Chapter 3 part B Flashcards
Orbital Energies
- Sublevels are not the same energy (except hydrogen)
- Exact values depend on elements
- Energy: s < p < d < f
-4s lower than 3d
-5s lower in energy than 4d
Effective nuclear charge
Sublevels differ in their effective nuclear charge (Zeff)
-Li has 2 shells:
- Inner shell of 2 electrons (1s), outer shell of 1
electron (2s)
Effective Nuclear charge
The Zeef is higher for lower l
-Zeff: s > p > d > f
-the radial probability of the 2s orb. penetrates into the RP of the 1s orb
Hund’s rule
Hund’s rule states that electrons are added to the atomic orbitals with the same energy levels (degenerate orbitals) in such a way that each orbital is occupied by a single electron with the same spin (either +½ or -½) before any orbital can be occupied by two electrons. In this student’s diagram, one of the orbitals has two electrons while three orbitals at equal energies contain zero electrons, which is in violation of Hund’s rule.
From slides:
When placing 2 or more electrons in orbitals of same energy (within a subshell, degenerate orbitals), lowest energy attained when the number of electrons with the same spin is maximized
✓Electrons are negatively charge, so repel each other
✓Electrons in different orbitals repel each other less
✓Quantum mechanics says there is increased stability from having pairs of electrons of parallel spin: exchange energy
Aufbau Principle
-Electron configurations of atoms are built up by placing electrons into atomic orbitals from lowest to highest energy according to the Pauli exclusion principle
Transition Metal Configurations
• In the 4th row, for K and Ca, the 4s orbital is lower in energy than the 3d, therefore 4s fills first:
K [Ar] 4s1
Ca [Ar] 4s2
• For Sc - Zn (transition metals), 3d actually lowerin energy than 4s, yet electrons still placed in 4sfirst:
Sc [Ar] 3d14s2
Ti [Ar] 3d24s2
• 4s orbital is more diffuse than 3d orbitals →more electron-electron repulsion in 3d orbitals →more favorable to always “stash” 2 electrons in the 4s
• Exceptions:
Cr [Ar] 3d^5 4s^1 - These are exceptions as the atom wants to have something in every obital
Cu [Ar] 3d^10 4s^1
Core vs Valence
Noble gas is core electrons, and the d block are core electrons once you reach p
S and P are always valance
Configs & Element Properties
• Number of valence electrons determines the
behavior of an element
• Elements with same number of valence
electrons are in the same group
• Elements in same group have similar
reactivity!
• Noble gases have filled valence shells →
unreactive
• Elements with one more or one less electron
are very reactive
Atomic Radius
- Electrons have no fixed radii, or upper limits on their
radii
✓no exact measurements for atomic or ionic radii - Bonding or covalent radius: ½ the distance
between the nuclei of atoms that are bonded to each
other
-Left tends to be larger-(the electrons aren’t getting farther away but there are more protons
-Largest radii are bottom left corner
Ionic radius
-Cations smaller than the neutral atom
-lose outermost electrons
-Ion size increases down the group
-higher valence shell, large
-Anions larger than neutral atoms
-adding electrons to the outermost shell
- electron-electron repulsion increases
- nuclear charge stays the same
Isoelectronic species: same electron configuration
* Ionic radius decreases as atomic number increases
✓same number of electrons, but increasing nuclear charge
✓increases the effective nuclear charge
-Ionic size increases down the group
Ionization energies
- Ionization energy (IE): the amount of energy needed
to remove an electron from an atom/ion
✓gas state
✓always requires energy (positive)
✓units of Joules
-First ionization energy (IE1): amount of energy to remove one electron from the neutral atom
X(g) → X+(g) + 1 e–
-Second ionization energy (IE2): the amount of energy
to remove one electron from a +1 ion
✓Always requires more energy than the first
X+(g) → X2+(g) + 1 e– IE2
- IE decreases down the group
✓valence electron farther from the nucleus - IE generally increases left → right
✓effective nuclear charge increases - IE of transition metals varies little
Electron Affinity
- Electron affinity: amount of energy gained or
released when an atom or ion gains an electron
Electron affinity increases from left to right across the periodic table as effective nuclear charge increases. Electron affinity decreases down the periodic table because the electrons are less attracted to the nucleus when they are in shells, or energy levels, that are further from the nucleus.
Electronegativity
Electronegativity is a measure of an atom’s ability to attract electrons.
Anions electron configuration
Anions: add an electron to the lowest energy unoccupied or partially occupied orbitals
Cations
Remove electrons from the highest quantum number n (highest quantum number l within a given value to n)
