Chem 105 Exam 2
Bond order (BO)
# of bonding e⁻ pairs participating in resonance (# of bonds) / # of equivalent bonds (locations) BO= 1 = single bond (share 2e⁻) BO= 2 = double bond (4e⁻) BO= 3 = triple bond (6e⁻) -As BO inc, length dec, & energy inc -The higher the BO the stronger the bond & the shorter the length (share more e⁻s so it's harder to break the bond)
Formal charge
(+) and (-) symbols sometimes placed on atoms in Lewis structures -> provide estimate of charge distribution in molecules FC= # of ve⁻ for atom - #e⁻ obtained from bond division
Electronegativity
-Ability of an atom in a compound to attract e⁻s to itself & away from a neighboring atom -> x= (IE₁+EA)/2 -Noble gases are exceptions since x follows trends of IE & EA
Planetary Model of the Atom (Rutherford 1911)
-All pos-charge & almost all mass is concentrated in the nucleus, imagined orbiting e- around nucleus -Explained large scattering angles due to repulsion forces due to hitting the nucleus (low chance explains infrequency) -Prevailing view of atomic structures in 1920 1) The electron follows a planetary-like orbit 2) Different e-s are in diff orbits Problem: This kind of atom wouldn't be stable, physics requires that such an atoms must collapse (due to opposite charges)
Electron spin quantum number (ms)
-Can only take on 2 values = -1/2 & +1/2 -Electron having ms= -1/2 vs +1/2 behave diff in magnetic fields -> spinning tops -2e⁻s may occupy any given orbital as long as one is spin up & the other is spin down
Principle quantum number (n)
-Can only take on positive integers -Determines the spatial extent of an orbital = determines the size of the orbital
Heisenberg uncertainty principle
-Can't know both the position and momentum of an electron at the same time due it's wave properties *impossible* -We can't assert that electrons are in defined orbits w/i atoms -> orbits aren't good descriptions of electrons in atoms
Apparent octet expansion
-Central atoms having Z>12, steric # >4 may exceed octet *Steric #: the sum of the # of bonded outer atoms & # of lone pairs that reside on the central atom* -Follow octet rule before minimizing FC -Octet may NOT be exceeded for main group central atoms
Screening/shielding
-Determine how much positive nuclear charge various electrons are exposed to -The orbitals in multi e⁻ atoms are like those in the H atom BUT are contracted/smaller by their exposure to additional positive nuclear charge More protons = more pull on e⁻s EX: H atom (Z=1): electron feels one pull positive charge so Z*=Z due to lack of screening
Quantum numbers
-Each electron in an atom has a unique set of 4 quantum #s (q. no.) 1) 3 q. no. describe the e⁻ wave (orbital) 2) 4th q. no. describes the e⁻ spin: not related to standing wave
Pauli exclusion principle
-Each e⁻ must have its own unique set of 4 quantum #s -Violation is forbidden/impossible** -Can share first 3 q. no.s but not ms if they want to share the same orbital
Aufbau principle
-Electron configuration or orbital occupancy of the elements may be built up by successive additions of e⁻s to available orbitals using the 3 prior rules
Quantum mechanical model
-Electron was shown to exhibit wave-particle duality -Some experiments detect electrons as particles, others detect them as waves -> always have both properties BUT MAY ONLY BE ABLE TO SEE ONE IN AN EXPERIMENT -Wave-particle duality is observed only for very small objects ie photons, e-, neutrons, & some atoms **Exists for all objects but can only be observed in small objects**
Electron waves
-Electron waves in atoms are standing waves that are "confined" around a single point = the nucleus *at a fixed time* 1) Spherical 2) A = - or + 3) Pure crest or trough wave 4) Don't have hard boundaries so they can spread out in space 5) Aren't completely confined in the atom -> the spheres drawn indicate where the e⁻s are likely to found 6) Oscillate b/t pos and neg amplitude in time 7) Every e⁻ in an atom is described by a wave like this -> they're called atomic orbitals -> meaning "related to orbits" not that e⁻s follow orbits
Hund's rule
-Electronic states/configurations w/ higher total spin (S) are more stable where S=Isum of msI *Higher spin = lower energy* -When e⁻s are placed in orbitals of equal energy they prefer to have same spin (remain unpaired or spin aligned) & diff ml -Violations produce excited states (also violates ground state rule)
Electron affinity (EA)
-Energy liberated/released when an e⁻ is added to an atom to form an anion (neg charged ion) *generally reported as pos value* -Generally inc across a row due to inc Z* (attracting free e⁻s causes release of E) -Generally dec down a column due to inc n (E release progressively dec w/ inc n) -Most elements have pos EA EXCEPTIONS: noble gases & group 2 (alkaline earth metals) -> Noble gases have neg EA bc by gaining an e⁻ they lose noble gas configuration (e⁻ is added to the n+1 outer shell which is even farther from the nucleus) -Halogens have largest EA bc for those elements by gaining 1e⁻ they achieve noble gas configuration; have largest Z* for same n for a given row
Energy level diagrams (H atom)
-For 1e⁻ atoms: energy depends only on n -> atomic orbitals w/ the same n are degenerate -When the e⁻ is in the 1s orbital (lowest energy) = results in the ground state of the atom/ion -When the e⁻ is in any other orbital above 1s = excited state (less stable) then it'll emit a photon to back to the ground state
Ground state configuration for transition metals
-For neutral TM atoms the ground state config. suggested by the periodic table (ns² (n-1)dⁿ) is sometimes incorrect BEST KNOWN EXCEPTIONS: Where half full (d⁵) or full (d^10) subshells can be created = often have extra stability (preferred) Group 6 (Cr & Mo) & Group 11 (Cu, Ag, Au) exceptions *Tungsten NOT an exception even though it's in Group 6* ns & (n-1)d orbitals are close in energy -> in neutral TM atoms the ns orbital is filled first; in TM cations (n-1)d filled first **
Electron waves in atoms
-Interact w/ each other: make chem bonds, determine molecule structure, determine chem reactivity -Determine organization of the periodic table & its trends -3D standing waves = electron waves
Resonance structures
-Invoked when a single lewis structure doesn't exist -Equivalent resonant forms = same # of FC & multiple bonds b/t same kinds of atoms -Non-equivalent: pushing e⁻s from diff atoms (composite usually can't be drawn) NOTE: only e⁻s may be moved around NOT atoms -Real structure is an avg of them -> composition representation (not a lewis structure)
Octet rule
-Main group atoms in 2nd period & below prefer to be surrounded by 8 valence electrons -> reach noble gas configuration/full shell -Shared (bonding) e⁻s OR unshared (lone-pair_ e⁻s H atoms follow the duet rule (only need 2 e⁻ to achieve noble gas config)
Angular momentum quantum number (l)
-Max values: n-1 s=0, sphere p=1, tangent spheres d=2, 4 tear drops f=3 -Shape of orbital depends on (l) *doesn't tell you e⁻ orbit/location*
Magnetic quantum number (ml)
-Only integers from -l to +l -The # of ml values determines the # of orbitals in a set -Orientation of orbitals are related to ml values -Orbitals in the set are identical in shape but diff orientation -> energetically degenerate = have the same energy EX: l=0, ml=0 (1 s orbital); l=1, ml=-1,0,1 (3 p orbitals) *Cartesian coordinate system*
Lewis dot structures
-Only valence e⁻s participate in bonding -Predict distribution of ve⁻s (not always correctly) -Do NOT predict the shape of the molecule -Don't tell us anything abt the orbitals participating in bonding -Pre-quantum mechanical model
Energy level diagrams for multi e- atoms
-Orbital energies depend on both n & l 1s<2s<3s... (diff n same l= lower energy than) ns<np<nd<nf (same n, diff l) -All orbitals in a set of p, d, or f orbitals are degenerate -These energy levels & their sequence determine orbital filling patterns & the arrangement of the periodic table
Ground state rule
-Orbital filling proceeds to give the lowest total energy of the atom (ie ground state) -> to do this: must fill orbitals from bottom to top -Violations produce an excited state (violation is allowed)
2 types of nodes
-Orbitals may contain 2 types of nodes (where A=0) 1) Radial nodes (spherical): inc size of the orbital w/ inc n -l is the same = shape is the same -Phase changes when crossing a node -The # of radial nodes in any orbital n-l-1 2) Angular nodes (planar): A=0 in the nodal plane b/t lobes -Phase change when crossing node -Nodal plane = containing nucleus Formula: # of AN in any orbital = l
Plum pudding model of the atom (Thomson 1887)
-Proposed positive charge to be smoothly distributed throughout the atom -Discovered electrons (all elements have identical negatively charged particles)
Traveling waves (eg ocean waves)
-Snapshot in time A-max = wave crest A-min = trough -Wave propagates in both space & time -Crests & troughs don't remain in fixed positions -> they travel = unconfined in space
Standing wave (eg violin strings)
-Snapshot still; different energy levels (each standing wave is a diff excitation of string) -Has boundary conditions where amplitude (A)=0 not to be confused w/ nodes -Nodes: the string remains motionless -> A=0 at all times; can only have integer # of nodes -For a standing wave w/ non-zero nodes: nodes are equally spaced Signs of amplitute change when crossing a node
First ionization energy (IE₁)
-The energy required to remove the 1st outermost electron from a neutral atom in gas phase to produce a cation -Increases across a row due to inc Z* -Dec down a column due to n & valence e⁻ higher in E X (g) -> X⁺(g) + e⁻
Effective nuclear charge (Z*)
-The net/effective positive nuclear charge experienced by any given electron in an atom/ion EX: Li (Z=3) -Occasionally the 2s e⁻s penetrates the 1s² core thus Z*2s > 1 -1s e⁻s partially screen each other (and occasionally by the penetrated 2s e⁻_ thus Z* 1s < Z -2s e⁻ is shielded by the 1s e⁻s -> Z* 2s < Z -The core e⁻s are contracted relative to the valence e⁻s bc of the greater Z* The energy (E1s) will be less than (more neg or less E) than E2s In general: Z*ns > Z*np > Z*nd if same principle principle q. no. thus E ns<E np<E nd
Atomic radii (r)
-The sizes of atoms/ions are determined by the volume occupied by their e⁻s (esp valence e⁻s) -As Z* inc, e⁻s are pulled closer to nucleus = r dec -As additional core shells are added (inc n) the valence e⁻s get further from nucleus = inc r
Losing an electron
-Usually the last e⁻ added is the first to be removed to form cations (pos charged ions) HOWEVER: for TM ions the *ns e⁻ are removed first, before (n-1)d e⁻s*
Periodic trends
-Z* increases from left to right bc a proton is added to the nucleus in each step causing Z to inc -For each row one additional e⁻ is added to the valence shell The left to right trends are primarily due to inc Z* -Z* are inc slightly top to bottom but doesn't dictate the trends down the column Top-down trends are due to inc energy & inc radial distances of the valence e⁻s as the core shells are progressively filled -> inc n used as proxy (higher E & father from nucleus)
Exceptions to octet
1) H only requires 2e⁻ for noble gas config (duet rule) 2) B & Be are e⁻ deficient atoms 3) Some molecules appear to have more than an octet around central atom = apparent octet expansion 4) Odd #ve⁻ molecules EX: NO
Preferred lewis structures
1) Obey octet if possible 2) Minimize FC (sum of absolute values of FC) while preserving octet 3) Place pos FC on electropositive atoms & neg FC on more electronegative atoms as much as possible
Rutherford alpha-scattering experiment
Alpha particles =He⁺² -Alpha particles can penetrate objects -> shot at gold foil w/ a detector OBSERVED: most particles passed through w/ some deviation of angle BUT some had angles >90° w/ some coming back to the source -Atoms as proposed by Thomson wouldn't produce large scattering angle
Partial charge
Sum of charges on each equivalent atom/ # of equivalent atoms
2nd ionization energy (IE₂)
X(g) -> X⁺(g) -> X²⁺(g) IE₂ is not E required to remove 2e⁻ all together IE₂> IE₁ -Always greater than IE₁ for a given element bc the removal of the first e⁻ inc Z* (less screening) providing one fewer screening e⁻s
Ionic radii (r)
anion>atom>cation for a given element (same Z) NOTE: in anion & cation formation Z doesn't change (# of protons remains the same) BUT # of e⁻ changes thus Z* changes -> Z* for outermost e⁻ inc for cations & dec for anions (more shielding, less attraction) Isoelectronic: same # of electrons
