Topic 1: Atomic Structure and Periodic Table Flashcards
Definition of relative atomic mass
The average mass of an atom relative to 1/12 of a C-12 Atom
Definition of relative isotopic mass
The average mass of an isotope relative to 1/12 of a C-12 atom
Definition of first ionisation energy
The energy needed to remove 1 mole of electron from 1 mole of gaseous atom
Definition of first ionisation energy
The energy needed to remove 1 mole of electron from 1 mole of gaseous atom to form 1 mole of gaseous +1 ion
Definition of successive ionisation energy
The energy needed to remove one mole of electron after another
What are ionisation energies influenced by?
Atomic Radius - Larger atomic radius = smaller I.E as the outer electrons are further from the nucleus so the attractive force from the nucleus is weaker
Nuclear Charge - Higher nuclear charge = higher I.E as there’s a greater positive charge in the nucleus so the attractive force is stronger felt by the electrons (Larger Proton to electron ratio)
Shielding - More shielding = smaller I.E as the inner shells repel the electrons in the outer shells
Definition of effective nuclear charge
The sum of protons - the sum of shielding electrons
Why does first I.E generally increase across the period?
Same period = Same number of shielding electrons, but the nuclear charge increases. Effective nuclear charge would increase. The P/e- ratio stays the same. So each electron would feel a greater force of attraction to the nucleus so I.E would increase.
When are the two decreases in I.E across the period?
Between s2p0 and s2p1 - The s2p1 is lower because further from nucleus and more shielded
Between s2p3 and s2p4 - one electron in each p orbital in group 5 while there’s one pair of e- in an orbital in group 6 so the repulsion makes the I.E lower
Why does the first I.E decrease down the group?
The same number of e- in the outer shell but more shielding e-. So larger atomic radius —> less attraction between e- and the nucleus and greater shielding effect
Describe the shape of an s and p orbital
S orbital - spherical
P orbital - dumbbell shaped (infinity shape)
What are the number of electrons that can fill and s, p and d orbital?
S - 2
P - 6
D - 10
What are the trend of successive ionisation energies?
Successive I.E increase between shells. Sudden rise in I.E after each shell as shielding e- decrease suddenly and so effective nuclear charge increases. Smaller atomic radius
Successive I.E increase within each shell. Larger P/e- ratio so more attraction between nucleus and e- and less repulsion between e-
Periodicity Trends
Atomic radius decreases —> same number of shells but more protons so stronger attraction between outermost e- and nucleus
Melting point trend according to structure -
Increase in metals with metallic bonding as smaller radius and more delocalised e- present
Giant covalent structure = highest m.p
Simple covalent molecules —> similar m.p
Monoatomic= lowest m.p
