Halogens are a group of elements that include fluorine, chlorine, bromine, and iodine. Even though they belong to the same family, these elements have unique properties. In this article, we’ll discuss the physical and chemical properties of halogens, including their atomic radius, melting and boiling points, electronegativity, volatility, and reactivity. We’ll also look at some of the ways halogens are used. Keep reading to learn more about these fascinating elements!
Halogens are a family of elements in the periodic table that all have five electrons in their outer p-subshell and tend to form -1 charged ions. You may also hear them referred to as group 7 or group 17. However, the technical term for group 7 according to the IUPAC is different, so it's simpler to stick with the name 'halogens' to prevent any confusion.
The halogen group is made up of five elements: fluorine (F), chlorine (Cl), bromine (Br), iodine (I), and astatine (At). Some scientists consider tennessine (Ts) to be a sixth halogen, but we won't be discussing it much because it shows some unusual behaviours, such as not forming negative ions, and because it is extremely unstable and costly to study. Astatine, in particular, is highly unstable and has a short half-life of just over eight hours, so its properties are mostly hypothetical. It's so radioactive that it hasn't been possible to collect a pure sample of it. Despite their differences, the halogens share many characteristics. In the following sections, we'll take a closer look at some of these shared properties.
The halogens are all non-metals. They show many of the physical properties typical of non-metals.They are poor conductors of heat and electricity.When solid, they are dull and brittle. They have low melting and boiling points.
However, I can provide a description of the halogens' distinct colours and states at room temperature. The halogens have characteristic colours: fluorine is a pale yellow gas, chlorine is a green gas, bromine is a dark red liquid that forms a red-brown vapour, and iodine is a grey-black solid that forms a purple vapour. Unlike most groups in the periodic table, the halogens can exist in all three states of matter (gas, liquid, and solid) at room temperature.
As you move down the group in the periodic table, the halogens increase in atomic radius. This is because they each have one more electron shell. For example, fluorine has the electron configuration 1s2 2s2 2p5, and chlorine has the electron configuration 1s2 2s2 2p6 3s2 3p5. Fluorine has just two main electron shells, whilst chlorine has three. Fluorine and chlorine, shown with their electron configurations. Notice how chlorine is a larger atom than fluorine. commons.wikimedia.org
As you can tell from their states of matter shown in the table earlier on, melting and boiling points increase as you go down the halogen group. This is because the atoms get larger and have more electrons. Because of this, they experience stronger van der Waals forces between molecules. These require more energy to overcome and so increase the element's melting and boiling points.
Volatility is very closely related to melting and boiling points - it is the ease with which a substance evaporates. From the data above, it is easy to see that the volatility of the halogens decreases as you move down the group. Once again, this is all thanks to van der Waals forces. As you move down the group, the atoms get larger and so have more electrons. Because of this, they experience stronger van der Waals forces, decreasing their volatility.
One trend in chemical properties within the halogen group is that electronegativity values increase as you move up the group. This means that the halogens closer to the top of the group (fluorine and chlorine) have a greater ability to attract electrons than those closer to the bottom of the group (iodine and astatine).
Another trend is that the reactivity of the halogens decreases as you move down the group. This is because as the atomic radius increases, the outer electrons become further away from the nucleus and are therefore less strongly attracted to it. This makes it harder for the halogens to attract electrons from other elements and participate in reactions. Finally, the boiling and melting points of the halogens increase as you move down the group. This is due to the increase in the size of the halogen atoms, which leads to stronger intermolecular forces between the molecules, requiring more energy to overcome and reach a higher temperature for boiling and melting.
That is correct. As you move down the halogen group, the atomic radius of the elements increases. This means that the outermost electrons are further away from the nucleus, which reduces the attractive forces between the electrons and the nucleus. Therefore, the ability of the halogens to attract a shared pair of electrons (i.e., their electronegativity) decreases as you move down the group.
This trend is also related to the polarity of molecules containing halogens. Since halogens have a high electronegativity, they tend to form polar covalent bonds with other elements. In a polar covalent bond, the electrons are shared unequally between the two atoms, resulting in a partial positive charge on one atom and a partial negative charge on the other. As the electronegativity of the halogen decreases, the bond becomes less polar, and the molecule becomes less polar overall. This trend can be observed in the physical properties of halogen compounds, such as their solubility in different solvents.
That is correct. Electron affinity is the amount of energy released when an atom gains an electron to form an anion. As you move down the halogen group, the atomic radius of the elements increases, which means that the incoming electron is further away from the nucleus and feels the attractive force of the nucleus less strongly. This reduces the energy released when the atom gains an electron and therefore decreases the electron affinity.
In addition, as you move down the group, the shielding effect of the inner electrons increases, which further reduces the attractive force between the nucleus and the incoming electron. This offsets the increase in nuclear charge, which also contributes to the decrease in electron affinity. It is worth noting that electron affinity values are always negative because energy is released when an atom gains an electron. This energy is usually in the form of heat or light, and the negative value of electron affinity reflects the fact that energy is being released from the system.
That is correct. Fluorine is the exception to the general trend of decreasing electron affinity as you move down the halogen group. This is because of the small size of the fluorine atom and the high electron density in the 2p subshell. The incoming electron is repelled by the other electrons in the subshell, which partially offsets the increased attractive force from the decreased atomic radius. As a result, the energy released when fluorine gains an electron is less than what would be predicted based on the trends in the rest of the group. It's also worth noting that while the electron affinity of fluorine is lower in magnitude than that of chlorine, it is still a highly electronegative element and forms strong polar covalent bonds with other elements. Fluorine's high electronegativity contributes to its ability to form compounds such as hydrogen fluoride, which is used in a variety of industrial processes.
That is correct. The reactivity of halogens is determined by their oxidizing and reducing abilities, which are related to their ability to gain or lose electrons. As you move down the halogen group, the oxidizing ability of the halogens decreases, while their reducing ability increases. This is because the atoms become larger and the electrons become further from the nucleus, making it easier for the halogens to lose electrons and harder for them to gain electrons.
However, as you noted, reactivity is not solely determined by electron affinity. Other factors, such as the size of the enthalpy changes involved in the reaction, also play a role. Fluorine, for example, may have a lower electron affinity than chlorine, but it is still the most electronegative element and has a very high oxidizing ability, making it extremely reactive. Overall, the reactivity of halogens follows a decreasing trend as you move down the group, but there may be exceptions due to other factors at play in specific reactions.
The final chemical property of halogens that we'll look at today is their bond strength. We'll consider both the strength of the halogen-halogen bond (X-X), and the hydrogen-halogen bond (H-X). Halogens forms diatomic X-X molecules. The strength of this halogen-halogen bond, also known as its bond enthalpy, generally decreases as you move down the group. However, fluorine is an exception - the F-F bond is much weaker than the Cl-Cl bond. Take a look at the graph below.
Bond enthalpy depends on the electrostatic attraction between the positive nucleus and the bonding pair of electrons. This in turn depends on the atom's number of unshielded protons, and the distance from the nucleus to the bonding electron pair. All halogens have the same number of electrons in their outer subshell and so have the same number of unshielded protons. However, as you move down the group in the periodic table, atomic radius increases, and so the distance from the nucleus to the bonding electron pair increases. This decreases the bond strength. Fluorine breaks this trend. Fluorine atoms have seven electrons in their outer shell. When they form diatomic F-F molecules, each atom features one bonding pair of electrons and three lone pairs of electrons. Fluorine atoms are so small that when two come together to form a F-F molecule, the lone pairs of electrons in one atom repel those in the other atom quite strongly - so much so that they decrease the F-F bond enthalpy.
Halogens can also form diatomic H-X molecules. The strength of the hydrogen-halogen bond decreases as you move down the group, as you can see from the graph below.
Once again, this is due to the increasing atomic radius of the halogen atom. As atomic radius increases, the distance between the nucleus and the bonding pair of electrons increases, and so bond strength decreases. But note that in this instance, fluorine follows the trend. Hydrogen atoms don't have any lone pairs of electrons, and so there isn't any additional repulsion between the hydrogen atom and the fluorine atom. Therefore, the H-F bond has the highest strength out of all of the hydrogen-halogen bonds. Thermal stability of hydrogen halides Let's take a moment to consider the relative thermal stabilities of hydrogen halides. As you move down the group in the periodic table, the hydrogen halides become less thermally stable. This is because the H-X bond decreases in strength and so is easier to break. Here's a table comparing the thermal stability and bond enthalpy of hydrogen halides:
To finish, we'll consider some of the uses of halogens. In fact, they have a number of applications.
Chlorine and bromine are used as disinfectants in a range of situations, from sterilising swimming pools and wounds to cleaning dishes and surfaces. In some countries, chicken meat is washed in chlorine to rid it of any harmful pathogens, such as salmonella and E. coli. Halogens can be used in lights. They improve the lifespan of the bulb.We can add halogens to drugs to make them dissolve in lipids more easily. This helps them cross through the phospholipid bilayer into our cells. Fluoride ions are used in toothpaste, where they form a protective layer around tooth enamel and prevent it from acid attack. Sodium chloride is also known as common table salt and is essential to human life. Similarly, we also need iodine in our body - it helps maintain optimum thyroid function.
Chlorofluorocarbons, also known as CFCs, are a type of molecule that were previously used in aerosols and refrigerators. However, they are now banned due to their negative effect on the ozone layer. You’ll find out more about CFCs in Ozone Depletion.
Properties of Halogens - Key takeaways The halogens are a group of elements in the periodic table, all with five electrons in their outer p-subshell. They commonly form ions with a charge of -1 and are also known as group 7 or group 17.The halogens are non-metals and form diatomic molecules. As you move down the halogen group in the periodic table: Atomic radius increases. Melting and boiling points increase. Volatility decreases. Electronegativity generally decreases. Reactivity decreases. The X-X and H-X bond strength generally decreases. Halogens aren’t very soluble in water, but are soluble in organic solvents such as alkanes.We use halogens for a variety of purposes, including sterilisation, lighting, medicines, and toothpaste.
What are the similar properties of halogens?
In general, halogens have low melting and boiling points, high electronegativities, and are sparingly soluble in water. Their properties show trends as you move down the group. For example, atomic radius and melting and boiling points increase down the group whilst reactivity and electronegativity decrease.
What are the chemical properties of halogens?
In general, halogens have high electronegativities - fluorine is the most electronegative element in the periodic table. Their electronegativity decreases as you go down the group. Their reactivity also decreases as you go down the group. Halogens all take part in similar reactions. For example, they react with metals to form salts and with hydrogen to form hydrogen halides. Halogens are sparingly soluble in water, tend to form negative anions, and are found as diatomic molecules.
What are the physical properties of halogens?
Halogens have low melting and boiling points. As solids they are dull and brittle, and they are poor conductors.
What are the uses of halogens?
Halogens are commonly used to sterilise things such as drinking water, hospital equipment, and work surfaces. They are also used in lightbulbs. Fluorine is an important ingredient in toothpaste as it helps protect our teeth from cavities whilst iodine is essential for supporting thyroid function.
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