for their non-stick and heat-resistant properties, making them invaluable in cookware and industrial applications.

Iodine and Its Compounds

Iodine and its compounds are essential in medicine and public health. Iodine tincture is a common antiseptic used for wound disinfection. Potassium iodide is added to table salt in many countries to prevent iodine deficiency disorders. In medical imaging, iodine-based contrast agents are used to enhance the visibility of blood vessels and organs during X-ray and CT scans.

Bromine Compounds

While less common in everyday life, bromine compounds have significant industrial applications. Methyl bromide, though now restricted due to environmental concerns, was once widely used as a pesticide. Brominated flame retardants are added to many plastics, textiles, and electronic equipment to reduce flammability. In pharmaceuticals, bromine is found in several medications, including those used to treat epilepsy and anxiety disorders.

Chlorine and Its Derivatives

Chlorine compounds are extensively used in water treatment and sanitation. Sodium hypochlorite, commonly known as bleach, is a powerful disinfectant used in households and public spaces. In industry, chlorine is a key component in the production of PVC (polyvinyl chloride), a versatile plastic used in construction, plumbing, and consumer goods. Chlorofluorocarbons (CFCs) were once widely used as refrigerants and propellants but have been phased out due to their ozone-depleting properties.

Organic Halogen Compounds

Organic halogen compounds, or organohalides, are carbon-based molecules containing one or more halogen atoms. These compounds have diverse applications in pharmaceuticals, agrochemicals, and materials science. For example, many antibiotics contain chlorine atoms, which enhance their effectiveness against bacteria. Fluorinated organic compounds are used in the development of new drugs due to their unique properties, such as increased lipophilicity and metabolic stability.

Environmental and Safety Considerations

While halogen compounds offer numerous benefits, their use also raises environmental and safety concerns. The production and disposal of certain halogenated compounds can lead to environmental pollution and health risks. Efforts are ongoing to develop more environmentally friendly alternatives and improve the safe handling and disposal of halogen-containing products.

Conclusion

Halogen compounds are integral to modern life, with applications spanning from simple table salt to complex pharmaceuticals. Their unique properties make them invaluable in various industries, including medicine, agriculture, and materials science. As research continues, new applications and safer, more sustainable halogen compounds are likely to emerge, further expanding their importance in our daily lives and technological advan

Halogen Displacement Reactions

Halogen displacement reactions are a fascinating aspect of chemical reactivity, showcasing the principle of more reactive elements replacing less reactive ones in compounds. These reactions are particularly important in the study of halogens, which include fluorine, chlorine, bromine, and iodine. Understanding these displacement reactions is crucial for both theoretical chemistry and practical applications in industry and laboratory settings.

At the core of halogen displacement reactions is the concept of the halogen reactivity series. This series arranges halogens in order of their reactivity, with fluorine being the most reactive, followed by chlorine, bromine, and iodine. The key principle is that a more reactive halogen can displace a less reactive one from its compounds. This phenomenon is based on the relative electron affinities and oxidizing powers of the halogens.

Let's explore some examples of halogen displacement reactions with balanced equations:

  1. Chlorine displacing bromine from potassium bromide:
    Cl + 2KBr 2KCl + Br
  2. Bromine displacing iodine from sodium iodide:
    Br + 2NaI 2NaBr + I
  3. Chlorine displacing iodine from potassium iodide:
    Cl + 2KI 2KCl + I

In each of these reactions, we observe that the more reactive halogen (on the left side of the equation) replaces the less reactive halogen in the compound. This results in the formation of a new halide compound and the liberation of the less reactive halogen as an element.

It's important to note that these reactions are not reversible under standard conditions. For instance, iodine cannot displace chlorine from potassium chloride because iodine is less reactive than chlorine.

The practical applications of halogen displacement reactions are numerous and significant in both industry and laboratory settings:

  • In the production of halogens: These reactions are used to extract less reactive halogens from their naturally occurring compounds. For example, bromine is often produced by displacing it from seawater using chlorine.
  • In organic synthesis: Halogen displacement reactions are valuable in creating organic halides, which are important intermediates in many synthetic processes.
  • In water treatment: Chlorine's ability to displace other halogens is utilized in water purification processes to remove bromide and iodide ions.
  • In analytical chemistry: These reactions are used to identify and quantify halide ions in unknown samples.
  • In the manufacture of pharmaceuticals: Many drug synthesis processes involve halogen displacement steps to introduce specific halogen atoms into molecular structures.

In laboratory settings, halogen displacement reactions serve as excellent demonstrations of chemical reactivity and the concept of oxidation-reduction reactions. They provide visual evidence of chemical change, often accompanied by color changes that make them particularly engaging for students learning about chemical reactions.

Understanding the principles behind halogen displacement reactions is crucial for chemists and chemical engineers. These reactions highlight the importance of electronegativity and electron affinity in determining chemical behavior. They also demonstrate how the periodic table's organization reflects and predicts chemical properties and reactivity trends.

In conclusion, halogen displacement reactions are a prime example of how understanding fundamental chemical principles can lead to practical applications across various fields. From industrial processes to laboratory techniques, these reactions play a vital role in modern chemistry. By mastering the concept of the halogen reactivity series and the principles governing these displacement reactions, chemists can predict outcomes, design new synthetic routes, and develop innovative applications in fields ranging from materials science to environmental protection.

Conclusion: The Significance of Halogens in Chemistry

Halogens stand out as a crucial group of elements in chemistry, characterized by their unique properties and widespread applications. These reactive nonmetals play vital roles in various chemical reactions, from industrial manufacturing to biological systems. The introduction video provides an essential foundation for understanding halogen importance, showcasing their distinctive features such as high electronegativity and tendency to form ionic compounds. By exploring halogen chemistry further, students and researchers can unlock a wealth of knowledge applicable to fields like pharmaceuticals, materials science, and environmental studies. The versatility of halogens in forming compounds and their impact on chemical reactions make them indispensable in modern chemistry. As we continue to advance in scientific research, the significance of halogens only grows, emphasizing the need for ongoing study and exploration of these fascinating elements. Delving deeper into halogen chemistry opens doors to innovative applications and breakthroughs in various scientific disciplines.

Halogens: A Summary

Introduction to Halogens

Step 1: Understanding the Basics of Halogens

Halogens, also known as Group 17 elements, are a group of highly reactive non-metals. This group includes fluorine, chlorine, bromine, iodine, and astatine. The term "halogen" is derived from Greek, meaning "salt-forming," which is a key characteristic of these elements. They are known for their ability to form salts when they react with metals.

Step 2: Properties of Halogens

The properties of halogens are quite distinct and vary as you move down the group. They are known for their high reactivity, which decreases from fluorine to astatine. Halogens have seven electrons in their outer shell, making them one electron short of a full octet, which explains their high reactivity. They readily gain an electron to achieve a stable electronic configuration.

Step 3: Trends in Properties

As you move down Group 17, several trends can be observed in the properties of halogens:

  • Reactivity: Reactivity decreases from fluorine to astatine. Fluorine is the most reactive, while astatine is the least.
  • Electronegativity: Electronegativity decreases down the group. Fluorine is the most electronegative element in the periodic table.
  • Melting and Boiling Points: Both melting and boiling points increase as you move down the group.
  • Color: The color of halogens becomes darker as you go down the group. Fluorine is pale yellow, chlorine is greenish-yellow, bromine is reddish-brown, and iodine is violet.

Step 4: Electronic Structure and Configuration

The electronic structure of halogens plays a crucial role in their chemical behavior. Each halogen has seven electrons in its outermost shell. This configuration makes them highly reactive as they seek to gain one more electron to achieve a stable octet. The number of inner electron shells increases as you move down the group, which affects their reactivity and other properties.

Step 5: Key Reactions of Halogens

Halogens are known for their reactivity and ability to form a wide variety of compounds. Some key reactions include:

  • Formation of Salts: Halogens react with metals to form ionic salts. For example, sodium reacts with chlorine to form sodium chloride (table salt).
  • Displacement Reactions: A more reactive halogen can displace a less reactive halogen from its compounds. For example, chlorine can displace bromine from sodium bromide.
  • Reactions with Hydrogen: Halogens react with hydrogen to form hydrogen halides, such as hydrogen chloride (HCl).

Step 6: Historical Context and Discovery

Historically, halogens were known in their compound forms long before they were isolated as elements. Their high reactivity made it challenging to isolate them in their elemental forms. For instance, fluorine is extremely reactive and was only isolated after several scientists faced significant challenges, including fatalities. Chlorine was used as a chemical weapon during World War I due to its toxic properties.

Step 7: Practical Applications and Safety

Despite their reactivity and potential hazards, halogens have numerous practical applications. Chlorine is widely used in water treatment and as a disinfectant. Fluorine compounds are used in toothpaste to prevent dental cavities. However, due to their toxic nature, handling halogens requires strict safety measures to prevent harmful exposure.

FAQs

  1. Why are halogens considered highly reactive elements?

    Halogens are highly reactive due to their electron configuration. They have seven valence electrons, making them eager to gain one more electron to achieve a stable octet. This high electronegativity and electron affinity make them readily react with many elements, particularly metals, to form ionic compounds.

  2. What is the order of reactivity for halogens?

    The reactivity of halogens decreases as we move down the group in the periodic table. The order from most to least reactive is: Fluorine > Chlorine > Bromine > Iodine > Astatine. This trend is due to decreasing electronegativity and increasing atomic size as we move down the group.

  3. How do halogens form compounds with metals?

    Halogens react with metals to form ionic compounds called halides. In these reactions, the metal loses electrons to the halogen, forming a positively charged metal ion and a negatively charged halide ion. For example, sodium reacts with chlorine to form sodium chloride (NaCl), commonly known as table salt.

  4. What are some common applications of halogen compounds?

    Halogen compounds have numerous applications in everyday life and industry. Some examples include: sodium chloride (table salt) for food seasoning, fluoride in toothpaste for preventing tooth decay, chlorine for water purification, bromine in flame retardants, and iodine in medical imaging and as a disinfectant.

  5. Why is fluorine the most reactive halogen?

    Fluorine is the most reactive halogen due to its small atomic size and high electronegativity. It has the smallest atomic radius among halogens, allowing its nucleus to exert a stronger pull on electrons. Additionally, fluorine's high electronegativity means it has a strong tendency to attract electrons from other atoms, making it extremely reactive in chemical reactions.

Prerequisite Topics

Understanding Group 17: The halogens requires a solid foundation in key chemistry concepts. Two crucial prerequisite topics are particularly relevant to this subject: the history and development of the periodic table and periodic trends in atomic radius. These topics provide essential context and principles that will enhance your comprehension of the halogens and their unique properties.

The history and development of the periodic table is fundamental to understanding the organization of elements, including the halogens. This knowledge helps you appreciate why the halogens are grouped together and how their position in the table reflects their shared characteristics. The periodic table's structure, based on electron configurations and chemical properties, directly influences the behavior of Group 17 elements.

Grasping periodic trends, particularly atomic radius, is crucial when studying the halogens. The trend in atomic radius across the group significantly impacts the reactivity and properties of these elements. As you move down Group 17, from fluorine to astatine, the atomic radius increases, affecting bond strengths, electronegativity, and overall reactivity patterns.

The concept of periodic table trends is especially relevant to the halogens. These trends explain why fluorine, despite being the smallest halogen, is the most reactive. Understanding how electron configurations change across the group helps predict and explain the chemical behavior of each halogen element.

Moreover, knowledge of atomic radius variations within Group 17 is essential for comprehending the physical and chemical properties of halogens. This understanding extends to concepts like ionization energy, electron affinity, and bond strengths all crucial aspects when studying halogen chemistry.

By mastering these prerequisite topics, you'll be better equipped to explore the fascinating world of halogens. You'll understand why these elements share similar properties, such as their tendency to form single covalent bonds and their high electronegativity. This foundation will also help you grasp more advanced concepts related to halogens, such as their role in organic chemistry, their importance in industrial applications, and their environmental impact.

In conclusion, a solid grasp of the periodic table's development and atomic radius trends will significantly enhance your study of Group 17 elements. These prerequisite topics provide the necessary context and principles to fully appreciate the unique characteristics and behaviors of the halogens, setting a strong foundation for more advanced chemistry studies.

In this lesson, we will learn:
  • To recall properties of the halogens
  • To understand the trend in properties found in the halogens
  • To apply knowledge of electronic structure and bonding to explain the trends in halogens.
  • To know the key reactions of the halogens.

Notes:

  • We have already seen that the Periodic Table is arranged, top-left to bottom-right, by proton number and number of outer shell electrons. The number of outer shell electrons dictates the chemical properties of an element.

  • Just like with the alkali metals, the halogens are another example of a well-studied group of elements which display trends in their common properties as you go down the group. As they are reactive elements, they were known and studied in their compounds before being isolated in their reactive, toxic elemental forms.

  • The halogens have the following properties:
    • They are non-metals stable as diatomic molecules (this means at room temperature and pressure, they exist as molecules made of two atoms, e.g. Cl2).
    • They have a valence of 1 and form covalent bonds with non-metals atoms, or ionic bonds with metal atoms. Halogen ions will usually have a single negative charge (X-), where they are known as halides.
    • They are colored.
    • They have relatively low melting and boiling points compared to other non-metals (except the noble gases).
    • Halogens in elemental form are relatively toxic, reactive substances.
    • They do not conduct heat or electricity.
    • They are brittle as solids.

  • As you go down the group, the properties of the elements change in the following ways:
    • The melting and boiling point gets higher – starting as gases, bromine is a liquid while iodine is a solid.
    • The color of the halogens gets darker – fluorine is pale yellow, followed by green chlorine, brown/purple bromine and purple iodine.
    • Electronegativity decreases down the group. The smallest halogen, fluorine, is the most electronegative element in the periodic table.
    • The halogens get less reactive – fluorine, top of the group, is the most reactive element known. Iodine is the least reactive halogen (besides astatine which is often ignored because it is extremely rare).

  • As with group 1 and 2, the trends in properties and GENERAL reactivity in group 7 can be explained by their electronic configuration:
    • The reason that melting and boiling points increase down the group is because the intermolecular forces between the halogen molecules (e.g. F2, Cl2, Br2) get stronger down the group.
      The increasingly large halogen atoms with more electrons produce stronger van der Waals forces between the molecules, so more heat energy is needed to overcome them.
      This is why fluorine and chlorine are gases at room temperature, bromine is a liquid and iodine a solid.
    • The reason that electronegativity decreases down the group is because each halogen further down the group has an extra inner electron shell shielding the nucleus from the outer electron shell.
      As explained in Periodic Trends: Electronegativity more inner shells make it more difficult to attract electrons into (and keep them held in) the outer shell, which is increasingly further away from the nucleus. This is why fluorine (with only one inner shell shielding the nucleus) is the single most electronegative element in the periodic table, and electronegativity drops from there.
    • Electronegativity decreasing down the group is the reason why reactivity decreases down the group! Halogens react by ‘pulling in’ electrons to their outer shell to complete it. Electronegativity measures this; how easily does an atom attract bonding electrons into its outer shell and hold onto them?
      Fluorine is the most electronegative halogen, so it is the most reactive halogen. Iodine is the least electronegative (except astatine), so it is the least reactive.

  • As one of the more reactive groups of elements, there are a variety of reactions the halogens take part in:
    • The halogens react well with group 1 and 2 metals because these have electron configurations that complement the halogens. The metals react by losing electrons; the halogens react by gaining them. These are vigorous, exothermic reactions.


      • Group 1 with halogens:
      2M + X2  \, →  \, 2MX


      Group 2 with halogens:
      M + X2  \, →  \, MX2

      In both equations:
      M = group 1 or group 2 metal.
      X = halogen (F2, Cl2, Br2, I2).

    • A more reactive halogen can displace a less reactive halogen in a salt. This is known as a displacement reaction.

      • X2 + 2MY  \, →  \, 2MX + Y2

      X = F, Cl, Br or I,
      Y = halogen atom below X in the group.
      X = For example:
      Cl2 + 2NaBr  \, →  \, 2NaCl + Br2

      This reaction works as long as the ‘X2’ halogen is higher up the group than the halogen in the ‘MY’ halide salt, where M is the metal and Y is the halide ion.
      If this reaction happens in water (‘aqueous solution’), it’s more accurate to write this reaction ignoring the metal ions because they are just spectator ions. Spectator ions don’t take part in the reaction – this is just an oxidation of the halide (Cl-, Br- or I-) ions by a halogen molecule (see the point just below).

    • The halogens are good oxidising agents, and more reactive halogens can oxidise halide ions further down the group.
      A general equation:
      X2 + 2Y- (aq)  \, →  \, 2X- (aq) + Y2

      Where X is a halogen atom higher up the group than Y.

      This reaction works as long as the halogen (X2) reactant is higher up the group than the halide ions (2Y-) they are oxidising. You can think of it as electrons (or negative charge) “moving up” the group.
      For example, Cl2 will react with Br- to produce chloride ions and bromine:

      Cl2 (g) + 2Br- (aq)  \, →  \, 2Cl- (aq) + Br2 (l)

      • Fluorine is such a strong oxidising agent that this reaction doesn’t practically work for fluorine because it literally oxidises the water instead of the halide ions!


      This reaction happens less readily going down the group. This is an effect of the enthalpy change of X2  \, →  \, 2X- (aq) getting less exothermic (more positive). The reasons given for why this happens is misleading though; there is a ‘short answer’ and a more accurate long answer.
      • The short answer is that in the halogen group, electronegativity is higher and there is less electron shielding of the nucleus at the top of the group. This means fluorine/chlorine is better at attracting electrons into its outer shell and forming an ion than the larger halogens like bromine/iodine below. Depending on your course/syllabus, this answer will probably be enough, but it isn’t completely true.

      • The short answer is inaccurate because it talks about only one part of the reaction called the electron affinity, and the enthalpy data2 doesn’t back it up. Fluorine isn’t better than chlorine at ‘attracting electrons into its outer shell’. The total enthalpy change still shows fluorine as the most reactive because…

      • The long answer is that this reaction has multiple enthalpy changes involved.
        The first is the enthalpy of atomisation: the energy needed to break all bonds within a substance into separate atoms in the gas phase. For the halogens this change can be written as:

        • X2 (g)  \, →  \, 2X (g)

        In elemental fluorine and chlorine, this is half of the bond enthalpy1 because they are diatomic gases at room temperature. F-F is a weaker bond than Cl-Cl, so F is more readily atomised.
        The second is the electron affinity: the energy released when an electron joins a gaseous atom to form a negative ion. This change can be written as:

        2X (g)  \, →  \, 2X- (g)

        Fluorine’s electron affinity is actually less than chlorine’s.2 Fluorine is a small atom with very high electron density where the bonding electrons are being attracted into. This decreases its electron affinity to being lower than that of Cl. Cl  \, →  \, Cl- releases slightly more energy than F  \, →  \, F-, so Cl is more readily ionized.
        The third is the enthalpy of hydration: the energy released when an ionic substance dissolves in and interacts with water. This change can be written as:

        2X- (g)  \, →  \, 2X- (aq)

        The enthalpy of hydration is much higher when F- ions dissolve than when Cl- ions dissolve.3 In other words, F- interacting with water releases A LOT more energy than Cl- interacting with water.
      In summary, when adding the three enthalpy values1,2,3 together to get total enthalpy change, you find fluorine is the most exothermic by far. The change gets less exothermic going down the group in a consistent manner.

    • The halogens, particularly chlorine, also react in disproportionation reactions. These are reactions where atoms of the same substance are both oxidised and reduced. In a reaction with Cl2, one atom is reduced when forming one product and the other Cl atom is oxidised when forming part of another product.

      The disproportionation of chlorine with water is below:

      • Cl2 + H2O  ⇌ \, \rightleftharpoons \, HClO + HCl

      The disproportionation of chlorine with cold sodium hydroxide is below:

      Cl2 + 2NaOH  \, →  \, NaClO + NaCl + H2O

      This reaction with cold sodium hydroxide produces sodium hypochlorite (NaClO), which is the active ingredient in bleach products. Hypochlorite salts and elemental chlorine (in very small quantities) are both used in water treatment as disinfectants.

      Disproportionation reactions are best understood using oxidation numbers (see Oxidation number for more):
      • Like all elemental forms, Cl in Cl2 has an oxidation state of zero.
      • Na, being in group 1, takes a +1 oxidation state; so does H when bonded to non-metals.
      • Oxygen, being the most electronegative element besides fluorine, has an oxidation state of -2 in compounds.
      • In the reactions above, no oxidation numbers change except for chlorine: chlorine goes from two atoms with an oxidation number of zero to one atom with an oxidation number of -1 and another with +1.

      The reaction of chlorine with hot sodium hydroxide is slightly different, producing sodium chlorate instead:
      3Cl2 + 6NaOH  \, →  \, 5NaCl + NaClO3 + 3H2O

      Just like in the reaction with water and cold sodium hydroxide, the only oxidation state change here is with the chlorine atoms.
      • 5 of the 6 Cl atoms in the reactants formed sodium chloride, going from an oxidation number of 0 to -1.
      • The last Cl atom forms sodium chlorate, with a change in oxidation number of 0 to +5.
      To balance this equation, memorize the products and then look at the oxidation number changes:
      • An oxidation number just tells you about electrons being transferred. Electrons don’t get created or destroyed, so the changes in oxidation number in a reaction must cancel to zero.
      • When Cl2 reacts with hot NaOH, NaClO3 forms which is an oxidation number change of 0 to +5 in the chlorine atom. If one atom of Cl changed 0  \, →  \, +5, there must be a 0  \, →  \, -5 change in the other chlorine atoms in total. In NaCl, chlorine has a 1- oxidation state, so it must be 5 moles of NaCl to make -5 overall change.
      • This gives you 5NaCl for every NaClO3. To produce this, you would need 3Cl2 for the chlorine atoms, 6 NaOH for 6 Na atoms and the leftover H and O make 3 H2O molecules.
      See below for the full equations:

      This disproportionation reactions can happen with bromine and iodine as well. The reaction to produce the chlorate/bromate/iodate (NaXO3, X = Cl, Br, I) occurs more readily going down the group.


      References:
      1: https://labs.chem.ucsb.edu/zakarian/armen/11---bonddissociationenergy.pdf - source for bond enthalpy data.
      2: http://hyperphysics.phy-astr.gsu.edu/hbase/Chemical/eleaff.html#c1 - source for electron affinity data.
      3: http://www.wiredchemist.com/chemistry/data/enthalpies-hydration - source for enthalpy of hydration.