Inorganic Chemistry ACS: Study Guide
Exam|334 Questions with Accurate
Answers
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 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
Exam|334 Questions with Accurate
Answers
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 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
