Atomic Structure and Periodic Trends - Chemistry

• Bohr postulated that each electron occupied only certain orbits and contains specific amounts of energies
• The further the orbit, the more energy contained in the electron
• Electrons drop from a higher-energy orbit to a lower-energy orbit. This causes a certain amount of energy in the form of light to be emitted (called a photon)
• A larger transition means that more energy is released
• Each transition corresponds to one line on the emission spectra
• Different elements have different emission spectra because they contain different numbers of electrons
• As the gaps between the lines on the emission spectra get smaller, the gaps between the energy levels also gets smaller
• Electrons do not move around the nucleus in circular orbits (we do not know about the detailed pathway of the electron). We are able to define the region around the nucleus where the electron might be found
• Atomic orbital = area around the nucleus where there is a high possibility of finding an electron
• n

  • principle quantum number
  • determines the size of the orbital
  • any whole number except 0 (from 1,2,3 and upwards)
• L

  • angular momentum number
  • determines the shape of the orbital
  • any whole number between 0 and n-1 inclusive
• m

  • magnetic quantum number
  • determines the orientation of the orbital
  • any whole number between -L and +L
• s

  • spin quantum number
  • determines the e- spin
  • either -1/2 or +1/2
• s-subshell-> sharp -> L = 0
• p-subshell-> principal -> L = 1
• d-subshell -> diffuse -> L=2
• f-subshell -> fundamental -> L=3
• If n=1, L=0 so only s-subshell is occupied
• If n=2, L=1 so s and p subshells are occupied
• If n=3, L=2 so s, p, and d subshells are occupied
• m effectively tells us how many orbitals make up a subshell and how many orbitals are in each shell
• s-subshell

  • L=0 so m=0 -> 1 number written for m so 1 orbital
  • L=1 so m=-1,0,+1 -> 3 numbers written for m so 3 orbitals
  • L=2 so m=-2,-1,0,+1,+2 -> 5 numbers written for m so 5 orbitals
• • n=1, L=0 -> 1 shell, 1 subshell
  • n=2, L=0 or 1 -> 1 shell, 2 subshells
  • n=3, L=0, 1, or 2 -> 1 subshell, 3 subshells
• Each orbital can only hold up to 2 electrons
• Each electron has a -1 charge so placing 2 electrons in close proximity to each other will cause e- - e- repulsions
• Electrons have a property called "spin"

  • The spin is described by the spin quantum number, s
  • It can have a value of -1/2 or +1/2
  • 2 electrons with opposing spins repel each other less than 2 electrons with the same spin
  • This explains why each orbital can hold a maximum of 2 electrons
• • This electron configuration chart covers all the elements in the periodic table
  • Example = H: 1s^1
  • (first) "1" represents the number of shells (n), the "s" represents the number of subshells (L), the superscript "1" represents the number of electrons

  
• Write the complete version of electron configuration for F and Ca

  • F: 1s^2 2s^2 2p^5
  • Ca: 1s^2 2s^2 2p^6 3s^2 3p^6 4s^2
• Write the condensed version of electron configuration for F and Ca

  • F: [He] 2s^2 2p^5
  • Ca: [Ar] 4s^2
• There are 2 exceptions in terms of electron configurations:

  • Cr: [Ar] 3d^5 4s^1
  • Cu: [Ar] 3d^10 4s^1
• Aufbau principle = states that all lower energy shells must be filled before starting the next subshell
• Distribution of electrons into different atomic orbitals follows the Aufbau Principle, Hund’s Rule, and Pauli Exclusion Principle.
• Pauli exclusion principle = states that no 2 electrons in the same orbital can have the same spin (i.e. every electron must have a unique combination of values for n, L, m and s)
• Hund's rule = each orbital must be half-filled before spin pairing can occur (doubling up)
• Each subshell contains orbitals and these orbitals have different shapes depending upon the type of subshell
• S-subshell -> 1 orbital

  
• P-subshell -> 3 orbitals

  A) px

  B) py

  C) pz

  
• D-subshell -> 5 orbitals

  A) dz^2

  B) dx^2 - y^2

  C) dxz

  D) dxy

  E) dyz

  
• Photoelectron Spectroscopy (PES):

  • Uses high energy radiation/light (x-ray or ultraviolet) enough to eject some electrons from the atom
  • It provides information about ionization energies of electrons in elements (i.e. how much energy is required to strip away an electron)
  • A detector is able to measure the number of electrons ejected from a material's surface
  • Units: MJ/mol (M = 10^6 J)
  • Larger binding energy (stronger force so more input of energy) = electrons are closer to the nucleus
• Using the information about ionization energies from PES, we can describe the:

  • The identity of a sample
  • The electronic configuration