Group VII properties Revision Notes

Group VII of the periodic table contains the halogens, a family of elements including chlorine (Cl), bromine (Br), and iodine (I). These elements are non-metals and exist naturally as diatomic molecules, meaning each molecule consists of two atoms (e.g., Cl₂, Br₂, I₂). This diatomic nature arises because halogen atoms share electrons to achieve a full outer electron shell, making the molecules more stable than individual atoms.

At room temperature and pressure, these halogens exist in different physical states: chlorine is a pale yellow-green gas, bromine is a red-brown liquid, and iodine is a grey-black solid. This variation in physical state is linked to their molecular size and intermolecular forces, which increase down the group.

• Halogens are diatomic molecules (Cl₂, Br₂, I₂).
• They are non-metal elements in Group VII.
• Physical states at room temperature vary: chlorine (gas), bromine (liquid), iodine (solid).

Density is defined as mass per unit volume (g/cm³). As you move down Group VII from chlorine to iodine, the density of the halogens increases. This is because the atoms become larger and heavier due to the addition of electron shells, increasing the relative atomic mass.

For example, chlorine gas has a density of about 3.2 g/L (at standard conditions), bromine liquid has a density around 3.1 g/cm³, and solid iodine has a density approximately 4.9 g/cm³. The increase in density reflects the increasing mass and stronger intermolecular forces that hold the molecules closer together in the liquid and solid states.

This trend is consistent with the general periodic trend where heavier elements have higher densities.

• Density increases down Group VII due to increasing atomic mass.
• Chlorine is less dense than bromine, which is less dense than iodine.
• Physical state changes (gas to liquid to solid) accompany density increase.

Reactivity refers to how readily an element undergoes chemical reactions. In Group VII, reactivity decreases as you move down from chlorine to iodine. This trend is linked to the halogens' ability to gain an electron to form a halide ion (X⁻).

Chlorine is the most reactive halogen because its atoms have a smaller atomic radius and stronger attraction for electrons, making it easier to gain an electron. As the atomic radius increases down the group, the outer electrons are further from the nucleus and shielded by inner shells, reducing the effective nuclear attraction. This makes it harder for bromine and iodine to attract electrons, thus lowering their reactivity.

For example, chlorine can displace bromide and iodide ions from solutions, while bromine can displace iodide but not chloride, demonstrating their relative reactivities.

• Reactivity decreases down Group VII.
• Smaller atoms attract electrons more strongly, increasing reactivity.
• Chlorine > bromine > iodine in reactivity order.

The trends in density and reactivity can be explained by changes in atomic structure down the group. Each successive element has an additional electron shell, increasing atomic size and mass.

The increased number of electrons and shells leads to greater van der Waals forces between molecules, raising melting and boiling points and density.

Conversely, the increased distance between the nucleus and outer electrons reduces the nucleus's pull on incoming electrons, lowering reactivity. This balance between atomic size, electron shielding, and nuclear charge governs the observed trends.

• Additional electron shells increase atomic size and mass.
• Stronger intermolecular forces increase density and physical state changes.
• Electron shielding reduces nuclear attraction, decreasing reactivity.

To summarise, Group VII halogens are diatomic non-metals with increasing density and decreasing reactivity down the group. Chlorine is a pale yellow-green gas, bromine a red-brown liquid, and iodine a grey-black solid at room temperature.

These trends are important for predicting chemical behaviour and practical applications, such as the use of chlorine as a disinfectant due to its high reactivity, and iodine in antiseptics where lower reactivity is advantageous.

• Halogens are diatomic non-metals.
• Density increases down the group: Cl < Br < I.
• Reactivity decreases down the group: Cl > Br > I.
• Physical states change from gas to liquid to solid down the group.

Halogens are elements found in Group VII (Group 17) of the periodic table. At room temperature and pressure (approximately 20–25°C and 1 atmosphere), the halogens chlorine, bromine, and iodine exist in different physical states: gas, liquid, and solid respectively. This variation is unusual within a single group and is linked to their molecular properties and intermolecular forces.

Understanding these physical appearances is important for recognising halogens in laboratory and environmental settings, as well as for predicting their behaviour in chemical reactions.

• Halogens are diatomic molecules (e.g., Cl₂, Br₂, I₂).
• Physical state varies down the group: chlorine (gas), bromine (liquid), iodine (solid).

Chlorine exists as Cl₂ molecules. At room temperature, chlorine is a gas with a distinctive pale yellow-green colour. This gaseous state is due to relatively weak intermolecular forces between chlorine molecules, specifically London dispersion forces, which are insufficient to hold the molecules in a liquid or solid state at room temperature.

Chlorine gas is denser than air and has a sharp, irritating smell. Its gaseous form allows it to diffuse quickly, which is important in its use as a disinfectant in water treatment.

Example: When a small amount of chlorine gas is released in a laboratory fume cupboard, it appears as a faint yellow-green cloud.

• State: Gas at room temperature and pressure.
• Colour: Pale yellow-green.
• Reason for state: Weak intermolecular forces between Cl₂ molecules.

Bromine (Br₂) is unique among the halogens because it is a liquid at room temperature. It has a deep red-brown colour and a strong, unpleasant odour. The liquid state arises because bromine molecules are larger than chlorine molecules, leading to stronger London dispersion forces that hold the molecules closer together in the liquid phase.

Bromine's liquid form is volatile, meaning it evaporates easily to form a coloured vapour. This volatility is a practical consideration when handling bromine, requiring careful containment to avoid inhalation of toxic vapours.

Example: A drop of bromine placed on a watch glass appears as a dark red-brown liquid that slowly evaporates, releasing coloured vapour.

• State: Liquid at room temperature and pressure.
• Colour: Red-brown.
• Reason for state: Larger molecules increase intermolecular forces, enabling liquid state.

Iodine (I₂) is a solid at room temperature, appearing as shiny grey-black crystals. Its molecular size is larger than bromine and chlorine, resulting in even stronger London dispersion forces. These forces hold iodine molecules tightly in a solid lattice.

When heated gently, iodine sublimates, turning directly from solid to purple vapour without becoming liquid at room temperature. This sublimation is a distinctive property used to identify iodine.

Example: Solid iodine crystals placed in a warm environment release a purple vapour, demonstrating sublimation.

• State: Solid at room temperature and pressure.
• Colour: Grey-black crystalline appearance.
• Reason for state: Strong intermolecular forces due to large molecular size.

The physical state of halogens at room temperature is determined by the strength of intermolecular forces between their diatomic molecules. As molecular size increases down the group (Cl₂ < Br₂ < I₂), the number of electrons increases, enhancing London dispersion forces.

These stronger forces require more energy (higher temperature) to overcome, so the melting and boiling points increase down the group. Consequently, chlorine is a gas, bromine a liquid, and iodine a solid at room temperature.

This trend illustrates how molecular structure influences macroscopic properties such as state and colour, which are critical for identification and practical use.

• Increasing molecular size increases London dispersion forces.
• Stronger forces raise melting and boiling points.
• Physical state changes from gas to liquid to solid down the group.

Halogens are elements in Group VII of the periodic table, including fluorine, chlorine, bromine, and iodine. They commonly exist as diatomic molecules (e.g., Cl₂, Br₂). Halide ions are negatively charged ions formed when halogens gain an electron, such as chloride (Cl⁻), bromide (Br⁻), and iodide (I⁻).

A displacement reaction occurs when a more reactive halogen atom displaces a less reactive halide ion from a compound. This type of reaction demonstrates the relative reactivity of halogens and is a key chemical property of Group VII elements.

• Halogens exist as diatomic molecules (X₂).
• Halide ions are negatively charged ions (X⁻).
• Displacement reactions involve a more reactive halogen replacing a less reactive halide ion.

The reactivity of halogens decreases down the group: fluorine is the most reactive, followed by chlorine, bromine, and iodine. This trend is due to the increasing atomic size and shielding effect, which make it harder for larger atoms to attract electrons.

Because of this trend, a halogen higher in the group can displace the halide ion of any halogen lower in the group from its compounds. For example, chlorine can displace bromide and iodide ions, but iodine cannot displace chloride ions.

This reactivity order is crucial for predicting the outcomes of displacement reactions involving halogens.

• Reactivity order: F₂ > Cl₂ > Br₂ > I₂.
• Higher reactivity means stronger oxidising ability.
• More reactive halogens can displace less reactive halide ions.

In a displacement reaction, the more reactive halogen molecule (X₂) gains electrons from the halide ions (Y⁻) in solution, oxidising them to halogen atoms (Y₂) and forming halide ions (X⁻) itself.

For example, when chlorine gas is bubbled through sodium bromide solution, chlorine molecules accept electrons from bromide ions, producing bromine and chloride ions.

This electron transfer is a redox process: the halogen molecule is reduced, and the halide ion is oxidised.

• More reactive halogen (X₂) oxidises halide ion (Y⁻).
• Electron transfer converts Y⁻ to Y₂ and X₂ to X⁻.
• Reaction is a redox process involving electron gain and loss.

A common example is chlorine displacing bromide ions from sodium bromide solution:

Cl₂ (aq) + 2Br⁻ (aq) → 2Cl⁻ (aq) + Br₂ (aq) (chlorine displaces bromine)

Similarly, chlorine can displace iodide ions: Cl₂ + 2I⁻ → 2Cl⁻ + I₂, and bromine can displace iodide ions: Br₂ + 2I⁻ → 2Br⁻ + I₂. However, iodine cannot displace chloride or bromide ions due to its lower reactivity, so no reaction occurs in those cases.

• Cl₂ + 2NaBr → 2NaCl + Br₂
• Cl₂ + 2NaI → 2NaCl + I₂
• Br₂ + 2NaI → 2NaBr + I₂
• I₂ does not displace Cl⁻ or Br⁻ ions.

Displacement reactions can be observed by changes in colour. For example, bromine solution is orange-brown, iodine solution is brown, and chlorine solution is pale green.

When chlorine displaces bromine from bromide ions, the solution changes from colourless to orange-brown due to bromine formation.

Safety precautions are important when handling halogens, as they are toxic and corrosive gases or liquids. Experiments should be conducted in a fume cupboard with appropriate protective equipment.

• Colour changes indicate displacement: e.g., colourless to orange-brown for bromine.
• Use fume cupboard and gloves for safety.
• Halogen solutions have characteristic colours useful for identification.

Group VII elements, also called halogens, include fluorine, chlorine, bromine, iodine, and astatine. These elements are non-metals with seven electrons in their outermost shell, making them highly reactive as they tend to gain one electron to achieve a full outer shell.

The halogens exist in different physical states at room temperature: fluorine and chlorine are gases, bromine is a liquid, and iodine and astatine are solids. This variation is part of a predictable trend down the group.

• Halogens have seven valence electrons.
• They are highly reactive non-metals.
• Physical states vary from gases to solids down the group.

As you move down Group VII from fluorine to astatine, the atomic number increases, adding more electron shells. This increase in electron shells causes the atomic radius to increase because the outer electrons are further from the nucleus.

The increased number of protons also increases nuclear charge, but the effect of additional inner shells (shielding) reduces the attraction felt by outer electrons, allowing atomic size to grow.

Consequently, the mass of the atoms increases down the group due to the addition of protons and neutrons in the nucleus.

• Atomic radius increases down the group due to added electron shells.
• Shielding reduces effective nuclear charge on outer electrons.
• Atomic mass increases with proton and neutron number.

Melting and boiling points of halogens increase down the group. This is because larger atoms have stronger van der Waals forces between molecules, requiring more energy to change state.

For example, fluorine and chlorine are gases at room temperature due to weak intermolecular forces, bromine is a liquid with stronger forces, and iodine is a solid with even stronger forces.

Density also increases down the group as atoms become heavier and more closely packed in the solid or liquid state.

• Melting and boiling points rise down the group.
• Intermolecular forces strengthen with larger atomic size.
• Physical state changes from gas to solid down the group.
• Density increases due to heavier atoms.

Reactivity of halogens decreases down the group. Fluorine is the most reactive, followed by chlorine, bromine, iodine, and astatine.

This trend occurs because the atomic radius increases, making it harder for the nucleus to attract an extra electron due to increased distance and shielding.

Thus, fluorine more readily gains an electron to form halide ions, while iodine and astatine are less reactive and less likely to gain electrons quickly.

• Reactivity decreases down the group.
• Smaller atoms attract electrons more strongly.
• Electron affinity decreases down the group.
• Halogens form halide ions by gaining one electron.

Using these trends, we can predict properties of halogens not explicitly studied. For example, astatine, below iodine, is expected to be a solid at room temperature with a higher melting point than iodine.

Its reactivity will be lower than iodine's, and it will have a larger atomic radius and higher density.

Similarly, if a new halogen were discovered below astatine, it would likely be even less reactive, with higher melting and boiling points, and exist as a solid under normal conditions.

• Astatine predicted to be a solid with high melting point.
• Reactivity continues to decrease down the group.
• Physical and chemical properties follow established trends.
• Predictions rely on atomic structure and periodic trends.