RECAP Schrodinger’s Standing Waves Louis De Broglie developed a theory that matter can have...
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Transcript of RECAP Schrodinger’s Standing Waves Louis De Broglie developed a theory that matter can have...
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 nucleuso 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 TIMEo 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
Summary of Quantum Numbers & Restrictions
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, ml: the orientation of the sublevel (-l to +l)
Spin quantum number, ms: the electron spin (-1/2 to +1/2)
Energy shell Orbital shape Orbital orientation
Electron Spin
1 0 0 +1/2,-1/2
201
0-1,0,+1
+1/2,-1/2+1/2,-1/2
3012
0-1,0,+1
-2,-1,0,+1,+2
+1/2,-1/2+1/2,-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
2s2p
4s3p
3d
4p
5s4d
5p
6s
E
Pictorial representation of
electron distribution in
orbitals
Aufbau principle – e- occupy the lowest energy
orbital availablep. 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 ml = 0 ms = -½
Creating Energy-Level Diagrams• 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
Energy Level Diagrams
1s
3s
2s2p
4s3p
3d
4p
5s4d
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
2s2p
4s3p
3d
4p
5s4d
5p
6s
E
Cations - Remove e- from sublevel with highest value of n.
Energy Level Diagrams
1s
3s
2s2p
4s3p
3d
4p
5s4d
5p
6s
E
Explaining the Periodic Table
Main Group metals are “S block”
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
6p5d
4f
5p4d
3d
2e-
2e-
6e-
2e-
6e-
2e-10e
-
6e-
2e-10e
-
6e-
2e-14e
-
10e-
6e-
2e-
8e-
8e-
18e-
18e-
32e-
Period 1
Period 2
Period 3
Period 4
Period 5
Period 6
Group1,2 13-18 3-12 La & Ac series
Electron configurations• Lists e- location from low to high energy
in the following format
iron atom: 1s2 2s2 2p6 3s2 3p6 4s2 3d6
Shorthand e- configurations
becomes . . .
Cl: 1s2 2s2 2p6 3s2 3p5
• e- config. written only for the outer shell electrons• Inner shell e- expressed as a noble gas. So . . .
Cl: 1s2 2s2 2p6 3s2 3p5
becomes . . .
Cl: [Ne] 3s2 3p5
1s 2s 2p 3s 3p
Cl (z = 17)
Shorthand: [noble gas]
•Completely filled sub-levels are more stable than half filled orbitals which are more stable than partially filled orbitals.
> >
Anomalous E Configurations
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-Electrostatic repulsion
e-South
Magnetic fieldMagnetic attractione-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] 4s2 3d7
Co: [Ar]
Co2+: [Ar]
Co3+: [Ar]
- Filled 4s and half filled 3d
- Half filled 4s & 3d less stable than Co2+ so less common ion
Explaining the multivalent ions
•Some materials exhibit strong magnetic properties naturally these are referred to as being ferromagnetic.
•As moving charges have magnetic fields so too do spinning electrons. Those spinning in one direction would have one magnetic polarity as those spinning in the opposite direction the opposing magnetic pole.
•For example – clockwise – south pole and counter clockwise –north pole
•Those atoms that have a number of similarly spinning electrons would have similar magnetic fields.
•If these “magnetic atoms” were free to align themselves with neighbouring atoms of similar characteristics they would create regions of magnetism in the material called domains.
•The arrangement of these domains in a material results in a magnet.
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.
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.