<P> Chlorine is the second halogen, being a nonmetal in group 17 of the periodic table . Its properties are thus similar to fluorine, bromine, and iodine, and are largely intermediate between those of the first two . Chlorine has the electron configuration (Ne) 3s 3p, with the seven electrons in the third and outermost shell acting as its valence electrons . Like all halogens, it is thus one electron short of a full octet, and is hence a strong oxidising agent, reacting with many elements in order to complete its outer shell . Corresponding to periodic trends, it is intermediate in electronegativity between fluorine and bromine (F: 3.98, Cl: 3.16, Br: 2.96, I: 2.66), and is less reactive than fluorine and more reactive than bromine . It is also a weaker oxidising agent than fluorine, but a stronger one than bromine . Conversely, the chloride ion is a weaker reducing agent than bromide, but a stronger one than fluoride . It is intermediate in atomic radius between fluorine and bromine, and this leads to many of its atomic properties similarly continuing the trend from iodine to bromine upward, such as first ionisation energy, electron affinity, enthalpy of dissociation of the X molecule (X = Cl, Br, I), ionic radius, and X--X bond length . (Fluorine is anomalous due to its small size .) </P> <P> All four stable halogens experience intermolecular van der Waals forces of attraction, and their strength increases together with the number of electrons among all homonuclear diatomic halogen molecules . Thus, the melting and boiling points of chlorine are intermediate between those of fluorine and bromine: chlorine melts at − 101.0 ° C and boils at − 34.0 ° C. As a result of the increasing molecular weight of the halogens down the group, the density and heats of fusion and vaporisation of chlorine are again intermediate between those of bromine and fluorine, although all their heats of vaporisation are fairly low (leading to high volatility) thanks to their diatomic molecular structure . The halogens darken in colour as the group is descended: thus, while fluorine is a pale yellow gas, chlorine is distinctly yellow - green . This trend occurs because the wavelengths of visible light absorbed by the halogens increase down the group . Specifically, the colour of a halogen, such as chlorine, results from the electron transition between the highest occupied antibonding π molecular orbital and the lowest vacant antibonding σ molecular orbital . The colour fades at low temperatures, so that solid chlorine at − 195 ° C is almost colourless . </P> <P> Like solid bromine and iodine, solid chlorine crystallises in the orthorhombic crystal system, in a layered lattice of Cl molecules . The Cl--Cl distance is 198 pm (close to the gaseous Cl--Cl distance of 199 pm) and the Cl Cl distance between molecules is 332 pm within a layer and 382 pm between layers (compare the van der Waals radius of chlorine, 180 pm). This structure means that chlorine is a very poor conductor of electricity, and indeed its conductivity is so low as to be practically unmeasurable . </P> <P> Chlorine has two stable isotopes, Cl and Cl . These are its only two natural isotopes occurring in quantity, with Cl making up 76% of natural chlorine and Cl making up the remaining 24% . Both are synthesised in stars in the oxygen - burning and silicon - burning processes . Both have nuclear spin 3 / 2 + and thus may be used for nuclear magnetic resonance, although the spin magnitude being greater than 1 / 2 results in non-spherical nuclear charge distribution and thus resonance broadening as a result of a nonzero nuclear quadrupole moment and resultant quadrupolar relaxation . The other chlorine isotopes are all radioactive, with half - lives too short to occur in nature primordially . Of these, the most commonly used in the laboratory are Cl (t = 3.0 × 10 y) and Cl (t = 37.2 min), which may be produced from the neutron activation of natural chlorine . </P>

What kind of bond is formed in pure clorine