Atomic Structure Flashcards
What is the degeneracy of the p; d; f orbitals
p orbital = 3 fold
d orbital = 5 fold
f orbital = 7 fold
What are the reasons for the energy ordering
- Repulsions among electrons
- Relative amplitudes of the radial portions of the atomic orbital wave functions - compare the extent to which the radial portions of the various wave functions penetrate to the nucleus
Explain Radial penetration of the wave functions
An electron in an atomic orbital with electron density close to the nucleus is stabilized by close interaction with the positive charge of the nucleus.
Explain the stability ordering
- For any n: s-orbital most electron density close to the
nucleus, then p, d, f. - Order of stability: s > p > d > f
Aufbau principle
- build up from the bottom
- orbitals of lowest energy first
Hund’s rule
- Where orbitals are available in degenerate sets, maximum spin multiplicity preserved (electrons are not paired until each orbital in a degenerate set is half-filled)
Pauli exclusion principle
- No two electrons may have the
same set of 4 quantum numbers - Two electrons in the same orbital
must have opposite signs
ms = +½ and ms = −½
- maximum 2 electrons per orbital
What is the orbital order
1s 2s. 2p. 3s. 3p. 4s. 3d 4p. 5s. 4d. 5p. 6s. 4f. 5d. 6p. 7s.
What electron configuration does period 3 start with
[ Ne ] 3s^1
Explain the electron configuration of each period
- Period 4 started with [ Ar ] 4s, then 3d and ended with 4p
- Period 5 starts with 5s, then 4d and ends with 5p.
What is the electron configuration for 24Cr
[Ar] 4s13d5 (not [Ar] 4s23d4)
What is the electron configuration for 29Cu
Ar] 4s13d10 (not [Ar] 4s23d9)
What is the electron configuration for 42Mo:
[Kr] 5s14d5 (not [Kr] 5s24d4)
What is the electron configuration for 47Ag:
[Kr] 5s14d10 (not [Kr] 5s24d9)
What is the electron configuration for 41Nb
[Kr] 5s14d4 (not [Kr] 5s24d3)
What is the electron configuration for 44Ru
[Kr] 5s14d7 (not [Kr] 5s24d6)
What is the electron configuration for 45Rh
[Kr] 5s14d8 (not [Kr] 5s24d9)
What is the electron configuration for 46Pd
[Kr] 5s04d10 (not [Kr] 5s24d8)
What is the electron configuration for 78Pt
[Xe] 6s15d9 (not [Xe] 6s25d8)
What is the electron configuration for 79Au
[Xe] 6s15d10 ( not [Xe] 6s25d9)
What is the electron configuration for 64Gd
[Xe] 6s2 5d1 4f7 (not [Xe] 6s25d04f8)
Explain Mendeleev
- Arranged in horizontal rows
- Related elements in vertical columns
Explain Moseley
- Proper sequence: atomic number
- Vertical columns contain chemically similar elements and electronically similar atoms
- Periodic arrangement of electron configurations give the same result as what Mendeleev deduced from chemical observations.
Explain Hund’s rule fully
- Electrons make use of different orbitals as far as possible
(within a set of degenerate orbitals)
A consequence of the charge of the electrons – electrons minimise repulsive forces by occupying different p orbitals
- Parallel spins (same value of ms) are used until the
Pauli exclusion principle requires pairing of spins
Early pairing of spins leads to less stable arrangements
Explain Most stable electronic state
The state with maximum spin multiplicity (largest number of (un)paired electron spins)
What does Hund’s rule imply
- Spreading out of electrons as wide as possible in the space surrounding the nucleus
- This spreading of electron density leads to extra stability
What does the amount of nuclear effective charge felt by an electron depends on
- The type of orbital in which the electron is housed
- The ability of other electrons in more penetrating orbitals to screen (or shield) the electron from the
nucleus
When will an electron experience the greatest effective nuclear charge
- An electron will experience the greatest effective nuclear charge when housed in
an s orbital (for any given n), than p, etc. (Reason for Aufbau) - Geometry and orientation of the p orbitals makes them poor at shielding one
another from the nucleus.
Explain poor electron shielding
Where screening of an electron is poor, the effective nuclear charge is high
( In terms of Hund’s rule: Electrons spread out into a degenerate set of orbitals in
order to experience the maximum effective nuclear charge )
Explain Slater’s screening constant, σ
- Slater proposed an empirical constant that represents the cumulative extent to which the other electrons of an atom shield (or screen) any particular electron from the nuclear charge
- Z* is always smaller than Z (the electron is screened or shielded from Z by an
amount σ). - The value of σ for the electron in question is the cumulative total provided by
the other electrons of the atom– calculated using Slater’s rules.
To conclude
Z* increase continually from left to right across the rows of the periodic table because of imperfect shielding
