Periodicity

What does s,p,d,f blocks show about an element?

That the outer electrons are in that orbital

Why do metals melting and boiling points increase across a period?

Due to the increasing strength of the metallic bonding which increase across the period because the ions have a larger positive charge, they have a smaller ionic radius, there will be more delocalised electrons, all these result in a stronger electrostatic attraction between metal cations and delocalised electrons which requires increasing energy to overcome

Why does atom radius decrease across a period?

As protons are added across a period the nuclear charge increases resulting in a stronger electrostatic attraction between the nucleus and the outer shell electrons drawing them closer to the nucleus. Electrons added across a period go into the outer energy level adding to shielding. Increasing nuclear charge and minimal shielding causes stronger electrostatic attraction causing atomic radius to decrease across period

What is ionisation?

The process of removing one or more electrons from an atom or molecule and requires an input of energy so is endothermic

What is first ionisation energy?

The energy needed to remove 1 electron from each atom in 1 mole of gaseous atoms to form 1 mole of gaseous 1+ ions

Three factors that affect ionisation energy?

1. Nuclear charge- atoms with more protons in their nucleus have a stronger positive charge which creates stronger electrostatic attraction between the nucleus and the outer electrons

2. Atomic radius- electrostatic attraction drop off steeply with increasing distance. Electrons in smaller atoms are held closer to the nucleus so the attraction is greater

3. Electron shielding- inner electron shells shield the outermost electrons from the full attractive force of the nucleus reducing the effective nuclear charge experienced by the outer electrons

Why does ionisation energy decrease down groups?

Because the atomic radius increases down the the group as more shells are added and this also increases shielding which are greater than the nuclear charge effect, leading to an overall decrease in ionisation energy as you move down the group

Why does ionisation energy increase across a period?

Nuclear charge increases as more protons are added and this outweighs the similar shielding across the period

Why has group 3 got a lower first ionisation energy then group 2?

As in group 3 the electron is removed from a p orbital rather than an s orbital like in group 2, p orbitals have slightly higher energy than s orbitals so the outermost electron is on average further from the nucleus, the p orbital also experiences additional shielding from the nucleus provided by the s electrons, as a result less energy is required to remove the outermost p electron from the group 3 element compared to removing the outermost s electrons from the group 2 element

Why has group 6 got a lower first ionisation energy than group 5?

In group 5 the electron is removed from a singly occupied orbital, in group 6 the electron is removed from an orbital containing two electrons, the paired electrons in the group 6 element experience greater electron electron repulsion, as a result less energy is needed to remove one of these paired electrons in the group 6 element compared to the unpaired electron in the group 5 element

What is successive ionisation energies?

The energy to remove each successive electron

What is second ionisation energy?

The energy needed to remove 1 electron from each ion in 1 mole of gaseous 1+ ions to form 1 mole of gaseous 2+ ions

Why is second ionisation energy greater than first?

As successive electrons are removed from the same shell the remaining electrons experience greater electrostatic attraction to the increasingly positive nucleus. This increased nuclear attraction requires more energy to remove the next electron from that shell

Why is there an increase in ionisation energy between shells?

Because the closer the shell is to the nucleus the harder it is to remove the electron so more energy needed