Modern Atomic Theory: How are an atom’s electrons configured?

Light is an electromagnetic wave

Speed = Frequency X Wavelength Since speed is a constant: Since speed is a constant: Low frequency waves have long wavelengths High frequency waves have short wavelengths Energy is proportional to frequency so Energy is proportional to frequency so Low frequency waves have low energy High frequency waves have high energy

Particle or wave? Light can behave as a wave in some instances, and as a particle, which is known as a photon, in other instances. We call this the wave-particle duality of light. Up to this point, we have assumed that electrons are particles. But electrons also have properties that can only be explained by assuming that they are waves. They also have this dual nature.

Line emission spectra could not be explained by Rutherford’s atomic model. In 1913, Niels Bohr showed that these spectra could be explained by assuming that an electron could only exist in certain energy states. In 1913, Niels Bohr showed that these spectra could be explained by assuming that an electron could only exist in certain energy states. He called the lowest energy state the ground state. If an electron acquires additional energy, it may move into an exited state. He called the lowest energy state the ground state. If an electron acquires additional energy, it may move into an exited state. When a photon of a specific energy is absorbed, the electron moves into a higher energy level. When a photon of a specific energy is absorbed, the electron moves into a higher energy level. When the electron moves from a higher energy level into a lower energy level, a photon is released. When the electron moves from a higher energy level into a lower energy level, a photon is released.

The only possible energy levels are defined by an equation: Energy of electron = X 10 J, Energy of electron = X J, n 2 n 2 where n is a positive whole number. This number n is called the principal quantum number. When energies can only adopt certain values, the energy is said to be quantized.

Modern Quantum Theory Electrons are found in orbitals, which are sometimes represented as clouds. Electrons are found in orbitals, which are sometimes represented as clouds. Orbitals are regions where there is a high probability of finding electrons of a given energy level. Orbitals are regions where there is a high probability of finding electrons of a given energy level. We can never know the exact location of an electron, only the probability of finding it in a given location. We can never know the exact location of an electron, only the probability of finding it in a given location.

Electron Configurations Each electron in an atom can be thought of as having a unique address, consisting of 4 quantum numbers- n, l, m, and m s. Each electron in an atom can be thought of as having a unique address, consisting of 4 quantum numbers- n, l, m, and m s. The orbitals are determined by n, l, and m. The orbitals are determined by n, l, and m. The Pauli exclusion principle states that each orbital can only hold 2 electrons. These two electrons must have different spins, denoted by m s, the spin quantum number. The Pauli exclusion principle states that each orbital can only hold 2 electrons. These two electrons must have different spins, denoted by m s, the spin quantum number.

n is the principal quantum number, discussed earlier. n is the principal quantum number, discussed earlier. l, which must be an integer, can range in value from 0 to n-1. This means that the first level can have only l value, the second level 2, the third level 3, etc. l, which must be an integer, can range in value from 0 to n-1. This means that the first level can have only l value, the second level 2, the third level 3, etc. These l values represent the type of orbital. Orbitals with l = 0 are s orbitals, those with l=1 are p orbitals, those with l=2 are d orbitals, those with l=3 are f orbitals. These l values represent the type of orbital. Orbitals with l = 0 are s orbitals, those with l=1 are p orbitals, those with l=2 are d orbitals, those with l=3 are f orbitals. Please refer to table 3-4 on p. 97 of your text.

The third quantum number, m, also an integer, can range from –l to l. This number represents the number of possible orbitals for the energy level. So, The third quantum number, m, also an integer, can range from –l to l. This number represents the number of possible orbitals for the energy level. So, s levels can only have 1 orbital each. s levels can only have 1 orbital each. p levels can have 3 orbitals each. p levels can have 3 orbitals each. d levels can have 5 orbitals each. d levels can have 5 orbitals each. f levels can have 7 orbitals each. f levels can have 7 orbitals each. Each orbital can have two electrons, one with spin up and one with spin down. Therefore, Each orbital can have two electrons, one with spin up and one with spin down. Therefore, s levels can hold 2 electrons (1x2). s levels can hold 2 electrons (1x2). p levels can hold 6 electrons (3x2). p levels can hold 6 electrons (3x2). d levels can hold 10 electrons (5x2). d levels can hold 10 electrons (5x2). f levels can hold 14 electrons (7x2). f levels can hold 14 electrons (7x2).

Each type of orbital has a characteristic shape.

These orbitals overlay one another.

Filling the energy levels The Aufbau principle says that energy levels fill from the lowest level up. However, sometimes orbitals with higher principal quantum numbers turn out to have lower energies than orbitals with lower principal quantum numbers. The following slide shows how they sometimes overlap.

Note how the 4s orbital fills before the 3d

The following chart shows the order in which the electrons fill the orbitals.

Writing electron configurations Determine the total number of electrons in the atom Determine the total number of electrons in the atom Fill levels from the lowest to the highest. Fill levels from the lowest to the highest. The superscript indicates the number of electrons in that sublevel. The superscript indicates the number of electrons in that sublevel. When you no longer have enough electrons to fill another sublevel, put the remaining electrons in the next sublevel. When you no longer have enough electrons to fill another sublevel, put the remaining electrons in the next sublevel.

Example: S (sulfur) S has a total of 16 electrons. S has a total of 16 electrons. The 1s sublevel can hold 2 electrons, leaving 14 remaining. The 1s sublevel can hold 2 electrons, leaving 14 remaining. The 2s sublevel can hold 2 electrons, leaving 12 remaining. The 2s sublevel can hold 2 electrons, leaving 12 remaining. The 2p sublevel can hold 6 electrons, leaving 6 remaining. The 2p sublevel can hold 6 electrons, leaving 6 remaining. The 3s sublevel can hold 2 electrons, leaving 4 remaining. The 3s sublevel can hold 2 electrons, leaving 4 remaining. These last 4 electrons go into the 3p sublevel. These last 4 electrons go into the 3p sublevel. The electron configuration of S is 1s 2 2s 2 2p 6 3s 2 3p 4. The electron configuration of S is 1s 2 2s 2 2p 6 3s 2 3p 4.