Inorganic Chemistry ACS Study Guide
Ionization energy energy required to remove the least tightly bound electron from a neutral atom in
the gas phase
periodic trend of ionization energy highest at top right-smaller electron=harder to remove
Why is a half filled subshell so stable? it serves to maximize the stabilizing interactions while minimizing the
destabilizing interactions among electrons
exchange interaction pie, stabilizing, result of electrons pairing in degenerate orbitals with parallel spin
pairing energy destabilizing, coulomb interaction, pic, energy of electron-electron repulsion in a
filled orbital
Is it easier to ionize a high energy or low energy high energy electron-already contains more energy so it requires less energy
electrons input
What happens when a 3d series metal is ionized? the first electron to be ionized will come from the 4s orbital, the other s electron
will enter the d orbital (4s03dn+1)
lanthanide contraction reduction in atomic radius following the lanthanide series, contrary to the overall
trend observed for the periodic table
lanthanides elements 57-71, first appearance of f orbitals, f orbitals are poor at shielding so
any electrons dded will have a higher Zeff, shrinking the radius
Slater's rules tell us what the effective nuclear charge will be, Zeff=Z-sigma, Z is the atomic
number, sigma=sum of the number of electrons in a given subtle multiplied by a
weighting coefficient (page 1)
Shielding the reduction in charge attraction between the nucleus and electrons due to
electrons between the nucleus and the electron in question, it is considered the
be between if it has a lower energy
penetration when an electron of a higher atomic orbital is found within the shell of electrons
of a lower atomic number, that is to say that an electron of higher energy is found
within an orbital of lower energy
electron affinity the difference in energy for a neutral gaseous atom, and the gaseous anion.
used interchangeably with electron gain enthalpy. more positive=more stable EA
with the additional electron, more positive EGE=more stable with extra electron
Combination of electron affinity and ionization energy electronegativity, overall measure of an atoms ability to attract electrons to itself
when part of a compound, fluorine has highest electronegativity
polarizability an atoms ability to be distorted by an electric field, regions of a molecule can
take on partial positive or partial negative charge
Why do we use the hydrogen system approximation systems involving multiple electrons are much more complex, and they require
the use of quantum mechanics
, Inorganic Chemistry ACS Study Guide
What is the formula for the energy of a hydrogen orbital E=-13.6(eV)*(Z^2/n^2), h is plancks constant (background on pg 4)
Energy can be expressed in... Joules, wavenumber, inverse centimeters
quantum number N principle quantum number, defines energy and size of orbital
quantum number L orbital angular momentum quantum number, defines the magnitude of the orbital
angular momentum, as well as the angular shape of the orbital, L can have
values of 0 to n-1.
quantum number Ml magnetic quantum number, describes the orientation of the angular momentum,
ml can have values of 0 to +/-1
quantum number Ms spin magnetic quantum number, defines intrinsic angular momentum of an
electron, Ms can have values of either +1/2 or -1/2
Radial wavefunction (R(r)), along with the angular wavefunction, gives us the orbitals. With a wave
function it is possible to completely characterize a particle, goes to zero at infinity,
produce characteristic shapes when graphed
Radial distribution function a plot of R^2(r)r^2, tells us probability of finding an electron at a certain distance
from the nucleus, every orbital has a different radial distribution function and a
node on the graph is a region of zero probability
Bohr radius the most probably distance to find the electron in a one proton, one electron
system (52.9 pico-meters)
What orbitals correspond to l=0 through l=4 L=0=s, L=1=p, L=2=d, L=3=f, L=4=g
Building up principle/Hund's rule when degenerate orbitals are available for occupation, electrons occupy
separate orbitals with parallel spin
Pauli exclusion principle no more than two electrons can occupy a single orbital, and to do so, their spins
must be paired
Descibe VSEPR purpose is to predict molecular geometries, basic assumption is that regions of
enhanced electron density take positions as far apart as possible in order to
minimize repulsive forces.
Relative repulsion strengths VSEPR lone pair> multiple bonds> single bonds
Valence bond theory explains chemical bonding by considering the overlap of tomic orbitals, wave
patterns of atomic orbitals interfere constructively to form a bond, sigma is
formed when orbital overlap has cylindrical symmetry, pi bond forms when they
overlap side by side after the formation of a sigma bond
How is hybridization used in valence bond theory explains bonding where the number of equivalent bonds exceeds the number of
valence orbitals
, Inorganic Chemistry ACS Study Guide
Effect of a lone pair on geometry? it pushes strongly against all other substituent. It is the strongest force governing
the shape of a molecule
Molecular orbital theory an improvement over valence bond theory in that the bonding description
extends to all atoms in a molecule and handles polyatomic molecules, atomic
orbitals combine to form molecular orbitals which are delocalized descriptions of
electron distribution
MO theory assumptions orbital approximation, linear combinations of atomic orbitals
Orbital approximation the wave function describing all of the electrons of a molecule can be written as a
product of the one electron avefunctions
linear combination of atomic orbitals the superposition of multiple atomic orbitals of same type along with weighting
coefficients
H2 and H2-like molecules the 1s orbitals are of equal energy so they lie at the same level of the diagram,
atomic orbitals combine to form one sigma orbital which is lower in energy than
one anti bonding sigma orbital which is higher in energy
Li2 through N2 (period 2) energy contributing to each atom is the same, 2s orbitals combine to form a
bonding 1sigmag and an anti bonding sigmau, as well as a bonding 2sigmag,
although that MO is mainly 2p in character. 2p combine as follows: lowest energy
orbitals are two degenerates 1piu followed by 2sigmag. Following that is anti
bonding orbital 1pig (doubly degenerate) ad 2sigmau
Period 2: O2, F2 etc, why not Ne2? difference is MO diagram's relative energies of the bonding orbitals piu and pig
H-X H is something that bonds through 2p and 2s orbitals, relative energies of the
atomic orbitals are taken into account because there are two separate atoms,
H1s is usually higher in energy, 4 atomic orbitals of X will mix with 1s of H to give
5 molecular orbital. Anti bonding is mostly H in character while the other 4 are
mostly X
X-Y relative energies are taken into account, more electronegative=lower energy of
orbitals
Relative energies of MOs 4sigma>2pi>3sigma>1pi>2sigma>1sigma
Plyatomic MO must use approximations of MO theory. Linear combinations of Atomic orbitals
create a cumulative approximation of the atomic orbitals of its constituent atoms,
compare orbitals to determine what will mix
LCAO notation a,b=nondegenerate, e=doubly degenerate, t=triple degenerate
LCAO energies lowest=a (nondegenerate)
intermediate=e
highest=t
LCAO linear combinations of atomic orbitals
Ionization energy energy required to remove the least tightly bound electron from a neutral atom in
the gas phase
periodic trend of ionization energy highest at top right-smaller electron=harder to remove
Why is a half filled subshell so stable? it serves to maximize the stabilizing interactions while minimizing the
destabilizing interactions among electrons
exchange interaction pie, stabilizing, result of electrons pairing in degenerate orbitals with parallel spin
pairing energy destabilizing, coulomb interaction, pic, energy of electron-electron repulsion in a
filled orbital
Is it easier to ionize a high energy or low energy high energy electron-already contains more energy so it requires less energy
electrons input
What happens when a 3d series metal is ionized? the first electron to be ionized will come from the 4s orbital, the other s electron
will enter the d orbital (4s03dn+1)
lanthanide contraction reduction in atomic radius following the lanthanide series, contrary to the overall
trend observed for the periodic table
lanthanides elements 57-71, first appearance of f orbitals, f orbitals are poor at shielding so
any electrons dded will have a higher Zeff, shrinking the radius
Slater's rules tell us what the effective nuclear charge will be, Zeff=Z-sigma, Z is the atomic
number, sigma=sum of the number of electrons in a given subtle multiplied by a
weighting coefficient (page 1)
Shielding the reduction in charge attraction between the nucleus and electrons due to
electrons between the nucleus and the electron in question, it is considered the
be between if it has a lower energy
penetration when an electron of a higher atomic orbital is found within the shell of electrons
of a lower atomic number, that is to say that an electron of higher energy is found
within an orbital of lower energy
electron affinity the difference in energy for a neutral gaseous atom, and the gaseous anion.
used interchangeably with electron gain enthalpy. more positive=more stable EA
with the additional electron, more positive EGE=more stable with extra electron
Combination of electron affinity and ionization energy electronegativity, overall measure of an atoms ability to attract electrons to itself
when part of a compound, fluorine has highest electronegativity
polarizability an atoms ability to be distorted by an electric field, regions of a molecule can
take on partial positive or partial negative charge
Why do we use the hydrogen system approximation systems involving multiple electrons are much more complex, and they require
the use of quantum mechanics
, Inorganic Chemistry ACS Study Guide
What is the formula for the energy of a hydrogen orbital E=-13.6(eV)*(Z^2/n^2), h is plancks constant (background on pg 4)
Energy can be expressed in... Joules, wavenumber, inverse centimeters
quantum number N principle quantum number, defines energy and size of orbital
quantum number L orbital angular momentum quantum number, defines the magnitude of the orbital
angular momentum, as well as the angular shape of the orbital, L can have
values of 0 to n-1.
quantum number Ml magnetic quantum number, describes the orientation of the angular momentum,
ml can have values of 0 to +/-1
quantum number Ms spin magnetic quantum number, defines intrinsic angular momentum of an
electron, Ms can have values of either +1/2 or -1/2
Radial wavefunction (R(r)), along with the angular wavefunction, gives us the orbitals. With a wave
function it is possible to completely characterize a particle, goes to zero at infinity,
produce characteristic shapes when graphed
Radial distribution function a plot of R^2(r)r^2, tells us probability of finding an electron at a certain distance
from the nucleus, every orbital has a different radial distribution function and a
node on the graph is a region of zero probability
Bohr radius the most probably distance to find the electron in a one proton, one electron
system (52.9 pico-meters)
What orbitals correspond to l=0 through l=4 L=0=s, L=1=p, L=2=d, L=3=f, L=4=g
Building up principle/Hund's rule when degenerate orbitals are available for occupation, electrons occupy
separate orbitals with parallel spin
Pauli exclusion principle no more than two electrons can occupy a single orbital, and to do so, their spins
must be paired
Descibe VSEPR purpose is to predict molecular geometries, basic assumption is that regions of
enhanced electron density take positions as far apart as possible in order to
minimize repulsive forces.
Relative repulsion strengths VSEPR lone pair> multiple bonds> single bonds
Valence bond theory explains chemical bonding by considering the overlap of tomic orbitals, wave
patterns of atomic orbitals interfere constructively to form a bond, sigma is
formed when orbital overlap has cylindrical symmetry, pi bond forms when they
overlap side by side after the formation of a sigma bond
How is hybridization used in valence bond theory explains bonding where the number of equivalent bonds exceeds the number of
valence orbitals
, Inorganic Chemistry ACS Study Guide
Effect of a lone pair on geometry? it pushes strongly against all other substituent. It is the strongest force governing
the shape of a molecule
Molecular orbital theory an improvement over valence bond theory in that the bonding description
extends to all atoms in a molecule and handles polyatomic molecules, atomic
orbitals combine to form molecular orbitals which are delocalized descriptions of
electron distribution
MO theory assumptions orbital approximation, linear combinations of atomic orbitals
Orbital approximation the wave function describing all of the electrons of a molecule can be written as a
product of the one electron avefunctions
linear combination of atomic orbitals the superposition of multiple atomic orbitals of same type along with weighting
coefficients
H2 and H2-like molecules the 1s orbitals are of equal energy so they lie at the same level of the diagram,
atomic orbitals combine to form one sigma orbital which is lower in energy than
one anti bonding sigma orbital which is higher in energy
Li2 through N2 (period 2) energy contributing to each atom is the same, 2s orbitals combine to form a
bonding 1sigmag and an anti bonding sigmau, as well as a bonding 2sigmag,
although that MO is mainly 2p in character. 2p combine as follows: lowest energy
orbitals are two degenerates 1piu followed by 2sigmag. Following that is anti
bonding orbital 1pig (doubly degenerate) ad 2sigmau
Period 2: O2, F2 etc, why not Ne2? difference is MO diagram's relative energies of the bonding orbitals piu and pig
H-X H is something that bonds through 2p and 2s orbitals, relative energies of the
atomic orbitals are taken into account because there are two separate atoms,
H1s is usually higher in energy, 4 atomic orbitals of X will mix with 1s of H to give
5 molecular orbital. Anti bonding is mostly H in character while the other 4 are
mostly X
X-Y relative energies are taken into account, more electronegative=lower energy of
orbitals
Relative energies of MOs 4sigma>2pi>3sigma>1pi>2sigma>1sigma
Plyatomic MO must use approximations of MO theory. Linear combinations of Atomic orbitals
create a cumulative approximation of the atomic orbitals of its constituent atoms,
compare orbitals to determine what will mix
LCAO notation a,b=nondegenerate, e=doubly degenerate, t=triple degenerate
LCAO energies lowest=a (nondegenerate)
intermediate=e
highest=t
LCAO linear combinations of atomic orbitals
