Inorganic Chemistry ACS: Study Guide
Verified Questions And Answers
onization 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 electrons - - high energy
electron-already contains more energy so it requires less energy 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
-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
-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
Verified Questions And Answers
onization 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 electrons - - high energy
electron-already contains more energy so it requires less energy 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
-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
-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
