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Shape of orbitals

Shape of orbitals. s-orbital p-orbitals.

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Shape of orbitals

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  1. Shape of orbitals s-orbital p-orbitals

  2. Electrons can only occupy so-called atomic orbitals with well defined energy levels corresponding to the principal quantum number, n. The lowest level will have n = 1, the next n = 2, and so on.Electrons must always enter the first available orbital of lowest energy.The first element, hydrogen, only has one electron, and so this electron must enter the 1s orbital.The electronic configuration of hydrogen in the ground state must therefore be: H 1s1.

  3. Boron 1s2 2s2 2p1 Atomic number 5 Carbon 1s2 2s2 2p2 Atomic number 6 According to Hund’s rule, fill a set of similar energy orbitals with as many unpaired electrons as possible (as in the p or d orbitals).

  4. The successive ionization energies of sodium [Ar] 3s1 • One electron furthest away from nucleus easiest to remove (3S1) • Eight nearer to nucleus, harder to remove (2p6, 2p2) • Two very close to the nucleus and most difficult to remove (1s2) • Proves numbers of electrons in each shell is 2,8,1 2p6, 2p2

  5. First ionisation energies (IE) of elements in periods 2 to 4/KJmol-1 IEs increase across period (nuclear charge increasing, shielding the same) IEs decrease down a group (distance of outer electrons from nucleus increasing)

  6. First ionisation energies (IE) of elements in periods 2 to 4/KJmol-1 • Look at group II and group III: the first electron for Al is a 3P1 electron and is less tightly held than 3S electron as it is further away from nucleus • Look at group V and group VI: a paired electron is easier to remove than an unpaired electron and the configuration for sulfur is 3P4, so it has a paired electron

  7. The Ionization Energies for elements in Period 3 I1 = first ionization energy; A(g) A+(g) + e- I2 = second ionization energy; A+(g) A2+(g) + e- I3 etc.

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