Chapter 8 - Electron Configuration and Electron Periodicity
chapter 8 - Electron Configuration and Chemical Periodicity
The laboratory work of nineteenth-century chemists created a massive amount of knowledge on the elements, which was arranged into the periodic table:
The first periodic table. Dmitri Mendeleev, a Russian scientist, organized the 65 elements known at the time into a table and described their behavior in the periodic law in 1870: when the elements are ordered by atomic mass, they display a periodic recurrence of identical qualities.
Mendeleev left vacant places in his table and was even able to anticipate the characteristics of some elements that were not found until later, such as germanium.
The contemporary periodic table. Today's periodic table (in the beginning of the book) includes 53 elements that were unknown in 1870 and, more crucially, organizes the elements by atomic number (number of protons) rather than atomic mass.
This modification is based on the work of British scientist Henry G. J. Moseley, who used electron bombardment to produce x-ray spectra of different metals. Moseley discovered that the greatest x-ray peak for each metal was connected to the nuclear charge, which rose by one with each subsequent element.
The three quantum numbers n, l, and ml, respectively, represent the size (energy), shape, and orientation in space of an atomic orbital.
A fourth quantum number describes spin, which is a characteristic of the electron rather than the orbital.
A beam of atoms with one or more lone electrons splits into two beams as it travels through a nonuniform magnetic field (produced by magnet faces with various shapes). The strong external magnetic field attracts half of the electrons while repelling the other half; the image attached above depicts this for a beam of H atoms.
The concept that an electron operates like a small magnet explains this finding.
Although an electron is not a little sphere, we may describe its magnetic behavior by picturing it spinning on its axis in one of two opposed directions, clockwise or counterclockwise; a spinning charge creates a tiny magnetic field whose direction is determined by the spin. As a result, two magnetic fields with opposite orientations exist, resulting in the two beams seen in the image attached above.
The spin quantum number (ms) has two possible values that correspond to the two orientations of the electron's magnetic field, +1/2 or - 1/2.
Helium (He; Z = 2) is the next element after hydrogen, and it is the first element with atoms that have more than one electron.
The first He electron has the same quantum numbers as the electron in the H atom, but the second He electron does not.
The Austrian physicist Wolfgang Pauli established a concept that specifies the arrangement of electrons in orbitals in atoms with more than one electron based on observations of excited states:
Pauli exclusion principle: no two electrons in the same atom can have the same four quantum numbers.
Therefore, the second He electron occupies the same 1s orbital as the first but has an opposite spin:
First electron: n = 1, l = 0, ml = 0, and ms = +1/2 (arbitrary assignment of ms value)
Second electron: n = 1, l = 0, ml = 0, and ms = –1/2
The main implication of the exclusion principle is that an atomic orbital can only carry a maximum of two electrons with opposing spins.
Two electrons can share the same orbital and hence have the same n, l, and m1 values, but each electron must have a distinct ms value. Because there are only two ms values, adding a third electron to that orbital requires repeating an ms value.
The 1s orbital in He, with its two electrons, is said to be full, and the electrons have paired spins. As a result, in an experiment like the one seen in the image attached above, a beam of He atoms is not divided.
Electrostatic phenomena, such as the attraction of opposing charges and the repulsion of similar charges, play a significant influence in defining the energy states of many-electron atoms.
Unlike the H atom, where the energy state is controlled only by the attraction between nucleus and electron, the energy states of many-electron atoms are additionally impacted by electron-electron repulsions.
You'll see in a moment how these extra interactions cause the splitting of energy levels into sublevels of varying energies: the energy of an orbital in a many-electron atom is primarily determined by its n value (size) and, to a lesser extent, by its l value (shape).
chapter 8 - Electron Configuration and Chemical Periodicity
The laboratory work of nineteenth-century chemists created a massive amount of knowledge on the elements, which was arranged into the periodic table:
The first periodic table. Dmitri Mendeleev, a Russian scientist, organized the 65 elements known at the time into a table and described their behavior in the periodic law in 1870: when the elements are ordered by atomic mass, they display a periodic recurrence of identical qualities.
Mendeleev left vacant places in his table and was even able to anticipate the characteristics of some elements that were not found until later, such as germanium.
The contemporary periodic table. Today's periodic table (in the beginning of the book) includes 53 elements that were unknown in 1870 and, more crucially, organizes the elements by atomic number (number of protons) rather than atomic mass.
This modification is based on the work of British scientist Henry G. J. Moseley, who used electron bombardment to produce x-ray spectra of different metals. Moseley discovered that the greatest x-ray peak for each metal was connected to the nuclear charge, which rose by one with each subsequent element.
The three quantum numbers n, l, and ml, respectively, represent the size (energy), shape, and orientation in space of an atomic orbital.
A fourth quantum number describes spin, which is a characteristic of the electron rather than the orbital.
A beam of atoms with one or more lone electrons splits into two beams as it travels through a nonuniform magnetic field (produced by magnet faces with various shapes). The strong external magnetic field attracts half of the electrons while repelling the other half; the image attached above depicts this for a beam of H atoms.
The concept that an electron operates like a small magnet explains this finding.
Although an electron is not a little sphere, we may describe its magnetic behavior by picturing it spinning on its axis in one of two opposed directions, clockwise or counterclockwise; a spinning charge creates a tiny magnetic field whose direction is determined by the spin. As a result, two magnetic fields with opposite orientations exist, resulting in the two beams seen in the image attached above.
The spin quantum number (ms) has two possible values that correspond to the two orientations of the electron's magnetic field, +1/2 or - 1/2.
Helium (He; Z = 2) is the next element after hydrogen, and it is the first element with atoms that have more than one electron.
The first He electron has the same quantum numbers as the electron in the H atom, but the second He electron does not.
The Austrian physicist Wolfgang Pauli established a concept that specifies the arrangement of electrons in orbitals in atoms with more than one electron based on observations of excited states:
Pauli exclusion principle: no two electrons in the same atom can have the same four quantum numbers.
Therefore, the second He electron occupies the same 1s orbital as the first but has an opposite spin:
First electron: n = 1, l = 0, ml = 0, and ms = +1/2 (arbitrary assignment of ms value)
Second electron: n = 1, l = 0, ml = 0, and ms = –1/2
The main implication of the exclusion principle is that an atomic orbital can only carry a maximum of two electrons with opposing spins.
Two electrons can share the same orbital and hence have the same n, l, and m1 values, but each electron must have a distinct ms value. Because there are only two ms values, adding a third electron to that orbital requires repeating an ms value.
The 1s orbital in He, with its two electrons, is said to be full, and the electrons have paired spins. As a result, in an experiment like the one seen in the image attached above, a beam of He atoms is not divided.
Electrostatic phenomena, such as the attraction of opposing charges and the repulsion of similar charges, play a significant influence in defining the energy states of many-electron atoms.
Unlike the H atom, where the energy state is controlled only by the attraction between nucleus and electron, the energy states of many-electron atoms are additionally impacted by electron-electron repulsions.
You'll see in a moment how these extra interactions cause the splitting of energy levels into sublevels of varying energies: the energy of an orbital in a many-electron atom is primarily determined by its n value (size) and, to a lesser extent, by its l value (shape).
