RECAP

Schrodinger’s Standing Waves Louis De Broglie developed a theory that matter can have wave-like properties Schrodinger extended this theory to electrons bound to a nucleus o Postulated that electrons resembled a standing wave Certain orbitals exist at whole wavelengths of electron vibrations

Orbitals - Redefined Orbital: region around the nucleus where there is a high probability of finding an electron o As per wave model of Schrodinger – because things are vibrating

Heisenberg Uncertainty Principle

Heisenberg studied statistics and developed matrix algebra Developed a statistical approach to explaining how electrons works and realized… IT IS IMPOSSIBLE TO KNOW THE EXACT POSITION AND SPEED OF ELECTRON AT A GIVEN TIME o At best, we can describe the probability of finding it at a specific place

Wave functions: the mathematical probability of finding an electron in a certain region of space Wave functions give us: Electron probability densities: the probability of finding an electron at a given location, derived from wave equations

Quantum Numbers Quantum Numbers: numbers that describe the quantum mechanical properties (energies) of orbitals o From the solutions to Schrodinger’s equation The most stable energy states is called the ground state

Principal quantum number, n: the main electron energy level or shell (n ) Secondary quantum number, l: the electron sublevels or subshells (0 to n-1) Magnetic quantum number, m l : the orientation of the sublevel (-l to +l) Spin quantum number, m s : the electron spin (-1/2 to +1/2) Energy shellOrbital shape Orbital orientation Electron Spin 100+1/2,-1/ ,0,+1 +1/2,-1/ ,0,+1 -2,-1,0,+1,+2 +1/2,-1/2

Electron Configurations and Energy Level Diagrams The four quantum numbers tell us about the energies of electrons in each atom Unless otherwise stated were are talking about ground state energies

Energy Level Diagrams 1s 3s 2s 2p 4s 3p 3d 4p 5s 4d 5p 6s E Pictorial representation of electron distribution in orbitals Aufbau principle – e - occupy the lowest energy orbital available p. 188 in text Aufbau principle – e - occupy the lowest energy orbital available p. 188 in text Hund’s rule – e - half-fill each orbital in a sublevel before pairing up Pauli exclusion principle – max 2 e - per orbital (spin up and spin down) n = 1 l = 0 m l = 0 m s = - ½

Pauli exclusion principle – o no two electrons in an atom may have the same four quantum numbers o no two electrons in the same orbital may have the same spin o only two electrons with opposite spins may occupy an orbital aufbau principle – (German for “building up’) o each electron is added to the lowest available energy orbital Hund’s rule – o one electron is placed in each orbital at the same energy level before the second electron is placed

1s 2s2p 3s3p O (z = 8) 1s 2s2p 3s3p P (z = 15) 1s 2s2p 3s3p Ar (z = 18)

Energy Level Diagrams 1s 3s 2s 2p 4s 3p 3d 4p 5s 4d 5p 6s E Anions - Add e - to lowest energy sublevel available.

Cations - Remove e - from sublevel with highest value of n. Energy Level Diagrams 1s 3s 2s 2p 4s 3p 3d 4p 5s 4d 5p 6s E

Cations - Remove e - from sublevel with highest value of n. Energy Level Diagrams 1s 3s 2s 2p 4s 3p 3d 4p 5s 4d 5p 6s E

Electron Configuration & The Periodic Table

s-block 2 groups p-block 6 groups d-block 10 groups f-block 14 groups

1s 2s 3s 4s 5s 6s 2p 3p 4p 6p 5d 4f 5p 4d 3d 2e - 6e - 2e - 6e - 2e - 10e - 6e - 2e - 10e - 6e - 2e - 14e - 10e - 6e - 2e - 8e - 18e - 32e - Period 1 Period 2 Period 3 Period 4 Period 5 Period 6 Group 1, La & Ac series

Lists e - location from low to high energy in the following format iron atom: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 6

becomes... Cl: 1s 2 2s 2 2p 6 3s 2 3p 5 e - config. written only for the outer shell electrons Inner shell e - expressed as a noble gas. So... Cl: 1s 2 2s 2 2p 6 3s 2 3p 5 becomes... Cl: [Ne] 3s 2 3p 5 1s 2s2p 3s3p Cl (z = 17) Shorthand: [noble gas]

Anomalous E Configurations

Anomalies in Configurations n=3 n=4 spd

The configuration of Copper

Explaining the multivalent ions As orbitals are being filled there are varying levels of stability due to the interaction of forces or electrostatic repulsion between electrons and force associates with the magnetic field due to the electron’s spin. Filled orbitals are most stable because the electrostatic repulsion is balanced against the magnetic attraction. e-e- e-e- Electrostatic repulsion e-e- South Magnetic field Magnetic attraction e-e- North Magnetic field The configuration of stable ions can also be predicted by electron configuration

Orbitals that are completely filled, like Noble gases, have the most stable structure due to the balanced forces between electrostatic repulsion and magnetic attraction. Transitional metals often have partially filled d-orbitals and complete s- orbitals with similar energy levels. Electrons are lost to achieve the best combination of stability. Best stability – completely filled orbitals (2 electrons/orbital) Electrostatic repulsion and Magnetic attraction Next best stability – half filled orbitals (1 electron/orbital) Electrostatic repulsion with minimal crowding Co: [Ar] 4s 2 3d 7 Co: [Ar] Co 2+ : [Ar] Co 3+ : [Ar] - Filled 4s and half filled 3d - Half filled 4s & 3d less stable than Co 2+ so less common ion Explaining the multivalent ions

Iron, nickel and cobalt are such naturally occurring atoms. Iron, nickel and cobalt are small enough atoms that they can realign themselves due to the magnetic properties of their surrounding and thereby create domains. There are other such atoms that have similar electron configurations but limited ability to migrate. Hence, they are reduced in their magnetic properties.

Ferromagnetic – materials with strong magnetic properties. Their presence increases a magnetic field substantially. Paramagnetic – materials with weak magnetic properties. Their presence only slightly strengthens a magnetic field. Diamagnetic – materials that have reduced magnetic properties. Their presence weakens a magnetic field.

Shape of orbitals The diagram we used to represent oxygen is; 8 Protons O

Shape of orbitals The diagram we might currently use to represent oxygen is;

p. 191 #3alt,4 p. 197 #5 p. 194 #6-11 p. 197 #1,2,6-10,13

Confidence building questions 1.Write out the shorthand notation for the electron configuration of B. 2.Write out the shorthand notation for the electron configuration of Cl. 3.Write out the shorthand notation for the electron configuration of F. 4.Write out the shorthand notation for the electron configuration of Ca. 5.Write out the shorthand notation for the electron configuration of Kr. 6.Write out the shorthand notation for the electron configuration of O2-. Notice that this is an anion! 7.Write out the shorthand notation for the electron configuration of Na+. Notice that this is a cation! 8.Why are Groups 1 and 2 referred to as the s-block of the periodic table? 9.Why are Groups 3 through 12 referred to as the d-block of the periodic table? 10.Using what you now know about electron configurations explain the notion that elements in the same column in the periodic table have similar chemical and physical properties.