Electron Configuration Chemistry Notes PDF

1/16/25 The Many Electron Atom, Effective Nuclear Charge, & Electron Configurations Tro, “Chemistry: Structure and Properties”: 3.2 The periodic law and the periodic table 3.3 Electron conf...

1/16/25 The Many Electron Atom, Effective Nuclear Charge, & Electron Configurations Tro, “Chemistry: Structure and Properties”: 3.2 The periodic law and the periodic table 3.3 Electron configurations: how electrons occupy orbitals 3.4 Electron configurations, valence electrons, and the periodic table 3.5 Electron configurations and elemental properties 3.6 Periodic trends in atomic size and effective nuclear charge 3.7: “Ions: electron configurations, magnetic properties, radii, and ionization energy” 1 Tro – Chemistry: structure and proper5es 2.6: The shapes of atomic orbitals Energy Levels in the Hydrogen Atom According to solu:ons of the Schrödinger equa:on, for the hydrogen atom: The energy of each shell is determined solely by: the principle quantum "n" , True for all 1-electron systems, accounBng for “Z” (or Zeff) of the same Therefore, subshells (!, ", #...) within the same shell are: degenerate , energy −'ℎ)" $! = +# n 3 = n = 2 En = energy of the orbital h = Planck’s constant Z = nuclear charge R = Rydberg’s constant n = principal quantum # U = 1 2 2 1  1/16/25 The Energy of Orbitals in non-Hydrogen Energy levels are “split” in systems with >1 electron due to: electrostatic interactions i.e. the subshells within a shell are no longer degenerate more lines = more 1) Energy increases as ____ N increases & electronic transitions e-further from the away nucleus H= 2 n= 1 3 3 Orbital Energy as a Function of Shape (Penetration Ability) l increases 2) Energy increases as ___ Penetration of electrons into the inner An orbital’s __________ shape significantly impacts its penetraBng ability shell reduces the magnitude of the shielding they experience, and increases Angular nodes reduce an electron’s: access to the nucleus/its ~ their ability to shield other "! The degree of penetraBon is significant for ___-orbitals, S but is attraction aOenuated for ____ P and ____ d -orbitals. to the nucleus higher A b higher Man than a VC 1 at ov probable to find Es Orbital m -more -- less e- likely to find e = e-have more access to close to the nucleus4 uncles > - unshielded > - more stable 4 2  1/16/25 Effective Nuclear Charge Z = nuclear charge = # protons in the nucleus Zeff = the actual magnitude of the positive charge that is “experienced” by an electron in an atom Why are these not the same thing?! Electrons are simultaneously attracted to the nucleus, but repelled by one another Shielding (or screening) = partial obstructions of nuclear charge by other electrons #"## = 2 -- ! = shielding constant Estimate for ! = # of core e- “Slater’s rules” can be used to calculate ! , and are much more specific and accurate 5 Screening (Shielding) EffecEveness Core electrons are most effective at shielding valence electrons. Electrons in a lower shell of the same subshell type are very effective at shielding. i.e. Is e are very good at 3se- shielding Electrons in the same shell but different subshell type are less effective at shielding one another. However, lower energy subshells provide some shielding to higher energy subshells i.e. 35 electrons minimally shield 3p electrons , which minimally shield sde- Electrons within in the same subshell do not shield one another. i.e. spelections do not shield 3p electrons (nor... 2 pet 6 6 3 Prac%ce: imagine an e- in a 5p orbital of a  1/16/25 total orbitals = u ? maximum no · of e = zn ? Electron Configurations An electron configuration is a list of an atom’s occupied atomic orbitals, and the number of electrons in each, as # of e-in that subshell described by the quantum numbers (+, ,, -$ , and -% ). Electrons are listed from lowest to highest energy A shorthand notation for the electronic state an atom energylever type ofsubsnell Defines the atom’s chemical and physical properties. 3 primary rules/principles for the arrangement of electrons in orbitals: Aufbau principle: place electrons in the lowest energy orbitals first i.e. "build-up" principle , LtR , top - > bottom Pauli exclusion principle: no two electrons can have the same four quantum numbers i.e. maximum # 2e- per orbital ( + 1 -1 ↑ V) of ,... Hund’s rule = the most stable arrangement of e- in degenerate orbitals is the one in which parallel spins are maximized i.e. do not pair e-until you have to , must be singly filled before being doubly filled 7 7 The Periodic Table: A Coordinate System Based on Quantum #s Orbitals in the !th period are filled from left to right, in order: ns (n-2)f (n-1) d , , , up " = 0 (#-orbitals) 2 " = 1 (%-orbitals) # n=1 1# % 6 % n=2 2# % 2%$ & = 3, " = 2 ('-orbitals) Row when n=3 3# % 10 3%$ occupa:on ' n=4 4# % 3'&' 4%$ of " -type orbitals n=5 5# % 4'&' 5%$ begins ) n=6 6# % 5'&' 6%$ n=7 7# % 6'&' 7%$ ↑ 4) &( 5) &( + 5 more 14 & = 4, " = 3 () -orbitals) 8 Row = period = “n” s-block l = 0, ml = o, (1 orbital) Max e- = 2 4  1/16/25 Writing Electron Configurations Practice: write the full electron configurations for: Orbitals in the &th period are Magnesium (Z = 12) 1522522p63s7 filled from left to right, in order: &, < & − 2 / < & − 1 0 < &1 Selenium (Z = 34) Is 2232 2p63323p64323d4p4 1522322p63533p64323d 4p65934915p66324445d3 * Tantalum (Z = 73) Period # 9 9 Why does “n” change for some blocks?! It depends on the first time that subshell is available to fill! Writing Noble Gas Electron Configurations Noble gas configuration: a simplified electron configuration! Includes [noble gas] in preceding row, followed by remaining " ! configuration The " ! up to the noble gas are called ______________ core and don’t participate in bonding The " ! outside of the noble has core are generally _______________ valence e- , and do participate in bonding For main group elements, valence 1 ! = outermost # and % 1 ! Is and p block) For '-block elements, valence 1 ! = outermost # and ' 1 ! Electron Configuration Noble Gas Configuration [Ne] 3523p24 valence Silicon: 1s2 2s2 2p6 3s2 3p2 4323d104p24 valence Germanium: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p2 [Ar] Nickel: 1s2 2s2 2p6 3s2 3p6 4s2 3d8 [Ar]43239810 valence e Prac%ce: write the noble gas configuraNon for: [Xe]6g24 + 145d3 5 valence e- Tantalum (Z = 73) 10 10 5  1/16/25 Common Exceptions in Electron Configurations Exceptions in electron configurations are common in the transition metals and lanthanides/actinides Apparent results: “special stability of half-filled subshells” Due to similarity of orbital energies for 4s/3d , 5s/4d, etc. Minimizes repulsions; “maximizes parallel spins” > hund's me Also due to shielding (Slater’s rules) Exceptions to recognize: Cr Mo , , Cu Ag , An , Chromium expected: [Ar]4s2 3d4 more reality: (Ar] 45'3d5 -stable as orbitals are singly filled Copper expected: [Ar]4s2 3d9 reality: [Ar]4s'3d10 11 11 Increasing metal ')** by ionization (or by partial removal of electrons onto the ligands) has a greater effect on stabilizing the energy of the &( orbital relative to the (& + 1)- orbital because nuclear penetration already exposes the --electrons to a substantial amount of nuclear charge. Orbital Diagram NotaEon An orbital diagram is a graphical representation of " 2 configurations Element #.2. 2 configuration orbital diagram Li 3 1s2 2s1 O 8 1s2 2s2 2p4 P 15 Is " Is 2p63523p3 [Ne] 3523p3 Fe 26 1522322p63323p64323d6 & IV 3p I …how many unpaired e- are in each of these atoms? 12 12 An orbital diagram is a graphical representation of e- configurations Each orbital is represented by a line or a box Must include all orbitals in a subshell, even if unoccupied! 6 E- are represented by “up” or “down” pointing arrows  1/16/25 Describing and Utilizing the Periodic Table Periods = rows Main group elements tend to gain, lose or share Groups (families) = columns the number of valence electrons needed to achieve the same number of electrons as the Analogous valence e- configura@ons à nearest noble gas similar periodic properBes Common charges can be predicted by group # Some families are referred to by common names Groups 1A-3A charges = group # (cation) I aroun Groups 5A-7A charges = group #- Olanion) 3+ 03 group 7 it - 2 - 1- 0 2+ group 13 ↑ > 13 [7ns np5 ↑ standard & halogen configuration Electron Configurations for Main Group Ions Ion = a non-neutral atom; an atom that has lost or gained valence " 2 Ca%on = a positive ion , where: # "3 & #. 2 Anion = a negative ion , where: # "3 L #.2 Wri:ng electron configura:ons for main group (s- and p-block) ions: Write the e- configuraNon of the neutral element Ca@on: remove "! from the highest energy valence orbital(s) Anion: add "! to the lowest energy available orbital(s) Atom: Ion: ? Na: 1s22s22p63s /1 Na+ : Is ? Is 2 p6 "isoelectronic same of e F- : = F: 1s22s22p5 Is22522pb 14 14 7  1/16/25 Electron Configurations for Transition Metal Ions For transition metal (d-block) ions, write the. 2 Atom: Ion: configuration of the neutral element. Then, for a… Zn: [Ar] 4s/2 3d10 Zn2+ [Ar] 3 do ! Cation: remove " from highest "n" orbital first (t) Zn2- : [Ar] 4523d104p 7 Anion: add " ! to lowest "e"available orbitals Pt: [Xe] 6s2 4f14 5d6 Pt+ : [xe] 65"4f14598 () first Pt4+ [Xe]4914596 Why?! In the ground state (neutral atom, oxidation state = 0), the (&). subshell is lower in energy than the (& − 1)0 subshell onlyADD Why electrons fill the “4#” before the “3'” For any transition metal ion (oxidation state > 0) , the (&). subshell is higher in energy than the (& − 1)0 subshell Why electrons are removed from the “4#” before the “3'” 15 15 a metal ion with a + 2 charge has the e-config. [Ar] 3d10 what is the metal ? &45 = zn 8