CHEM-S3.1

• Periodicity


  Repeated pattern of properties in every period
• Period


  Each row of periodic table
• Group


  Each column of periodic table
• General trends of atomic radius

  1. Down the group: Outer electrons are in higher energy level which further away from nucleus, Atomic radius increase

  2. Across the period: Number of protons increases, nuclear charge increases, Electrons are added to same energy level, Shielding of electrons is about the same, Nuclear attraction to outer electron increase, Atomic radius decreases
• Cation


  Smaller size than the atom
• Anion


  Larger size than the atom
• Cation and anion examples

  • Na+ ion is smaller than Na atom
  • O2- ion is larger than O atom
  • Mg2+ ion is smaller than Cl- ion
  • Cl- ion is smaller than S2- ion
  • C4+ ion is smaller than C4- ion
• Electronegativity


  Ability of atom to attract the bond pair electrons
• Noble gases are not assigned electronegativity values because they are unreactive and do not form bonds
• Trend in electronegativity


  1. Across the period: Number of protons increase, nuclear charge increase, atomic radius decreases, nuclear attraction to electrons in bond becomes stronger, Electronegativity increase

  2. Down the group: Atomic radius increase, outer electrons are in higher energy level, nuclear attraction to electrons in bond becomes weaker, electronegativity decrease
• Relationship of electronegativity difference of compound and bond character

  • Higher difference in electronegativity: Higher ionic character
  • Lower difference in electronegativity: Higher covalent character

  
• First ionization energy

  Minimum energy required to remove one mole electrons from one mole gaseous atoms
• Second ionization energy


  Minimum energy required to remove one mole electrons from one mole gaseous singly-charged positive ions of an element
• Successive ionization energy


  Energy required to remove each electron
• Why the successive ionization energy of neon increase: First electron is removing electron from Ne atom, the other electron is removing electron from more positive ions, Repulsion between electrons decrease, Size of the ions formed decrease as electrons are removed, the attraction from the positive ions to the remaining electrons is stronger
• The electronic configuration of neon is 2,8 based on the large energy gap between 8th and 9th ionization energy
• Trend of ionization energy down the group

  Atomic radius increases, there is more shielding of outermost electrons, nuclear attraction to the outermost electrons decreases, ionization energy decrease
• Trend of ionization energy across the period

  Number of protons increase, nuclear charge increases, outermost shell electrons are in same electron shell/energy level, shielding of outermost electrons is about the same, atomic radius decreases, nuclear attraction to the outermost electrons increases, ionization energy increase
• Exceptions to ionization energy trends: Ionization energy of boron is lower than beryllium, ionization energy of aluminium is lower than magnesium, ionization energy of oxygen is lower than nitrogen, ionization energy of sulfur is lower than phosphorus
• First electron affinity


  Energy released when one mole of gaseous atoms acquires one mole electron
• Trend of first electron affinity down the group: Atomic radius of atom increases, added electron is further away from the nucleus, more shielding of electrons, nuclear attraction to incoming electrons decreases, electron affinity become less exothermic (negative)
• Second electron affinity


  Energy required to add one mole electron to one mole of gaseous 1- ions
• Second electron affinity is endothermic because there is repulsion between the negative ions and the extra electrons
• Group 1: The alkali metals


  • Soft metals with low density
  • React vigorously with water/acids to form hydrogen
  • React with in water and form alkaline solution
  • Form basic oxides
  • Reactivity increases down the group
  • Boiling point/melting point decrease down the group
• Group 17: The halogens


  • Diatomic molecules
  • Form negative ion, X-
  • Form acidic oxides
  • Reactivity decreases down the group
  • Boiling/melting point increase down the group
  • Perform displacement reaction
• The most vigorous reaction occurs between the elements that are the most reactive, e.g. Fluorine + francium (caesium)
• Displacement reaction: When chlorine gas is bubbled through a solution of sodium iodide, the chlorine displaces the iodine from the solution
• Gas bubbles
• Solid dissolve/disappear
• Form basic oxides
• Reactivity increases down the group (higher reaction rate and more vigorously)
• Caesium is more reactive than rubidium
• Caesium has a lower ionization energy
• Caesium is easier to remove electron
• Boiling point/melting point decrease down the group (across period)
• Halogens


  Group 17 elements
• Halogens


  • Diatomic molecules(F2, Cl2, I2, Br2)
  • Form negative ion, X-
  • Form acidic oxides
  • Reactivity decreases down the group(reverse trend of Group 1 and 2)
  • Boiling/melting point increase down the group(reverse trend of Group 1 and 2)
  • Perform displacement reaction
• Reaction with group 1 metals

  1. Ionic halides are formed

  2. 2Na(s) + Cl2(g) -> 2NaCl(s)

  3. The most vigorous reaction occurs between the elements that are the most reactive
• Displacement reaction


  1. When chlorine is added to aqueous potassium bromide

  2. Colour of the solution change from colourless to orange/brown(bromine)

  3. Chlorine is more reactive than bromine and able to displace bromine

  4. Reaction equation: Cl2 + 2KBr -> 2KCl + Br2

  5. Ionic equation: Cl2 + 2Br- -> 2Cl- + Br2
• Halogen

  • Potassium chloride
  • Potassium bromide
  • Potassium iodide
  • Chlorine
  • Bromine
  • Iodine