Quantum mechanical model of atom

The Quantum Mechanical Model of the Atom

Energy Is Quantized

After Max Planck determined that energy is released and absorbed by atoms in certain fixed amounts known as quanta, Albert Einstein took his work a step further, determining that radiant energy is also quantized—he called the discrete energy packets photons. Einstein’s theory was that electromagnetic radiation (light, for example) has characteristics of both a wave and a stream of particles.

The Bohr Model of the Atom

In 1913, Niels Bohr used what had recently been discovered about energy to propose his planetary model of the atom. In the Bohr model, the neutrons and protons are contained in a small, dense nucleus, in which the electrons orbit in defined spherical orbits. He referred to these orbits as “shells” or “energy levels” and designated each by an integer: 1, 2, 3, etc. An electron occupying the first energy level was thought to be closer to the nucleus and have lower energy than one that was in a numerically higher energy level. Bohr theorized that energy in the form of photons must be absorbed in order for an electron to move from a lower energy level to a higher one, and is emitted when an electron travels from a higher energy level to a lower one. In the Bohr model, the lowest energy state available for an electron is the ground state, and all higher-energy states are excited states.

Orbitals and Quantum Numbers

In the 1920s, Werner Heisenberg put forth his uncertainty principle, which states that, at any one time, it is impossible to calculate both the momentum and the location of an electron in an atom; it is only possible to calculate the probability of finding an electron within a given space. This meant that electrons, instead of traveling in defined orbits or hard, spherical “shells,” as Bohr proposed, travel in diffuse clouds around the nucleus.

When we say “orbital,” the image below is what we picture in our minds.

To describe the location of electrons, we use quantum numbers. Quantum numbers are basically used to describe certain aspects of the locations of electrons. For example, the quantum numbers n, l, and ml describe the position of the electron with respect to the nucleus, the shape of the orbital, and its unique orientation, while the quantum number ms describes the direction of the electron’s spin within a given orbital.

Below are the four quantum numbers, showing how they are depicted and what aspects of electrons they describe.

Principal quantum number (n)	Has positive values of 1, 2, 3, etc. As n increases, the orbital becomes larger—this means that the electron has a higher energy level and is less tightly bound to the nucleus.
Second quantum number or azimuthal quantum number (l )	Has values from 0 to n – 1. This defines the shape of the orbital, and the value of l is designated by the letters s, p, d, and f, which correspond to values for l of 0, 1, 2, and 3. In other words, if the value of l is 0, it is expressed as s; if l = 1 = p, l = 2 = d, and l = 3 = f.
Magnetic quantum number (ml)	Determines the orientation of the orbital in space relative to the other orbitals in the atom. This quantum number has values from -l through 0 to +l.
Spin quantum number (ms)	Specifies the value for the spin and is either +1/2 or -1/2. No more than two electrons can occupy any one orbital. In order for two electrons to occupy the same orbital, they must have opposite spins.

Orbitals that have the same principal quantum number, n, are part of the same electron shell. For example, orbitals that have n = 2 are said to be in the second shell. When orbitals have the same n and l, they are in the same subshell; so orbitals that have n = 2 and l = 3 are said to be 2f orbitals, in the 2f subshell.

Finally, you should keep in mind that according to the Pauli exclusion principle, no two electrons in an atom can have the same set of four quantum numbers. This means no atomic orbital can contain more than two electrons, and if the orbital does contain two electrons, they must be of opposite spin.

 

Types of chemical bonds

Important Points

• What holds atoms together?
• Review the electron shell model of the atom, valence electrons
• Define chemical bond
• Six types of bonds: ionic, metallic, covalent, polar, hydrogen, and van der Waals. Be able to define each, provide examples, characterize their properties

Atoms in combination: the chemical bond

• So far we discussed the nature of the atom in some detail, and have a qualitative sense of how it looks. However, a lone, non-interacting atom is rare. Most atoms are found in combination with others.
• After the big bang event, the universe began to rapidly expand, and quite soon, within a few minutes, a major component was neutrons. Neutrons are not stable by themselves, so many of them split into protons and electrons, which form a significant component of the universe after about 10-15 minutes. It took about 100,000 years for the temperature of the universe to cool enough for the electrons to attach themselves to the protons and actually form atoms. So, about 100,000 years after the big bang, atoms became a significant component of the universe.
• This tells us that the energies of keeping electrons around a nucleus are much smaller than those associated with the nucleus or the formation of electrons.
• We live in a world of electrons, all our senses, and life itself, is manifest by variations in electronic interactions. Therefore life and humanity can only exist at the lower energy conditions in which electrons are bound to nuclei, i.e. Earth-like conditions and not Sun-like conditions.
• Electrons are the glue that holds groups of atoms together.

Electron shells and chemical bonding

• Let's review the nature of the atom from the point of view of the electrons.
• The atom is mainly low-density space with a very small but dense nucleus that defines the center of the atom.
• Electrons are located around the nucleus.
• These electrons can be classified in terms of shells that correspond to the rows in the periodic table. Each shell can fit a certain number of electrons, depending on how far away from the nucleus it is. A shell that is close to the nucleus can only contain a small number of electrons, otherwise, the electrons are too close together, and electrostatic repulsive forces push them apart. 
• Shells of electrons are most stable when they contain the maximum number of electrons that they can hold. On the one hand, if there are too few electrons, the electrons are constantly whizzing about, trying to fill all available space. Therefore, they have high kinetic energy. On the other hand, if there are too many electrons then electrostatic repulsion takes over and pushes them apart.
• Define the electrons in an unfilled outer shell as valence electrons.
• Since these are the outermost electrons, these are the ones that are perturbed by bringing another atom close by. These are the ones that are involved in bonding.
• Define: a chemical bond is the result of a redistribution of electrons that leads to a more stable configuration between two or more atoms.