Chapter 7: Multi-Electron Species and Periodic Properties Flashcards
1
Q
What does the one-electron wavefunction consider and ignore? (2)
A
- attraction between electrons and the nucleus is considered
- electron-electron repulsion is ignored
2
Q
Where is the electron repulsion greatest?
A
- it is greatest in the overlapping areas of the probability distributions between orbitals
3
Q
What does electron-electron repulsion depend on and what does it affect?
A
- depends on subshell (l)
- affects the energies of the subshells
4
Q
Does the energy formula from the one-electron species apply for multi-electron species?
A
- no longer applies and subshells will have different energies
5
Q
effective nucleur charge
A
- nuclear charge that the valence electron “feels”, Z_eff
6
Q
What is an approximation of Zeff`
A
Z - # of shielding electrons
7
Q
shielding electrons
A
- electrons that are not in the valence shell
8
Q
spin quantum number (4)
A
- m_s
- allowed values: +1/2 and -1/2
- concept of electron spin
- represented by half-headed arrows
9
Q
orbital diagram
A
- each orbital is represented by a horizontal line and electrons are shown as half arrows pointing up or down, name of the orbital is written beneath the line
10
Q
What does 1s^22s^1 tell us (2)
A
- two electrons in the 1s orbital
- one electron in the 2s orbital
11
Q
electron configuration
A
- the name of the orbital is followed by the # of electrons in the orbital superscript
12
Q
Aufbau principle (the “building up” principle) (2)
A
- filling the orbitals with electrons from the lowest energy first: in a ground-state-multi-electron atom or ion, the electrons always occupy the lowest energy orbitals first available
- a maximum of 2 electrons can occupy each orbital, as long as they have opposite spin quantum numbers
13
Q
Pauli exclusion principle
A
- no 2 electrons in an atom or ion may have the same 4 quantum numbers
14
Q
What 4 quantum numbers are multi-electron species dependent on?
A
- n, l, m_s, m_l
15
Q
Hund’s rule
A
- when orbitals that have the same energy are available, electrons occupy them singly with the same spin before being paired within an orbital
16
Q
denegerate
A
- when orbitals have the same energy
17
Q
What is the general procedure for writing electronic configurations of multi-electron species?
A
- locate the element in the periodic table noting the period (n) the element is in and the block (s,p,d,f)
- identify the noble gas in the previous period (n-1) and write its symbol in square brackets
- starting from the left of the periodic table in period n, write the subshells and occupancies in order until the element is reached
- for ionic species (of the s or p blocks), simply add to (for anions) or subtract from (cations) the configuration for the neutral atom
18
Q
What are the exceptions to the regular electron filling? (2) (pt1)
A
- 4s and 3d orbitals are close in energy and their relative energies can change depending upon the particular electronic configuration of a d-block atom
- chromium: more energetically favourable to have one electron in the 4s orbital and one electron in each of the five 3d orbitals than to have two electrons in the 4s orbital and four electrons in the five 3d orbitals
- copper: configuration of [Ar]4s^13d^10, rather than the expected [Ar]4s^23d^9
19
Q
What are the exceptions to the regular electron filling? (pt2)
A
- transition metal ions: when transition metals are ionized, the remaining electrons will be found in the d subshell of the valence shell
- for example Sc: [Ar]4s^23d^1, Sc[I] ion Sc+; [Ar]3d^2
20
Q
Excited state of a muti-electron species
A
- when a species violates either Aufbau principle and/or Hund’s rule
21
Q
Which electrons possess an overall magnetic moment?
A
- only species with unpaired electrons
22
Q
paramagnetic
A
- species with one or more unpaired electrons
23
Q
diamagnetic
A
- species with no unpaired electrons
24
Q
What can be explained by considering the effect of increasing nuclear charge (Z), which decreases total energy, and increasing electronic repulsion, which increases total energy, within atoms and ions
A
- atom and ion size, ionization energy, and electron affinity
25
How does increasing nuclear charge (Z) affect energy?
- decreases total energy
26
How does increasing electronic repulsion affect energy?
- increases total energy
27
covalent radius
- estimate for radius which is applied to single-bonded homonuclear species
28
ionic radius
- estimate made from the crystal lattice structure and uses crystallography measurements of distances between nuclei in crystal lattice
29
van der waal radius
- uses distances between non-bonded atoms to estimate the boundary of an atom
30
What happens to atomic radii as we move in the periodic table and why? (2)
- atomic radii decreases as we move from left to right because distributions shift closer to the nucleus as Z and Z_eff increase
- the number of shells (n) in a species affects the radius more than the nuclear charge Z: as you go down the periodic table, n increases and the atomic radius increases
31
Are cations smaller or larger than the neutral atom and why? (2)
- cations experience less electron-electron repulsion (greater Z_eff) then neutral counterparts and removal of all valence shell electrons leaves core electrons in the inner (n-1) shell
- cations are smaller than neutral atoms
32
Are anions smaller or larger than the neutral atom and why? (2)
- anions experience more electron-electron repulsion (smaller Z_eff) then corresponding neutral atoms
- anions are larger than neutral atoms
33
Why is the atomic radius of Al have a greater atomic radius than Ga, breaking the trend?
- Ga has more protons(?) in 3d, 4d and 5d cells
34
Is an input or output of energy needed to move an electron further from the nucleus?
- given that electrons are held to the nucleus by electrostatic forces of attraction, an input of energy is required
35
ionization energy
- minimum energy required to remove a single electron from an atom, molecule, or ion in its gaseous state
- proportional to the magnitude of electrostatic attraction between the electron being removed and the nucleus, combined with the electrostatic repulsion from the remaining electrons
36
why is ionization energy and electron affinities measured in the gaseous state?
- so that energy from interactions between species (liquid and solid) does not affect the measurement
37
what is the trend of ionization energy as we move left to right on the periodic table and why?
- ionization energy generally increases
- valence orbitals are closer to the nucleus on the right side so the electrons feel a stronger force of attraction to the nucleus, with effective Z_eff, and require more energy to be ionized
38
what is the trend of ionization energy as we move from top to bottom on the periodic table and why?
- generally decreases
- size of the atom increases from top to bottom because the energy of attraction between the nucleus, with effective Z_eff, and valence electrons is inversely proportional to the distance between them
39
electron affinity
- energy change that results from the addition of a single electron to an atom or ion in its gaseous state
40
Is an input or output of energy needed to move an electron closer to the nucleus?
- output or release of energy
- if an electron is added to the right side of the periodic table, it will feel a greater attraction to the nucleus because the orbital it will occupy is lower in energy so more energy is released on the right side when elements gain an electron
41
What is the trend of electron affinity on the periodic table?
- electron affinity increases from left to right until the halogens as the noble gases have very low electron affinities
- electron affinity decreases going down the periodic table as the added electron will go into an orbital that is further away from the nucleus (higher in energy) and will experience a lower force of attraction
