ACS INORGANIC CHEMISTRY STUDY GUIDE
1. Period 2: O2, F2 etc, why not Ne2?: difference is MO diagram's relativeenergies of the
bonding orbitals piu and pig
2. H-X: H is something that bonds through 2p and 2s orbitals, relative energies ofthe 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 5molecular orbital. Anti bonding
is mostly H in character while the other 4 are mostlyX
3. X-Y: relative energies are taken into account, more electronegative=lower ener- gy of
orbitals
4. Relative energies of MOs: 4sigma>2pi>3sigma>1pi>2sigma>1sigma
5. 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
6. LCAO notation: a,b=nondegenerate, e=doubly degenerate, t=triple degenerate
7. LCAO energies: lowest=a (nondegenerate)intermediate=e
highest=t
8. LCAO: linear combinations of atomic orbitals
9. Bond order: a method to assessing overall bond strength between two atomsin a molecule,
higher bond order=stronger bond, assesses net number of bonds between two atoms
10. Calculating bond order: (bonding-antibonding)/2
11. How to assign bonding character-bonding: a bonding orbital will lie lower inenergy than
its substituent atomic orbitals
12. How to assign bonding character-nonbonding: a nonbonding orbital willequal in energy
to its corresponding atomic orbitals
13. How to assign bonding character-antibonding: an antibonding orbital will liehigher in
energy than its substituent atomic orbitals
14. HOMO interactions: highest occupied molecular orbital, highest energy elec-trons reside
15. LUMO interactions: lowest unoccupied molecular orbital, it is the energy leveldirectly
above the homo
16. Importance of HOMO/LUMO: critical to electron structure, when excited an electron is
promoted from HOMO to LUMO. bonding and anti bonding characteristicsdetermine stability,
the frontier orbitals are the point of interaction and basis of understanding for molecular
,interaction
17. paramagnetism: if there are two unpaired electrons
18. metallic bonding: the bonding and ordering of metals into pure solids or solidsolutions,
viewed as enormous molecules with continually overlapping atomic or- bitals
19. ionic bonding: ions of different elements held together in rigid, symmetricalarrays as a
result of attraction between their opposite charges
20. lattice: a 3D infinite array of lattice points which define the repeating structureof a crystal
21. unit cell: an imaginary, parallel sided region from which the entire crystal canbe built,
such that it fits perfectly together, giving rise to a crystal system.
22. 7 types of crystal systems: cubic, tetragonal, orthorhombic, monoclinic, triclin-ic,
rhombohedral, hexaganol
23. Ionization energy: energy required to remove the least tightly bound electronfrom a
neutral atom in the gas phase
24. periodic trend of ionization energy: highest at top right-smaller electron=harderto remove
25. Why is a half filled subshell so stable?: it serves to maximize the stabilizinginteractions
while minimizing the destabilizing interactions among electrons
26. exchange interaction: pie, stabilizing, result of electrons pairing in degenerateorbitals
with parallel spin
27. pairing energy: destabilizing, coulomb interaction, pic, energy of electron-elec-tron
repulsion in a filled orbital
28. Is it easier to ionize a high energy or low energy electrons: high energyelectron-already
contains more energy so it requires less energy input
29. 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)
30. lanthanide contraction: reduction in atomic radius following the lanthanide se-ries,
contrary to the overall trend observed for the periodic table
31. lanthanides: elements 57-71, first appearance of f orbitals, f orbitals are poor atshielding
so any electrons dded will have a higher Zeff, shrinking the radius
32. 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)
33. 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
34. 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
, 35. electron affinity: the difference in energy for a neutral gaseous atom, and the gaseous
anion. used interchangeably with electron gain enthalpy. more posi-tive=more stable EA with
the additional electron, more positive EGE=more stablewith extra electron
36. 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
37. polarizability: an atoms ability to be distorted by an electric field, regions of amolecule
can take on partial positive or partial negative charge
38. Why do we use the hydrogen system approximation: systems involving multiple
electrons are much more complex, and they require the use of quantummechanics
39. 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)
40. Energy can be expressed in...: Joules, wavenumber, inverse centimeters
41. quantum number N: principle quantum number, defines energy and size oforbital
42. quantum number L: orbital angular momentum quantum number, defines themagnitude of
the orbital angular momentum, as well as the angular shape of the orbital, L can have values
of 0 to n-1.
43. quantum number Ml: magnetic quantum number, describes the orientation ofthe angular
momentum, ml can have values of 0 to +/-1
44. quantum number Ms: spin magnetic quantum number, defines intrinsic angularmomentum
of an electron, Ms can have values of either +1/2 or -1/2
45. Radial wavefunction: (R(r)), along with the angular wavefunction, gives us theorbitals.
With a wave function it is possible to completely characterize a particle, goes to zero at
infinity, produce characteristic shapes when graphed
46. 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 radialdistribution function
and a node on the graph is a region of zero probability
47. Bohr radius: the most probably distance to find the electron in a one proton, oneelectron
system (52.9 pico-meters)
48. What orbitals correspond to l=0 through l=4: L=0=s, L=1=p, L=2=d, L=3=f,L=4=g
49. Building up principle/Hund's rule: when degenerate orbitals are available foroccupation,
electrons occupy separate orbitals with parallel spin
50. Pauli exclusion principle: no more than two electrons can occupy a singleorbital, and to
do so, their spins must be paired
51. Descibe VSEPR: purpose is to predict molecular geometries, basic assumption is that
1. Period 2: O2, F2 etc, why not Ne2?: difference is MO diagram's relativeenergies of the
bonding orbitals piu and pig
2. H-X: H is something that bonds through 2p and 2s orbitals, relative energies ofthe 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 5molecular orbital. Anti bonding
is mostly H in character while the other 4 are mostlyX
3. X-Y: relative energies are taken into account, more electronegative=lower ener- gy of
orbitals
4. Relative energies of MOs: 4sigma>2pi>3sigma>1pi>2sigma>1sigma
5. 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
6. LCAO notation: a,b=nondegenerate, e=doubly degenerate, t=triple degenerate
7. LCAO energies: lowest=a (nondegenerate)intermediate=e
highest=t
8. LCAO: linear combinations of atomic orbitals
9. Bond order: a method to assessing overall bond strength between two atomsin a molecule,
higher bond order=stronger bond, assesses net number of bonds between two atoms
10. Calculating bond order: (bonding-antibonding)/2
11. How to assign bonding character-bonding: a bonding orbital will lie lower inenergy than
its substituent atomic orbitals
12. How to assign bonding character-nonbonding: a nonbonding orbital willequal in energy
to its corresponding atomic orbitals
13. How to assign bonding character-antibonding: an antibonding orbital will liehigher in
energy than its substituent atomic orbitals
14. HOMO interactions: highest occupied molecular orbital, highest energy elec-trons reside
15. LUMO interactions: lowest unoccupied molecular orbital, it is the energy leveldirectly
above the homo
16. Importance of HOMO/LUMO: critical to electron structure, when excited an electron is
promoted from HOMO to LUMO. bonding and anti bonding characteristicsdetermine stability,
the frontier orbitals are the point of interaction and basis of understanding for molecular
,interaction
17. paramagnetism: if there are two unpaired electrons
18. metallic bonding: the bonding and ordering of metals into pure solids or solidsolutions,
viewed as enormous molecules with continually overlapping atomic or- bitals
19. ionic bonding: ions of different elements held together in rigid, symmetricalarrays as a
result of attraction between their opposite charges
20. lattice: a 3D infinite array of lattice points which define the repeating structureof a crystal
21. unit cell: an imaginary, parallel sided region from which the entire crystal canbe built,
such that it fits perfectly together, giving rise to a crystal system.
22. 7 types of crystal systems: cubic, tetragonal, orthorhombic, monoclinic, triclin-ic,
rhombohedral, hexaganol
23. Ionization energy: energy required to remove the least tightly bound electronfrom a
neutral atom in the gas phase
24. periodic trend of ionization energy: highest at top right-smaller electron=harderto remove
25. Why is a half filled subshell so stable?: it serves to maximize the stabilizinginteractions
while minimizing the destabilizing interactions among electrons
26. exchange interaction: pie, stabilizing, result of electrons pairing in degenerateorbitals
with parallel spin
27. pairing energy: destabilizing, coulomb interaction, pic, energy of electron-elec-tron
repulsion in a filled orbital
28. Is it easier to ionize a high energy or low energy electrons: high energyelectron-already
contains more energy so it requires less energy input
29. 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)
30. lanthanide contraction: reduction in atomic radius following the lanthanide se-ries,
contrary to the overall trend observed for the periodic table
31. lanthanides: elements 57-71, first appearance of f orbitals, f orbitals are poor atshielding
so any electrons dded will have a higher Zeff, shrinking the radius
32. 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)
33. 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
34. 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
, 35. electron affinity: the difference in energy for a neutral gaseous atom, and the gaseous
anion. used interchangeably with electron gain enthalpy. more posi-tive=more stable EA with
the additional electron, more positive EGE=more stablewith extra electron
36. 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
37. polarizability: an atoms ability to be distorted by an electric field, regions of amolecule
can take on partial positive or partial negative charge
38. Why do we use the hydrogen system approximation: systems involving multiple
electrons are much more complex, and they require the use of quantummechanics
39. 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)
40. Energy can be expressed in...: Joules, wavenumber, inverse centimeters
41. quantum number N: principle quantum number, defines energy and size oforbital
42. quantum number L: orbital angular momentum quantum number, defines themagnitude of
the orbital angular momentum, as well as the angular shape of the orbital, L can have values
of 0 to n-1.
43. quantum number Ml: magnetic quantum number, describes the orientation ofthe angular
momentum, ml can have values of 0 to +/-1
44. quantum number Ms: spin magnetic quantum number, defines intrinsic angularmomentum
of an electron, Ms can have values of either +1/2 or -1/2
45. Radial wavefunction: (R(r)), along with the angular wavefunction, gives us theorbitals.
With a wave function it is possible to completely characterize a particle, goes to zero at
infinity, produce characteristic shapes when graphed
46. 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 radialdistribution function
and a node on the graph is a region of zero probability
47. Bohr radius: the most probably distance to find the electron in a one proton, oneelectron
system (52.9 pico-meters)
48. What orbitals correspond to l=0 through l=4: L=0=s, L=1=p, L=2=d, L=3=f,L=4=g
49. Building up principle/Hund's rule: when degenerate orbitals are available foroccupation,
electrons occupy separate orbitals with parallel spin
50. Pauli exclusion principle: no more than two electrons can occupy a singleorbital, and to
do so, their spins must be paired
51. Descibe VSEPR: purpose is to predict molecular geometries, basic assumption is that
