Group 7

• Reduction


  The process of electron gain (Decrease in oxidation number) (loss of oxygen OR gain of hydrogen)
• Group 7 ions are larger than the neutral atoms because electrons are gained to form the negative ion, leading to repulsion between the electrons moving them further apart from each other. Also, the effective nuclear charge decreases as there are more electrons with the same number of protons
• Group 7 elements are called the halogens
• The halogens are in the p block of the periodic table because their highest energy electron is in a p orbital
• Redox reaction

  A reaction in which reduction and oxidation occurs
• The halogens react with metals to form ionic compounds and form ions with a single negative charge called halide ions (F-, Cl-, Br-, I-)
• Reducing agents

  Electron donors
• Oxidation


  The process of electron loss (Increase in oxidation number) (gain of oxygen OR loss of hydrogen)
• Atomic radius increases down the Group 7
• Disproportionation

  When an element is both reduced and oxidised in the same reaction
• Oxidising agents

  Electron acceptors
• Electronegativity


  The power of an atom to attract the pair of electrons in a covalent bond
• As you go down the group, the number of shells increases, leading to a greater distance between the nucleus and outer electrons, increased shielding, weaker attraction between nucleus and outer electrons, and therefore an increase in atomic radius
• The halogens are diatomic (e.g., F2, Cl2, Br2, I2)
• Fluorine is the most electronegative element on the periodic table
• First ionisation energy


  Decreases down the group as atomic radius increases and shielding increases, resulting in weaker attraction between nucleus and outer electrons, requiring less energy to remove the outer electron
• If you shake an aqueous solution of a halogen with an organic solvent, an aqueous layer and an organic layer will form as the two solvents are immiscible. The aqueous layer originally containing the halogen will decolourise as the halogen moves into the organic layer, where the colour of the halogen will then be seen due to increased solubility in the organic layer
• Appearance and solubility of the halogens
• Group 7 ions

  Are larger than the neutral atoms because electrons are gained to form the negative ion; the repulsion between the electrons moves them further apart from each other. Also, the effective nuclear charge decreases as there are more electrons with the same number of protons
• State and appearance of halogens

  • F: Yellow gas
  • Cl: Green gas
  • Br: Red/brown liquid, with red/brown vapour
  • I: Shiny grey/black solid, sublimes to violet vapour
• Definition: First ionisation energy - The energy required to remove one mole of electrons from one mole of atoms in the gaseous state to form 1 mole of unipositive gaseous ions
• Melting point


  Increases down the group as halogen molecules have a bigger Mr and more electrons, leading to stronger van der Waals forces between the molecules, requiring more energy to overcome these intermolecular forces, resulting in a higher melting point
• The halogens are sparingly soluble in water as they are non-polar and become less soluble as you go down the group
• Electronegativity


  Decreases down the group due to an increase in atomic radius and more shielding, resulting in weaker electrostatic attraction between the nucleus and the shared pair of electrons in a covalent bond
• Reactivity


  Decreases down the group as atomic radius increases and shielding increases, leading to weaker attraction between nucleus and outer electrons, making it harder to form negative ions in a chemical reaction
• Definition: Electronegativity - The power of an atom to attract a shared pair of electrons in a covalent bond
• The halogens are soluble in non-polar solvents such as cyclohexane
• Halide ions have eight electrons in their outer shell. When they form halogen atoms, they donate/lose an electron, making them reducing agents
• The oxidising power of the halogens decreases down the group due to larger atomic radius and more shielding, making it more difficult to accept electrons
• Displacement reactions

  Reactions are determined by the oxidising ability of the halogen and the reducing ability of the halide ion
• Displacement reactions

  Chlorine (Cl2) is a good oxidising agent, so it will oxidise bromide (Br-) and iodide (I-) ions
• Displacement reactions

  Iodine (I2) is a poor oxidising agent, so it will not oxidise chloride (Cl-) or bromide (Br-) ions
• Halogens have seven electrons in their outer shell. When they form halide ions, they accept/gain an electron, making them oxidising agents
• Displacement reactions

  Cl- ions are poor reducing agents, so they will not reduce iodine or bromine
• The reducing power of the halide ions increases down the group due to larger ionic radius and more shielding, making it easier to donate electrons
• Displacement reactions


  1. ions are good reducing agents, so they will reduce chlorine and bromine
• When chlorine water reacts with aqueous potassium bromide, an orange solution appears
• Displacement reactions

  Bromine (Br2) will not oxidise chloride (Cl-) ions, but will oxidise iodide (I-) ions
• Halogen states

  • F: Yellow gas
  • Cl: Green gas
  • Br: Red/brown liquid, with red/brown vapour
  • I: Shiny grey/black solid, sublimes to violet vapour
• The strongest reducing agent is I-