Thermochem Flashcards
(14 cards)
What happens to atomic radii, electrostatic attraction, and electronegativity across a period?
The number of protons increases, increasing the nuclear charge. Electrons are being added to the same energy level with no additional shielding. So, since the charge on the nucleus increases, the electrostatic attraction between the positive nucleus and valence e- increases. This pulls valence e- closer to the nucleus, and atomic radius decreases.
What is electronegativity?
The ability of an atom to attract a pair of bonding electrons to itself. GREATER ELECTROSTATIC ATTRACTION, MEANS GREATER ELECTRONEGATIVITY.
Define shielding
The effect of the inner shells ‘blocking’ the valence electrons. They repel the outer electrons because they have the same charge.
What happens to atomic radii, electrostatic attraction, and electronegativity down a group?
The number of protons increases going down a group, increasing the nuclear charge. But, this is offset by the valence e- being further from the nucleus due to more energy levels being added. Shielding increases as the full inner energy levels shield the valence e- from the positive nucleus. So, the electrostatic attraction is less, so the atomic radius is larger.
Across a period, electronegativity…
Increases.
Down a group, electronegativity…
Decreases. (Increases up a group)
Compare cation radius with that of it’s parent atom
The atom and it’s cation have the same nuclear charge. The cation has had valence electrons removed, which decreases inter-electron repulsion. This means there’s a stronger attraction between the valence e- and nucleus. So, they’re pulled closer to the nucleus, making the cation is smaller than it’s parent atom.
Compare anion radius with that of it’s parent atom
The atom and it’s anion have the same nuclear charge. The anion has had valence electrons added, which increases inter-electron repulsion. This means there’s less attraction between the valence e- and nucleus. So, they move further from the nucleus, making the anion larger than it’s parent atom.
Define first ionisation energy
The minimum energy needed to remove one electron from from each atom in a mole of atoms in the gaseous state.
What happens to ionisation energy down a group?
The nuclear charge increases due to more protons in the atoms going down a group, but is offset by the increasing distance of valence electrons from the nucleus as more energy levels are added. The full inner energy levels shield the outer electrons from the protons in the nucleus so the electrostatic attraction decreases. So, the atomic radius increases. The further the valence e- are from the nucleus, the less energy needed to remove it.
What happens to ionisation energy across a period?
The nuclear charge increases due to more protons in the atoms going down a group. Electrons are being added to the same energy level with no additional shielding. So, since the charge on the nucleus increases, the electrostatic attraction increases. Thus atomic radius increases. The closer the valence e- are to the nucleus, the more energy needed to remove it.
What are the exceptions in electron configuration? Why?
Chromium (Cr) and Copper (Cu). Having half filled/fully filled orbitals is more stable, so they move electrons from 4s to 3d.
Explain the trend in ionisation energies going from Be to B and Mg to Al
Going from Be to B and Mg to Al there’s a decrease in ionisation energy.
The second element (B/Al) has valence electrons in p subshell, instead of s. P is further from nucleus, (larger atomic radius) therefore less ionisation energy.
Explain the trend in ionisation energies between Gp15 and Gp16 elements
Between N, O, P, S there’s a small decrease in IE.
Gp 15 elements have 3 unpaired electrons in different orbitals (this is in the p subshell). This is more stable than Gp 16 elements, which have 2 unpaired and 2 paired electrons in the p orbital. This has a little more electron-electron repulsion, so less IE.
