Characteristic Properties of the Halogens

1. Characteristic Properties of the Halogens

2. halogens Introduction • Group VIIA elements include • fluorine •  chlorine •  bromine •  iodine •  astatine (Salt producers)

3. Introduction • Astatine •  chemistry not much known • radioactive •  the total amount present in the Earth's crust is probably less than 30 g at any one time.

4. Halogens are p-block elements  outermost shell electronic configuration of ns2np5

5. Halogens are p-block elements one electron short of the octet structure

6. Introduction • In the free elemental state they complete their octets by sharing theirsingle unpaired p-electrons

7. When halogens react with other elements they either gainan additional electron to form halide ions or sharetheir single unpaired p-electrons to form single covalent bonds

8. High Electronegativity / Electron Affinity highest among the elements in the same period • have a high tendency to attract electrons • strong oxidizing agents

9. High Electronegativity / Electron Affinity • -1 is the most common oxidation state of halogens in their compounds Ionic : NaF, NaCl, NaBr, NaI Covalent : HF, HCl, HBr, HI

10. Variable Oxidation State All halogens (except fluorine) can expand their octet of electrons by utilizing the vacant, energetically low-lying d-orbitals.

11. “Electrons-in-boxes” diagrams of the electronic configuration of a halogen atom of the ground state and various excited states

12. +3 +1 The half-filled orbital(s) overlap(s) with those of more electronegative atoms (e.g. O)  positive oxidation state (+1, +3, +5, +7) +7 +5

13. Various oxidation states of halogens in their ions or compounds

14. Various oxidation states of halogens in their ions or compounds

15. Fluorine (1) • the most electronegative element • only one unpaired p electron available for bonding • oxidation state is limited to –1

16. Fluorine (1) • cannot expand its octet •  no low-lying empty d orbitals available •  the energy required to promote electrons into the third quantum shell is very high • Absence of HFO, HFO2, HFO3, HFO4

17. Variation in Physical Properties 1. Melting point / boiling point  down the group

18. Variations in melting point and boiling point of the halogens

19. Variation in Physical Properties 1. Melting point / boiling point  down the group • The molecular size  down the group • The electron cloud is more easily polarized • Induced dipoles are formed more easily • Stronger London dispersion forces

20. 2. Colour becomes darker down the group

21. chlorine Appearances of halogens at room temperature and pressure: chlorine

22. bromine Appearances of halogens at room temperature and pressure: bromine

23. iodine Appearances of halogens at room temperature and pressure: iodine

24. Colour • All halogens • coloured •  the absorption of radiationin the visible light region of the electromagnetic spectrum • The colour is due to the unabsorbed radiation in the visible light region

25. Colour • Fluorine atom • has thesmallest size • absorbs the radiation of relativelyhigh frequency (i.e. blue light) • appears yellow(the unabsorbed radiation)

26. Colour • Atoms of other halogens • larger sizes • absorb radiation of lower frequency

27. Colour • Iodine • absorbs the radiation of relativelylow frequency (i.e. yellow light) • appears violet

28. Q.1 The colour of astatine is black.

29. Colour • Halogens • different colours when dissolved in different solvents

30. Colours of halogens in pure form and in solutions

31. Colour • Halogens •  non-polar molecules not very soluble in polar solvents (such as water) but very soluble in organic solvents (such as 1,1,1-trichloroethane)

32. (a) (b) (c) Colours of halogens in water:(a) chlorine; (b) bromine; (c) iodine

33. (a) (b) (c) Colours of halogens in 1,1,1-trichloroethane:(a) chlorine; (b) bromine; (c) iodine

34. 3. Electron Affinity  down the group

35. The number of electron shells and size of atoms  down the group  The nuclear attraction for the additional electron  down the group  Electron affinity  from Cl to I

36. Atoms of fluorine have the smallest size among the halogens  The addition of an extra electron to the small quantum shell(n=2) results in great repulsion among the electrons.  Fluorine has a lower electron affinity than Cl and Br.

37. 4. Electronegativity  down the group

38. The number of electron shells and size of atoms  down the group  The nuclear attraction for the bonding electrons  down the group  Electronegativity  down the group

39. Fluorine has the highest electronegativity because it is the most reactive elements. The electronegativity of fluorine is arbitrarily assigned as 4.0.

40. Variation in Chemical Properties Reactivity : F2 > Cl2 > Br2 > I2 React by gaining electrons Oxidizing power : F2 > Cl2 > Br2 > I2

41. 1. Reactions with Sodium • All halogens • combine directly with sodium to form sodium halides • the reactivity decreases down the group from fluorinetoiodine

42. 1. Reactions with Sodium • Fluorine • react explosively to form sodium fluoride • 2Na(s) + F2(g)  2NaF(s)

43. 1. Reactions with Sodium • Chlorine • reacts violentlyto form sodium chloride • 2Na(s) + Cl2(g)  2NaCl(s)

44. 1. Reactions with Sodium • Bromine • burns steadily in bromine vapour to form sodium bromide • 2Na(s) + Br2(g)  2NaBr(s)

45. 1. Reactions with Sodium • Iodine •  burns steadily in iodine vapour to form sodium iodide • 2Na(s) + I2(g)  2NaI(s)

46. Na+(g) + X(g) E.A. B.E. Na+(g) + X(g) Na+(g) + X2(g) I.E. Na(s) + X2 NaX(s) • Vigor of reaction depends on • The activation energy (endothermic) • The lattice energy (exothermic)  Activation energy

47. Na+(g) + X(g) E.A. B.E. Na+(g) + X(g) Na+(g) + X2(g) F has an exceptionally low B.E. & zero I.E. Na(s) + X2 NaX(s) F is the most reactive (g)

48. Na+(g) + X(g) E.A. B.E. Na+(g) + X(g) Na+(g) + X2(g) The lattice enthalpy of NaF is most negative I.E. Na(s) + X2 NaX(s)

49. Na+(g) + X(g) E.A. B.E. Na+(g) + X(g) Na+(g) + X2(g) Cl has zero I.E. Na(s) + X2 NaX(s) Cl is more reactive than Br & I (g)

50. Na+(g) + X(g) E.A. B.E. Na+(g) + X(g) Na+(g) + X2(g) I.E. Na(s) + X2 NaX(s) Lattice enthalpy : NaCl > NaBr > NaI