Introduction to Quantum Chemistry, Chemistry tutorial


Quantum chemistry mathematically describes the fundamental behaviour of matter at the molecular scale. It is, in principle, possible to explain all chemical systems using this theory. In practice, only the simplest chemical systems might realistically be examined in purely quantum mechanical terms, and approximations must be made for most practical purposes.

Definition of Quantum Chemistry

Quantum chemistry is the application of quantum mechanical principles and equations to the learn of molecules. In order to understand matter at its most fundamental level, we must utilize quantum mechanical models and techniques. There are 2 features of quantum mechanics that make it dissimilar from previous models of matter. The 1st is the idea of wave-particle duality; that is, the notion that we require to think of extremely tiny objects (these as electrons) as having characteristics of both particles and waves.

2nd, quantum mechanical models properly expect that the energy of atoms and molecules is always quantized, meaning that they may have only specific amounts of energy. Quantum chemical theories permit us to explain the structure of the periodic table, and quantum chemical calculations let us to precisely predict the structures of molecules and the spectroscopic behavior of atoms and molecules.

Quantum chemistry is a branch of theoretical chemistry, which applies quantum mechanics and quantum field theory to address questions and problems in chemistry. The description of the electronic behavior of atoms and molecules as pertaining to their reactivity is one of the functions of quantum chemistry. Quantum chemistry lies on the border between chemistry and physics, and important contributions have been made by scientists from both fields. It has a strong and active overlap through the field of atomic physics and molecular physics, as well as physical chemistry.

In quantum mechanics, the Hamiltonian, or the physical condition of a particle can be expressed as the sum of 2 operators, one analogous to kinetic energy and the other to potential energy. The Hamiltonian in the Schrödinger wave equation utilized in quantum chemistry doesn't enclose terms for the spin of the electron.

Solutions of the Schrödinger equation for the hydrogen atom provide the form of the wave function for atomic orbitals, and the relative energy of the various orbitals. The orbital approximation can be utilized to understand the other atoms for instance helium, lithium and carbon.

History of Quantum Mechanics

The history of quantum chemistry essentially began with the year 1838 discovery of cathode rays via Michael Faraday. The year 1859 statement of the black body emission problem via Gustav  Kirchhoff,  the 1877 suggestion via Ludwig Boltzmann, which the energy states of a physical system could be discrete, and the 1900 quantum hypothesis by  Max Planck that any energy radiating atomic system can  theoretically be divided into a  number of discrete energy elements  ε  such that each of these energy elements is proportional to the frequency ν with which they each individually radiate energy, as described via the subsequent formula:

Where h is a numerical value called Planck's constant. Then, in the year 1905, to explain the photoelectric effect (1839), i.e., that shining light on certain materials can function to eject   electrons from the material, Albert Einstein postulated, based on Planck's quantum hypothesis, that light itself consists of individual quantum particles, which afterward came to be called photons (1926). In the years to follow, this theoretical basis gradually began to be applied to chemical structure, reactivity, and bonding.

Electronic structure

The 1st step in solving a quantum chemical difficulty is usually solving the Schrödinger equation (or Dirac equation in relativistic quantum chemistry) by the electronic molecular Hamiltonian. This is said determining the electronic structure of the molecule. It can be said that the electronic structure of a molecule or crystal implies essentially its chemical properties. An precise solution for the Schrödinger equation can only be obtained for the hydrogen atom. Since all other atomic or molecular systems, engage the motions of 3 or more 'particles', their Schrödinger equations cannot be solved precisely and so estimated solutions must be sought.

Wave model

The foundation of quantum mechanics and quantum chemistry is the wave model, in that the atom is a tiny, dense, positively charged nucleus surrounded by electrons.  Unlike the previous Bohr model of the atom, however, the wave model describes electrons as "clouds" moving in orbitals, and their positions are symbolized through probability distributions rather than discrete points. The strength of this model lies in its predictive power. Particularly, it expects the pattern of chemically similar elements found in the periodic table. The wave model is so named since electrons exhibit properties (such as interference) traditionally associated with waves.  See wave-particle duality.

Valence bond

Although the mathematical basis of quantum chemistry had been laid by Schrödinger  in the year 1926,  it  is  generally  accepted  that  the  first  true calculation in quantum  chemistry  was  that of the German physicists Walter Heitler and Fritz London on the hydrogen (H2) molecule in year 1927. Heitler and London's process was extended via the American theoretical physicist John C. Slater and the American theoretical chemist Linus Pauling to become the Valence-Bond (VB) [or Heitler-London-Slater- Pauling (HLSP)] process. In this technique, attention is primarily devoted to the pair wise interactions between atoms, and this process therefore correlates closely with classical chemists' drawings of bonds.

Molecular orbital

An alternative approach was developed in 1929 by Friedrich Hund and Robert S. Mulliken, in that electrons are described via mathematical functions delocalized over an entire molecule.  The Hund-Mulliken approach or molecular orbital (MO) process is less intuitive to chemists, but has turned out capable of predicting spectroscopic properties superior than the VB process. This approach is the conceptional basis of the Hartree-Fock method and further posts Hartree-Fock methods.

Density functional theory

The Thomas-Fermi model was developed independently by Thomas and Fermi in 1927.  This was the first attempt to describe many-electron systems on the basis of electronic density instead of wave functions, although it was not very successful in the treatment of entire molecules. The method did provide the basis for what is now known as density functional theory.  Though this method is less developed than post Hartree-Fock methods, its lower computational requirements allow it to tackle larger polyatomic molecules and even macromolecules, which has made it the most used method in computational chemistry at present.

Chemical dynamics

A further step can consist of solving the Schrödinger equation with the total molecular Hamiltonian in order to study the motion of molecules. Direct solution of the Schrödinger equation is called quantum molecular dynamics, within the semi classical approximation semi classical molecular dynamics, and within the classical mechanics framework molecular dynamics (MD). Statistical approaches, using for example Monte Carlo methods, are also possible.

 Adiabatic chemical dynamics

In adiabatic dynamics, interatomic interactions are represented by single scalar potentials called potential energy surfaces. This is the Born- Oppenheimer approximation introduced by Born and Oppenheimer in 1927. Pioneering applications of this in chemistry were performed by Rice and Ramsperger in the year 1927 and Kassel in 1928, and generalized into the RRKM theory in the year 1952 by Marcus who took the transition state theory developed by Eyring in 1935 into account. These methods enable simple estimates of unimolecular reaction rates from a few characteristics of the potential surface.

Non-adiabatic chemical dynamics

Non-adiabatic  dynamics  consists  of  taking  the  interaction  between several coupled  potential energy surfaces (corresponding to different electronic  quantum  states  of  the  molecule).  The coupling terms are called vibronic couplings. The pioneering work in this field was done through Stueckelberg, Landau, and Zener in the 1930s, in their work on what is now identified as the Landau-Zener transition. Their formula permits the evolution probability between 2 diabatic potential curves in the neighborhood of an avoided crossing to be computed.

Quantum Chemistry and Quantum Field Theory

The application of quantum field theory (QFT) to chemical systems and theories has happen to increasingly general in the modern physical sciences. One of the 1st and most essentially explicit appearances of this is seen in the theory of the photomagneton. In this scheme, plasmas that are ubiquitous in both physics and chemistry are studied in order to find out the essential quantization of the underlying bosonic field. Though, quantum field theory is of attention in many fields of chemistry, as well as: nuclear chemistry, astrochemistry, sonochemistry, and quantum hydrodynamics. Field theoretic techniques contain as well been serious in developing the ab initio effective Hamiltonian theory of semi- empirical pi-electron methods.

Usefulness of Quantum Mechanics

The growth of quantum mechanics has led to the possibility of calculating energy levels and other properties of molecules. Its purpose has made it probable to comprehend the nature of chemical bonds. It as well gives the foundation for understanding spectroscopy (the learn of interaction of electromagnetic radiation through molecules). In summary, the quantum theory is the basis of all our notions about the make-up of atoms and molecules, the building blocks of nature.

The Postulates of Quantum Mechanics

Quantum mechanics is depends on a set of unproved assumptions termed postulates. Such are validated via the extent to that the laboratory measurements agree by the computed consequences. The postulates are:

Postulates 1

A quantum mechanical system is explained as entire when a function, termed the wave function exists or is particular. For any physical circumstances, must be finite, single-valued and continuous. The particular values of satisfying equation below are said eigen functions while the values of the energy corresponding to these are termed eigenvalues of the system.

The physical consequence of is vague but 2 or * where * is the complex conjugate of and can be interpreted as the probability of finding a particle inside a specified domain. For a one dimensional system, the domain is the interval: x and x+dx; for a 2-dimensional difficulty or system, and component of area is required whereas for a three-dimensional trouble or system, an element of volume is needed. It is assumed that

∫  -  *d

exists where d  = dxdydz

and that is normalized by requiring that ∫ -  *d  =1.

That means the integral of * over all space is 1.

Postulates 2

To every observable physical quantity of a given system, A, for example, there exists a corresponding linear operator, Â such that

                         Â = A

A is the eigenvalue, is the eigenfunction and Â, the operator. An operator is a sign that instructs one to do something. It is utilized in mathematical manipulation. For instance, the symbol d/dx, that instructs us to obtain the derivative by respect to x of whatever follows the symbol. Operators are often symbolized via superscript symbol, ?.

Commutation of Operators

Two operators  and G commute if ÂG= GÂ


Suppose Â=d/dx, G=x and (x) = x3 do the operators commute?

For the operators to commute ÂG (x) = GÂ (x)

ÂG (x) = d(x. x3)/dx = d(x4)/dx = 4x3

GÂ (x) = xd(x3)/dx = 3x2.x=3x3

Since 3x3 ≠4x3, the operators don't commute.

Linearity of an Operator

An operator is linear if Â(a  +bØ) = a + bÂØ  Where a, b are constants and , Ø are arbitrary constants.

Postulates 3

If 1and 2 represent possible states of a system, so does a linear combination of them, such that

3= a   1+b   2

where a, b are constants.

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