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Chemistry ELECTRONIC STRUCTURE Electron Configurations. DO NOW: what are the four quantum numbers? Calculators not required Periodic tables not required. Quantum Numbers and electron placement. From last time… No two electrons may be doing the same two things: Pauli exclusion principle
ChemistryELECTRONIC STRUCTUREElectron Configurations DO NOW: what are the four quantum numbers? Calculators not required Periodic tables not required
Quantum Numbers and electron placement • From last time… • No two electrons may be doing the same two things: Pauli exclusion principle • Each orbital can hold two electrons • s subshell: ____ electrons • p subshell: ____ electrons • d subshell: ____ electrons • f subshell: ____ electrons • So, how do we determine which electrons go in which orbitals? • Works best if we think of orbitals as rooms in a hotel…
AUFBAU PRINCIPLE • The Aufbauprinciple states that electrons will enter orbitals in order from lowest energy to highest energy. • If you are booking a room in a hotel, you would prefer to spend less money for your room. • The “atomic currency” is energy. Less energy to spend is preferred.
Orbital Splitting • The Aufbau principle is necessary because the orbitals split in energy when there are many electrons repelling against other in the atom. • This makes some orbitals in an energy level higher in energy than others. • Or, some rooms on the same floor cost more to book.
Atomic “price schedule” • To the right is the energy “price schedule” of the atomic orbitals. • To continue our analogy… • Energy level = floor # • Sublevel = type of room (suite, penthouse, deluxe, fantastic) • Orbital = single room • Spin = bed in room • 2 beds per room
Predicting the price schedule • The Aufbau principle has a set pattern. • Following and re-creating the diagram to the right will help you memorize the “price schedule” and which orbitals are lower in energy. • 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p
Orbital BOX Diagrams • Each box in the diagram represents one orbital. • Half-arrows represent the electrons. • The direction of the arrow represents the relative spin of the electron.
Hund’s Rule • The lowest energy is attained when the number of electrons with the same spin is maximized. • Make all orbitals in a subshell (s, p, d, etc.) half-filled before adding the opposite spin electrons. • Half-filled subshells are especially stable. • In other words… • Electrons will always pick the lowest price room, but will prefer to have single rooms until they are forced to double up.
Drawing Orbital Box Diagrams • Follow the atomic “price schedule” for oxygen:
Magnetism • Magnetism results from unpaired electrons in the ground state (as opposed to the excited state, like we do in the hydrogen gas light). • When there are multiple unpaired electrons (like in iron), this results in a ferromagnet, or permanent magnet, or a paramagnet, or temporary magnet.
Practice • Draw the full orbital box diagram for iron. • See if you can tell whether or not iron is magnetic.
Anatomy of an Electron Configuration • An electron configuration gives the orbital placement of every electron in an atom. • This is one piece of an electron configuration.
Electron Configurations • Each component consists of • A number denoting the energy level;
Electron Configurations • This shows the distribution of all electrons in an atom. • Each component consists of • A number denoting the energy level, • A letter denoting the type of orbital,
Electron Configurations • This shows the distribution of all electrons in an atom. • Each component consists of • A number denoting the energy level, • A letter denoting the type of orbital, • A superscript denoting the number of electrons in those orbitals.
Sample Electron Configuration • Remember, we can only fit so many electrons in each type of orbital, and that each energy level only has certain subshells: • n =1 only has s 2 electrons in s • n = 2 can have s and p 6 in p • n = 3 has s, p, and d 10 in d • n = 4 has s, p, d, f 14 in f • You keep putting electrons in orbitals from low to high energy until you run out. Example: Neon (10 electrons) • 1s22s22p6 (total of 10 electrons)
Periodic Table • We fill orbitals in increasing order of energy. • Different blocks on the periodic table (shaded in different colors here) correspond to different types of orbitals.
Practice • Give the ground-state electron configuration for sodium. • Note: ground-state is when every electron is at its lowest energy orbital. If an electron is at a higher energy orbital, we say it is in an excited state, because energy had to be put in to make it go up.
Shortened Electron Configurations • As a shortcut, we can use a noble gas (He, Ne, Kr, etc..) as a starting point for the electron configuration. • Sodium’s full ground-state electron configuration: • 1s22s22p63s1 • Neon: 1s22s22p6 • Shortened ground-state electron configuration: • [Ne]3s1
Practice • Give the shortened (noble gas) ground-state electron configurations for uranium.
Exception: half-filled and filled d orbitals • When a subshell is half-filled or completely filled, it lowers the overall energy of the atom. • In other words, this is favored. • This usually doesn’t effect electron configurations, except in the d and f subshells. • Example: Copper (29 electrons) • Normal electron configuration: 1s22s22p63s23p64s23d9 • Added stability if 3d is full (electron gets promoted to 3d from 4s) • 1s22s22p63s23p64s13d10
Exception: half-filled and filled d orbitals • The same happens with half-filled d orbitals: • Example : Chromium • 1s22s22p63s23p64s23d4 not as stable as a half-full d-subshell • Becomes: 1s22s22p63s23p64s13d5 • In summary: • Configurations that end in s2d9 become s1d10. • Configurations that end in s2d4 become s1d5. • This should also be reflected in orbital box diagrams.