Periodicity Flashcards
Bonding and structure (first 20 elements)
Metallic lattice
Li, Be, Na, Mg, Al, K, Ca
Covalent molecular
H2, N2, O2, F2, Cl2, P4, S8 and fullerenes (eg C60)
Covalent network
B, C (diamond and graphite), Si
Monatomic
Noble gases
Describe the term covalent radius and explain the trends on the periodic table
The covalent radius is a measure of the size of an atom.
Across a period, the covalent radius decreases as the nuclear charge increases (higher nuclear charge pulls the outer electrons in closer).
Going down a group the covalent radius increases as the number of occupied shells increase (more shells = bigger atom).
State the definition of the terms first and second ionisation energy
First ionisation energy is the energy required to remove one mole of electrons from one mole of gaseous atoms.
Second ionisation energy is the energy required to remove the 2nd mole of electrons.
Explain the trends in ionisation energy on the periodic table
Across a period, the ionisation energy tends to increase as the nuclear charge increases.
Down a group the ionisation energy decreases as there are more shells, which means there is an increased screening effect due to the inner electrons (making it easier to remove the outer ones).
There are problem solving questions relating to difference in ionisation energies – remember that it is harder to remove outer electrons from ‘full’ shells.
State the definition of the term electronegativity
Electronegativity is a measure of the attraction which an atom (involved in a bond) has for the electrons in the bond.
Explain the trends in electronegativity on the periodic table
Across a period the electronegativity values increase due to the nuclear charge increasing (higher positive nuclear charge will mean the negative electrons will be more attracted).
Going down a group the electronegativity decreases, this is due to the screening effect (of the inner shell electrons) – they will ‘shield’ the nuclear charge.
