Properties, reactions and trends of Group 7 (halogens)
Atomic structure and the periodic table • The periodic table
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Outer shell electrons determine reactivity
Each halogen atom has seven electrons in its outermost shell and requires one more electron to achieve a full (noble gas) configuration. Because of this electronic configuration, halogen atoms attract an extra electron from other atoms; this attraction causes high electronegativity and oxidising behaviour. Down the group, atomic radius increases and inner electron shielding increases, so the nucleus attracts an extra electron less strongly and reactivity decreases. Cause → effect: seven outer electrons → strong tendency to gain one electron → formation of halide ions (X−) or covalent bonds; increased atomic size and shielding down the group → weaker attraction for an additional electron → lower reactivity for bromine and iodine compared with chlorine.
Physical properties and trends down the group
All halogens form diatomic molecules (F2, Cl2, Br2, I2). Molecular size and mass increase down the group, which increases London (van der Waals) forces between molecules. Cause → effect: larger, more polarisable molecules → stronger intermolecular forces → higher melting and boiling points down the group. Chlorine is a gas at room temperature, bromine is a red-brown liquid, and iodine is a purple-black solid that sublimes. Other trends: electronegativity decreases down the group because increased distance and shielding reduce nuclear attraction for bonding electrons. Oxidising strength also decreases down the group because the ability to gain electrons becomes weaker.
Reactions of halogens with metals - ionic halides
Halogens react with most metals by gaining electrons from metal atoms to form halide ions (Cl−, Br−, I−) and producing ionic halide salts. Cause → effect: metal atoms lose electrons to form positive ions; halogen atoms gain electrons to form negative halide ions; electrostatic attraction between ions produces ionic lattices. Examples include sodium chloride (NaCl) and magnesium bromide (MgBr2). Ionic halide salts often form white crystalline solids and many dissolve in water to give aqueous halide ions; solubility depends on the metal and lattice energy. Limiting factors: metal reactivity and ionic charge influence lattice strength and solubility; transition-metal halides may show different colours and variable solubility due to covalent character or hydration.
Reactions of halogens with non-metals - covalent compounds
Halogens form covalent bonds with non-metals by sharing electrons. Cause → effect: two non-metal atoms share pairs of electrons to achieve complete outer shells. Examples include hydrogen halides (HCl, HBr, HI), which exist as diatomic molecules and dissolve in water to form strong acids (hydrochloric acid etc.). Halogens also form covalent bonds with carbon in organic halides; bond polarity depends on the electronegativity difference and may influence physical properties and reactivity. Limiting factors: bond polarity and molecular polarity influence solubility and reactivity; heavier halogens form weaker, more polarisable bonds and often show different reactivity patterns in organic reactions.
Displacement reactions and oxidising ability
A more reactive halogen displaces a less reactive halide ion from solution. Cause → effect: stronger oxidising halogen molecules oxidise halide ions to form X− → for example, chlorine displaces bromide or iodide from their salts, while iodine cannot displace bromide or chloride. The order of oxidising strength and reactivity is: fluorine > chlorine > bromine > iodine. Practical observation: adding chlorine water to a solution of potassium bromide yields brown bromine; adding bromine water to a potassium iodide solution yields purple iodine.
Key notes
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