6. Carbon The first 4 electrons will be 1s22s2 which leaves 2 electrons to be placed in the available 2p orbitals. Some possibilities are: 2px 2py 2pz

6. Carbon The first 4 electrons will be 1s22s2 which leaves 2 electrons to be placed in the available 2p orbitals. Some possibilities are: 2px 2py 2pz 2px 2py 2pz

6. Carbon The first 4 electrons will be 1s22s2 which leaves 2 electrons to be placed in the available 2p orbitals. Some possibilities are: 2px 2py 2pz 2px 2py 2pz 2px 2py 2pz

6. Carbon The first 4 electrons will be 1s22s2 which leaves 2 electrons to be placed in the available 2p orbitals. Some possibilities are: 2px 2py 2pz 2px 2py 2pz 2px 2py 2pz 2px 2py 2pz 2px 2py 2pz 2px 2py 2pz

6. Carbon The first 4 electrons will be 1s22s2 which leaves 2 electrons to be placed in the available 2p orbitals. Some possibilities are: 2px 2py 2pz 2px 2py 2pz 2px 2py 2pz 2px 2py 2pz 2px 2py 2pz 2px 2py 2pz 2px 2py 2pz 2px2py 2pz 2px 2py 2pz

Choices can be restricted to: 2px 2py 2pz

To predict which one of the above will give the greatest stability, requires the use of Hund’s rule.

To predict which one of the above will give the greatest stability, requires the use of Hund’s rule. Hund’s Rule: The most stable arrangement of electrons in subshells is the one with the greatest number of parallel spins.

To predict which one of the above will give the greatest stability, requires the use of Hund’s rule. Hund’s Rule: The most stable arrangement of electrons in subshells is the one with the greatest number of parallel spins. 2px 2py 2pz This is the preferred configuration.

Hence, for carbon the electronic configuration is 1s2 2s22px2py The orbital diagram is: C 1s 2s 2px 2py

7. Nitrogen: The electronic configuration is 1s22s22px2py2pz The orbital diagram is: N 1s 2s 2px 2py 2pz

7. Nitrogen: The electronic configuration is 1s22s22px2py2pz The orbital diagram is: N 1s 2s 2px 2py 2pz Note that Hund’s rule applies to get the lowest energy configuration.

8. Oxygen: The electronic configuration is 1s22s22px22py2pz The orbital diagram is: O 1s 2s 2px 2py 2pz

8. Oxygen: The electronic configuration is 1s22s22px22py2pz The orbital diagram is: O 1s 2s 2px 2py 2pz Hund’s rule applies to get the lowest energy configuration.

8. Oxygen: The electronic configuration is 1s22s22px22py2pz The orbital diagram is: O 1s 2s 2px 2py 2pz Hund’s rule applies to get the lowest energy configuration. It is known from experiment that oxygen atoms are paramagnetic.

9. Fluorine: The electronic configuration is 1s22s22px22py22pz The orbital diagram is: F 1s 2s 2px 2py 2pz

10. Neon: The electronic configuration is 1s22s22px22py22pz2 The orbital diagram is: Ne 1s 2s 2px 2py 2pz

Maximum number of electrons that can be assigned to the various orbitals

Maximum number of electrons that can be assigned to the various orbitals For each shell (with principal quantum number n) there are n subshells.

Maximum number of electrons that can be assigned to the various orbitals For each shell (with principal quantum number n) there are n subshells. For example, if n = 3, there are three subshells (three values of l) having l = 0, l = 1, and l = 2, corresponding to 3s, 3p, and 3d.

Maximum number of electrons that can be assigned to the various orbitals For each shell (with principal quantum number n) there are n subshells. For example, if n = 3, there are three subshells (three values of l) having l = 0, l = 1, and l = 2, corresponding to 3s, 3p, and 3d. Each sublevel of quantum number l contains 2 l + 1 orbitals. For example, if l = 1 then there are three p orbitals.

Finally, each orbital can take two electrons Finally, each orbital can take two electrons. Therefore the total number of electrons that may be placed in a given number of orbitals is twice the total number of available orbitals.

Finally, each orbital can take two electrons Finally, each orbital can take two electrons. Therefore the total number of electrons that may be placed in a given number of orbitals is twice the total number of available orbitals. Example: What is the maximum number of electrons that can be present in the principal level for which n = 3?

When n = 3, we have: l = 0, l = 1, l = 2 The number of orbitals for each value of l is

When n = 3, we have: l = 0, l = 1, l = 2 The number of orbitals for each value of l is value of l number of orbitals (2 l + 1)

When n = 3, we have: l = 0, l = 1, l = 2 The number of orbitals for each value of l is value of l number of orbitals (2 l + 1) 0 1

When n = 3, we have: l = 0, l = 1, l = 2 The number of orbitals for each value of l is value of l number of orbitals (2 l + 1) 0 1 1 3

When n = 3, we have: l = 0, l = 1, l = 2 The number of orbitals for each value of l is value of l number of orbitals (2 l + 1) 0 1 1 3 2 5

When n = 3, we have: l = 0, l = 1, l = 2 The number of orbitals for each value of l is value of l number of orbitals (2 l + 1) 0 1 1 3 2 5 Total number of orbitals = 1 + 3 + 5 = 9

When n = 3, we have: l = 0, l = 1, l = 2 The number of orbitals for each value of l is value of l number of orbitals (2 l + 1) 0 1 1 3 2 5 Total number of orbitals = 1 + 3 + 5 = 9 The maximum number of electrons that can fit into these orbitals is 2x9 = 18.

When n = 3, we have: l = 0, l = 1, l = 2 The number of orbitals for each value of l is value of l number of orbitals (2 l + 1) 0 1 1 3 2 5 Total number of orbitals = 1 + 3 + 5 = 9 The maximum number of electrons that can fit into these orbitals is 2x9 = 18. (number = 2n2)

Summary for electronic configurations No two electrons in the same atom have the same four quantum numbers (Pauli Exclusion Principle).

Summary for electronic configurations No two electrons in the same atom have the same four quantum numbers (Pauli Exclusion Principle). 2. Each orbital can be occupied by a maximum of two electrons. They must have opposite spins (different values of the spin quantum number ms).

Summary for electronic configurations No two electrons in the same atom have the same four quantum numbers (Pauli Exclusion Principle). 2. Each orbital can be occupied by a maximum of two electrons. They must have opposite spins (different values of the spin quantum number ms). 3. The maximum number of electrons that can occupy each energy level is determined by the number of orbitals and the Pauli Exclusion Principle.

4. Orbitals are filled in the order of increasing energy.

4. Orbitals are filled in the order of increasing energy. 5 4. Orbitals are filled in the order of increasing energy. 5. The most stable arrangement of electrons in subshells is the one which has the greatest number of parallel spins (Hund’s rule).

4. Orbitals are filled in the order of increasing energy. 5 4. Orbitals are filled in the order of increasing energy. 5. The most stable arrangement of electrons in subshells is the one which has the greatest number of parallel spins (Hund’s rule). 6. Atoms (or molecules in general) which have one or more unpaired electrons are paramagnetic. Atoms (or molecules in general) which have all of the electron spins paired are diamagnetic.

4. Orbitals are filled in the order of increasing energy. 5 4. Orbitals are filled in the order of increasing energy. 5. The most stable arrangement of electrons in subshells is the one which has the greatest number of parallel spins (Hund’s rule). 6. Atoms (or molecules in general) which have one or more unpaired electrons are paramagnetic. Atoms (or molecules in general) which have all of the electron spins paired are diamagnetic. Diamagnetic: The tendency of a species not to be attracted (or to be slightly repelled) by a magnetic field.

7. In a many-electron atom, the energy of an electron depends on both n and l. This means that, for example, the 3s orbital is filled before the 3p orbital, and usually the 4s orbital is filled before the 3d orbital.

7. In a many-electron atom, the energy of an electron depends on both n and l. This means that, for example, the 3s orbital is filled before the 3p orbital, and usually the 4s orbital is filled before the 3d orbital. 8. For electrons of the same principal quantum number, the penetrating power, or proximity to the nucleus , decreases in the order s > p > d > f This means that, for example, more energy is required to remove a 3s than a 3p electron, and so on.

Shorthand Notation for Electronic Configurations

Shorthand Notation for Electronic Configurations Potassium has 19 electrons; its electronic configuration is 1s22s22p63s23p64s1

Shorthand Notation for Electronic Configurations Potassium has 19 electrons; its electronic configuration is 1s22s22p63s23p64s1 The inert gas preceding potassium in the periodic table is argon, for which the electron configuration is 1s22s22p63s23p6.

Shorthand Notation for Electronic Configurations Potassium has 19 electrons; its electronic configuration is 1s22s22p63s23p64s1 The inert gas preceding potassium in the periodic table is argon, for which the electron configuration is 1s22s22p63s23p6. So the electronic configuration is [Ar]4s1

The electronic configuration of Sr is 1s22s22p63s23p63d104s24p65s2 This can be written in shorthand notation as [Kr]5s2

Electronic configurations of the transition metals

Electronic configurations of the transition metals The elements from Sc to Zn form the first transition metal period. The 3d subshell is filled for this group.

Electronic configurations of the transition metals The elements from Sc to Zn form the first transition metal period. The 3d subshell is filled for this group. From Sc [Ar]4s23d1 to V [Ar]4s23d3 the electrons are placed in the 3d subshell according to Hund’s rule.

First exception:

First exception: For Cr we would expect the electronic configuration [Ar]4s23d4.

First exception: For Cr we would expect the electronic configuration [Ar]4s23d4. From experiment it is known that the electronic configuration of Cr is [Ar]4s13d5.

First exception: For Cr we would expect the electronic configuration [Ar]4s23d4. From experiment it is known that the electronic configuration of Cr is [Ar]4s13d5. Second exception: For Cu we might expect the electronic configuration [Ar]4s23d9.

First exception: For Cr we would expect the electronic configuration [Ar]4s23d4. From experiment it is known that the electronic configuration of Cr is [Ar]4s13d5. Second exception: For Cu we might expect the electronic configuration [Ar]4s23d9. However experiments indicate the electronic configuration of Cu is [Ar]4s13d10.

The reason for these irregularities is a slight extra stability associated with the half-filled (3d5) and completely filled (3d10) subshells.

The orbital diagram for Cr is: [Ar] 4s 3d

The orbital diagram for Cu is: [Ar] 4s 3d

After Cu, filling the 4s and 4p orbitals proceeds in a straightforward manner and is complete with Kr.

Core Electrons