3.1

• What is the atomic number range of elements in the periodic table?


  From 1 to 118
• How are elements arranged in the periodic table?


  By increasing atomic number and similar electronic configurations
• What is a period in the context of the periodic table?


  A row of elements arranged by increasing atomic number
• What do the group numbers in the periodic table indicate?


  The number of valence electrons in the outer shell
• What are valence electrons?


  Electrons in the outermost shell of an atom
• Why are hydrogen and helium placed in their respective groups?


  Based on similarities in physical and chemical properties
• What are the four main blocks of the periodic table?

  • s-block
  • p-block
  • d-block
  • f-block
• What characterizes s-block elements?


  They have only s electrons in the outer shell
• What characterizes p-block elements?


  They have at least one p-electron in the outer shell
• What characterizes d-block elements?


  They have at least one d-electron and one s-electron in the outer shell
• What characterizes f-block elements?


  They have at least one f-electron and one s-electron in the outer shell
• What are periodic trends?


  • Patterns in physical and chemical properties of elements
  • Trends repeat in each period of the periodic table
  • The study of these trends is known as periodicity
• How can the electron configuration of an element be deduced?


  From its position in the periodic table
• What is the electron configuration of germanium?


  It has four valence electrons in the fourth shell
• What does the position of an element in the periodic table indicate about its electron configuration?


  It indicates the number of valence electrons and the shell they occupy
• What is the first ionisation energy (IE)?


  The energy required to remove one mole of electrons from one mole of gaseous atoms
• Under what conditions are ionisation energies measured?


  At standard conditions of 298 K and 100 kPa
• What are the units of ionisation energy?


  Kilojoules per mole (kJ mol^-1)
• How does ionisation energy change across a period?


  It increases across a period
• How does ionisation energy change down a group?


  It decreases down a group
• What factors influence ionisation energy?


  • Size of the nuclear charge
  • Distance of outer electrons from the nucleus
  • Shielding effect of inner electrons
  • Spin-pair repulsion
• Why does ionisation energy increase across a period?


  Due to increased nuclear charge and constant distance to outer electrons
• What causes the rapid decrease in ionisation energy between periods?


  Increased distance and shielding outweigh the increased nuclear charge
• What is the first ionisation energy of beryllium?


  900 kJ mol^-1
• What is the first ionisation energy of boron?


  801 kJ mol^-1
• What is the first ionisation energy of nitrogen?


  1402 kJ mol^-1
• What is the first ionisation energy of oxygen?


  1314 kJ mol^-1
• What are the trends in ionisation energy down a group?


  • Ionisation energy decreases
  • Due to increased distance from nucleus
  • Increased shielding by inner shell electrons
• What are the trends in ionisation energy across a period?


  • Ionisation energy increases
  • Due to increased nuclear charge
  • Distance and shielding remain constant
• What is the electron configuration of beryllium?


  1s 2s
• Why does boron have a lower first ionisation energy than beryllium?


  Because the 2p subshell in boron is further from the nucleus than the 2s subshell in beryllium
• What causes the slight decrease in first ionisation energy between nitrogen and oxygen?


  Spin-pair repulsion in the 2p subshell of oxygen
• What are the factors affecting ionisation energy down a group?


  • Distance between the nucleus and outer electron increases
  • Shielding by inner shell electrons increases
  • Effective nuclear charge decreases as shielding increases
• What happens to ionisation energy as you go down a group?


  It decreases
• What is the trend in successive ionisation energies of an element?


  • Successive ionisation energies increase
  • Removing an electron from a positive ion is more difficult
  • Attractive forces increase due to decreasing shielding and increased proton to electron ratio
• Why do successive ionisation energies become very large and difficult to represent meaningfully?


  Because removing electrons from increasingly positive ions requires more energy
• What is the first electron removed from calcium characterized by?


  It has a low ionisation energy due to spin-pair repulsion in the 4s orbital
• How does the removal of the second electron from calcium compare to the first?


  The second electron is more difficult to remove than the first
• What does a large increase in successive ionisation energy indicate?


  It indicates a change of shell
• What is electron affinity (EA)?


  • Energy released when one mole of electrons is gained by one mole of atoms
  • Measured under standard conditions (298 K and 100 kPa)
  • Units are kilojoules per mole (kJ mol^-1)