Chem 120 Exam 2 Study Guide
• Electron configuration
o Description of the orbitals occupied by electrons
o 1
1 s – electron configuration for hydrogen atom
▪ 1S is the orbital
▪ Little 1 is the number of electrons in orbital
• Orbital diagram – a square representing each orbital and a half-arrow representing each
electron in the orbital
• Find out wtf degenerate means
• 3 Rules of Electron Configuration
o Aufbau Principle
▪ Electrons ordinarily occupy orbitals of the lowest energy available first
▪ 1s, 2, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p etc.
▪ When the electrons occupy the lowest energy orbitals, the atom is in its
ground state
▪ When the electrons occupy any other orbitals the atom is in an excited
state
o Pauli Exclusion Principle
▪ One atomic orbital can only accommodate two electrons, and these
electrons must have opposite spins
o Hund’s Rule
▪ Electrons in the same subshell occupy orbitals singly, with spins
parallel, before they are paired
• Number of shell corresponds with the period on the periodic table
• Condensed electron configurations – just write the previous noble gas and then continue
from wear that noble gas ends
o Cl: 1 s2
▪ Finish this
• Some irregularities occur when there are enough electrons to half-filled or filled s and d
orbitals
o Cr and Cu are exceptions
, o Half- filled d subshell
• Valence electrons – the electrons in the highest principal energy level (for main group
elements
o Most important in chemical bonding
• Core electrons – electrons in lower energy level
• Elements in the same group on the periodic table have the same valence electron
configuration and display similar chemical behavior
o For main group elements, # of valence electrons = group #
• Orbital Blocks in the Periodic Table
o Main group elements
▪ S-block: ns subshell fills
▪ P-block: np subshell fills
, o Transition elements:
▪ Dblock: (n-1)d subshell fills
o Put the other one hoe on ppt
• Writing Electron Configuration from the Periodic table
• Periodic Treds in the Atomic Properties
o Atomic radius
▪ The covalent atomic radius is one-half of the distance between covalently
bonded nuclei
▪ Increases in going down in a group
• Orbital size increases in successive principal quantum levels
▪ Decreases in going across a period from left to right
• Effective nuclear charge increases
• Valence electrons are drawn closer to the nucleus, decreasing the
size of the atom
▪ Size depends on attraction between valence electrons and effective
nuclear charge
• Effective nuclear charge – a net positive charge attracting a
particular electron
o Ionization energy
▪ Amount of energy required to remove an electron from the ground state
of a gaseous atom or ion
▪ Na (g) – Na+(g) +e- L=495 kJ/mol
▪ Called 1st ionization energy because it is the energy to remove the first
electron
▪ In general l1 decreases going down a group
• The electrons being removed are farther from the nucleus
▪ L1 increases going left to right
• Electrons added to the same principal quantum level cannot
completely shield the increasing nuclear charge and are generally
more strongly bound from left to right on the periodic table
o Electron affinity : energy change associated with the addition of an electron to a
gaseous atom
• Chemical bond: force that holds atoms together in a compound
• Electron configuration
o Description of the orbitals occupied by electrons
o 1
1 s – electron configuration for hydrogen atom
▪ 1S is the orbital
▪ Little 1 is the number of electrons in orbital
• Orbital diagram – a square representing each orbital and a half-arrow representing each
electron in the orbital
• Find out wtf degenerate means
• 3 Rules of Electron Configuration
o Aufbau Principle
▪ Electrons ordinarily occupy orbitals of the lowest energy available first
▪ 1s, 2, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p etc.
▪ When the electrons occupy the lowest energy orbitals, the atom is in its
ground state
▪ When the electrons occupy any other orbitals the atom is in an excited
state
o Pauli Exclusion Principle
▪ One atomic orbital can only accommodate two electrons, and these
electrons must have opposite spins
o Hund’s Rule
▪ Electrons in the same subshell occupy orbitals singly, with spins
parallel, before they are paired
• Number of shell corresponds with the period on the periodic table
• Condensed electron configurations – just write the previous noble gas and then continue
from wear that noble gas ends
o Cl: 1 s2
▪ Finish this
• Some irregularities occur when there are enough electrons to half-filled or filled s and d
orbitals
o Cr and Cu are exceptions
, o Half- filled d subshell
• Valence electrons – the electrons in the highest principal energy level (for main group
elements
o Most important in chemical bonding
• Core electrons – electrons in lower energy level
• Elements in the same group on the periodic table have the same valence electron
configuration and display similar chemical behavior
o For main group elements, # of valence electrons = group #
• Orbital Blocks in the Periodic Table
o Main group elements
▪ S-block: ns subshell fills
▪ P-block: np subshell fills
, o Transition elements:
▪ Dblock: (n-1)d subshell fills
o Put the other one hoe on ppt
• Writing Electron Configuration from the Periodic table
• Periodic Treds in the Atomic Properties
o Atomic radius
▪ The covalent atomic radius is one-half of the distance between covalently
bonded nuclei
▪ Increases in going down in a group
• Orbital size increases in successive principal quantum levels
▪ Decreases in going across a period from left to right
• Effective nuclear charge increases
• Valence electrons are drawn closer to the nucleus, decreasing the
size of the atom
▪ Size depends on attraction between valence electrons and effective
nuclear charge
• Effective nuclear charge – a net positive charge attracting a
particular electron
o Ionization energy
▪ Amount of energy required to remove an electron from the ground state
of a gaseous atom or ion
▪ Na (g) – Na+(g) +e- L=495 kJ/mol
▪ Called 1st ionization energy because it is the energy to remove the first
electron
▪ In general l1 decreases going down a group
• The electrons being removed are farther from the nucleus
▪ L1 increases going left to right
• Electrons added to the same principal quantum level cannot
completely shield the increasing nuclear charge and are generally
more strongly bound from left to right on the periodic table
o Electron affinity : energy change associated with the addition of an electron to a
gaseous atom
• Chemical bond: force that holds atoms together in a compound
