periodicity: content Flashcards
(11 cards)
how does atomic radius change going down a group
INCREASES due to more principal energy levels
atomic radius going across a period
DECREASES
- more protons = greater attraction to electrons
equation for ionisation energy
[ ]ⁿ⁺ (g) →[ ]⁽ⁿ⁺¹⁾⁺ (g) + e⁻
trends in successive ionisation energy
successive ionisation energies INCREASE
- stronger nuclear charge and smaller ionic radius = increased attraction on the outer electron (the one getting removed)
what governs atomic radius
the attraction between nucleus and electrons
- shielding
- nuclear charge
- no. of electron shells
trends in first ionisation energy going down a group
DECREASES: less energy is needed
- electron removed from a higher energy level (further from nucleus)
- more shielding
- weaker attraction between nucleus and outermost electron
trends in first ionisation energy going across a period
INCREASES; more energy is needed
- atomic radius decreases (proton number increases but same no. of shells)
- shielding is constant
- stronger attraction between nucleus and outermost electron
deviations/anomalies in first ionisation energy trends going across the period
slight decrease going from group 2 to 3
- electron is removed from a higher energy subshell (p block > s block) so is further from the nucleus and easier to remove
slight decrease from group 5 to 6
- p electrons can no longer occupy orbitals singly. like charge repulsion means it’s easier to remove e⁻
how can we use ionisation energies to find out about the structure of the atom
large increases/jumps in successive ionisation energies of a particular element: the electron is removed from a shell significantly closer to the nucleus
(eg between 3ʳᵈ and 4ᵗʰ IE of aluminium)
properties of metals and why
good conductors (delocalised e⁻ can carry charge)
high melting and boiling points (strong electrostatic forces of attraction)
malleable and ductile (layers of ions can slide over one another)
