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)
Effective Nuclear Charge Revisited
Zeff=Z – # of core electrons
-Zeff increases as you go right
-Valence electrons are shielded from the full strength of nuclear charge
* Effective nuclear charge (ENC, Zeff): nuclear charge (Z) minus
sum of shielding from other electrons (Si)
* Core electrons are very effective at shielding
* Electrons in same valence shell contribute shield only weakly
Coulomb’s Law
Coulomb’s Law states that the attraction between charges increases as the size of the charge increases.
a law stating that like charges repel and opposite charges attract, with a force proportional to the product of the charges and inversely proportional to the square of the distance between them.
Successive Ionization Energy
- Removal of each successive valence electron costs more energy
✓ higher effective nuclear charge
- Large increase in energy when start removing core electrons
✓ much higher effective nuclear charge
✓ break apart noble gas configuration
Zeff
Effective nuclear charge, Zeff, is defined as the true nuclear charge minus the charge that is shielded by electrons
As the core charge of an atom increases, what happens to the radii?
The radii decreases
What is the principle quantum number?
The principal quantum number is the shell number
Electron affinity
The electron affinity is the change in energy when an electron is added to a neutral gaseous atom to form an ion.
Electronegativity
Allen’s definition: average energy of valence electrons of
atoms in their ground electron configuration (configuration
energy)
✓ makes sense: valence electrons are involved in bonding, their
energies are what matter the most
✓ easy to measure experimentally
As you go down the periodic table it decreases
As you go right across the periodic table it increases
Summarize periodic trends
- Work with a partner to summarize the trends
on the periodic table (down and L→R) for:
✓Zeff
✓Atomic Radius
✓Ionic Radius
✓First Ionization Energy
✓Electron Affinity
✓Electronegativity
Periodic Law
When the elements are arranged in order of increasing atomic number, chemical properties recur periodically
Division of Periodic table:
- Metals: Left hand of the periodic table
✓solids at room temperature (except Hg)
✓shiny, conduct heat & electricity, malleable, ductile
✓Like to form cations - Nonmetals: right hand of the periodic table
✓found in all 3 states
✓do not conduct heat or electricity - Metalloids: zig-zags non metals
✓all solids
✓are semiconductors
✓other properties in between metals and nonmetals
Names of sections on the periodic table
- Alkali metals
- Alkaline earth metals
-Transition metals
-Halogens
-Noble gases
Ionic bonds
electrons are transferred from one atom to another, the resulting ions are held together by electrostatic forces
Covalent bonds
atoms are held together by the sharing of electrons (usually in pairs)
Metallic bond:
“Sea of electrons” free to move throughout the metal
Ionic compounds:
cations and anions held together by electrostatic forces.
✓Often (but not always), these form between metals and nonmetals.
* Ionic compounds are solids typically w/ high melting pts.
* Represented by the formula unit: the smallest electrically neutral collection of ions. Note that the total number of pos. charges must = the total number of neg. charges.
