periodic trend in electron configuration and ionisation energy Flashcards
define first ionisation energy
removal of 1 mol of
electrons from 1 mol of gaseous atoms
what happens to atomic radius across a period?
decreases
- increased nuclear charge (increase in number of protons in the nucleus) creates greater attraction to pull valence electrons in closer
what happens to atomic radius down a group?
increases
* each group an electron shell is added, increases distance between valence electrons and nucleus, reducing attraction.
+ more shells increases shielding
therefore nuclear attraction reduces + atomic radius increases
what happens to ionisation energy down a group?
IE decreases due to an increasing atomic radius and electron shielding which reduces the effects of electrostatic forces of atrraction
what happens to first ionisation energy along a period?
IE increases due to an decreasing atomic radius and greater electrostatic forces of attraction
why do successive ionisation energies increase?
atomic radius decreases + there is a greater attraction between outer shell electrons and nucleus
what does a large jump in successive ionisation energies suggest?
which group an element is in
state and explain the trend in atomic radius across period 2
- decreases
- same shielding
- increased charge
- increased nuclear attraction
what are the exceptions in period 2?
Boron and oxygen have lower ionisation energies than expected
why does boron have a lower IE than expected?
- the energy difference between the 2s and 2p sub-shells
- the electron is being removed from a higher energy level that is further from nucleus so e- is held less strongly
define ionisation
the removal of one or more electrons from an atom.
explain how the properties of a metal link to structure & bonding
- giant metallic lattice
- metallic bonding
- lots of energy required to break bonds = high mp/bp
- delocalised electrons free to move = conducts electricity in solution & solid
explain how the properties of a non metal link to structure & bonding
- simple molecular
- covalent bonding
- weak induced dipole dipole forces
- require little energy to overcome so low mp/bp
- no mobile charge carriers so do not conduct electricity
explain how the properties of an ionic compound link to structure & bonding
- giant ionic lattice
- ionic bonding
- stronger than london forces, require more energy to break so higher mp/bp
- ions in fixed position in solid so cannot conduct electricity
- ions mobile in solution so conduct electricity
explain why reactivity increases down group 2
- increasing atomic radius
- with more shielding
- so nuclear attraction is weaker
- ionisation energy decreases
- outer e- lost more easily
explain why magnesium has a greater melting point than sodium
- Mg has more outer electrons
- Mg has a higher nuclear charge
- so nuclear attraction is greater
explain why phosphorus has a greater melting point than chlorine
- P4 has more electrons than Cl2
- so there are greater induced dipole dipole forces
- which require more energy to overcome
describe the trend in ionisation energy across a period
- increases
- nuclear charge increases
- similar shielding
- atomic radius slightly decreases due to greater attraction
- more energy required to remove outer e-
state and explain the trend in atomic radius across period 2
- decreases
- same shielding
- increased charge
- increased nuclear attraction
why do successive ionisation energies increase with ionisation number?
- radius decreases
- so nuclear attraction increases
explain why Mg has a greater first ionisation energy than Al
- Mg outer electrons in 3s subshell
- Al outer electrons in 3p subshell
- p subshells has higher energy
explain why there is a slight decrease in IE at aluminium & boron
- evidence of subshells
- outer electron sits in higher energy subshell
- which is further from the nucleus
explain why there is a slight decrease in IE at sulfur & oxygen
- evidence of electron repulsion
- P&S / N&O have outer electron in 3p orbital
- removing electrons takes it from a 2 electron orbital
- electrons repel so less energy is required to remove the electron with 2 in an orbital than 1 in an orbital
