Bohr's Model of an Atom in Inorganic Chemistry

Department of Chemistry College of Education University of Salahuddin-Erbil Principle of Inorganic Chemistry Lecturer (3) (1stYear) Lecturer's name: Beriwan Muhammad Hammad Ameen Academic Year: 2021-2022

What is Bohrs Model of an Atom? Bohr model of the atom was proposed by Neil Bohr in 1915. It came into existence with the modification of Rutherford s model of an atom. Rutherford s model introduced the nuclear model of an atom, in which he explained that a nucleus (positively charged) is surrounded by negatively charged electrons. Bohr s model consists of a small nucleus (positively charged) surrounded by negative electrons moving around the nucleus in orbits. Bohr found that an electron located away from the nucleus has more energy, and electrons close to the nucleus have less energy.

Postulates of Bohrs Model of anAtom 1. In an atom, electrons (negatively charged) revolve around the positively charged nucleus in a definite circular path called orbits or shells. 2. Each orbit or shell has a fixed energy and these circular orbits are known as orbital shells. 3. The energy levels are represented by an integer (n=1, 2, 3 ) known as the quantum number. This range of quantum number starts from nucleus side with n=1 having the lowest energy level. The orbits n=1, 2, 3, 4 are assigned as K, L, M, N . shells and when an electron attains the lowest energy level, it is said to be in the ground state. 4. The electrons in an atom move from a lower energy level to a higher energy level by gaining the required energy and an electron moves from a higher energy level to lower energy level by losing energy.

Stationary states exist in which the energy of the electron is constant; such states are characterized by circular orbits about the nucleus in which the electron has an angular momentum mvr given by equation 1.6. The integer, n, is the principal quantum number. Energy is absorbed or emitted only when an electron moves from one stationary state to another and the energy change is given by below equation where n1 and n2 are the principal quantum numbers referring to the energy levels En1 and En2 respectively. E = En2 En1 = hv Where h is Planck s constant equal to 6.63 x 10 Js and v is the frequency of light.

Bohrs theory of the atomic spectrum of hydrogen: Bohr s model explains the spectral lines of the hydrogen atomic emission spectrum. While the electron of the atom remains in the ground state, its energy is unchanged. When the atom absorbs one or more quanta of energy, the electron moves from the ground state orbit to an excited state orbit that is further away. Energy levels are designated with the variable n. The ground state is n = 1, the first excited state is n = 2, and so on. The energy that is gained by the atom is equal to the difference in energy between the two energy levels. When the atom relaxes back to a lower energy state, it releases energy that is again equal to the difference in energy of the two orbits (see Figure 1). Figure 1. Bohr model of the atom: electron is shown transitioning from the n = 3 energy level to the n = 2 energy level. The photon of light that is emitted has a frequency that corresponds to the difference in energy between the two levels.

Based on the wavelengths of the spectral lines, Bohr was able to calculate the energies that the hydrogen electron would have in each of its allowed energy levels. He then mathematically showed which energy level transitions corresponded to the spectral lines in the atomic emission spectrum (Figure 2). Figure 2. The electron energy level diagram for the hydrogen atom.

He found that the four visible spectral lines corresponded to transitions from higher energy levels down to the second energy level (n = 2). This is called the Balmer series. Transitions ending in the ground state (n = 1) are called the Lyman series, but the energies released are so large that the spectral lines are all in the ultraviolet region of the spectrum. The transitions called the Paschen series and the Brackett series both result in spectral lines in the infrared region because the energies are too small.

Emission Spectrum of Hydrogen When an electric current is passed through a glass tube that contains hydrogen gas at low pressure the tube gives off blue light. When this light is passed through a prism (as shown in the figure below), four narrow bands of bright light are observed against a black background.

These narrow bands have the characteristic wavelengths and colors shown in the table below. Wavelength Color 656.2 486.1 blue-green 434.0 410.1 red blue-violet violet

Rydberg-Equation: (calculate wave number of hydrogen spectral line emissions due to the transition of an electron from one orbit to another). In 1901, Planck suggested that energy could be absorbed or emitted only in quanta of magnitude. Related to the frequency of the radiation ( ), The proportionality constant is h, the PlanckConstant (h = 6:626 X 10-34 J s). 1/ = RH(1/n12- 1/n22)

Q1/ What is the wavelength of light emitted when the electron, in a hydrogen atom undergoes transition from the energy level with n = 4 to the energy level with n = 2? In which region of electromagnetic spectrum does this radiation fall? Ans. According to Rydberg Equation = 486.3 nm visible region (Balmer series) Solving for the wavelength of this light gives a value of 486.3 nm, which agrees with the experimental value of 486.1 nm for the blue line in the visible spectrum of the hydrogen atom.

Zeeman effect : The "Zeeman Effect Experiment" is the splitting up of the spectral lines of atoms when they are placed in a magnetic field The idea of an e circling the nucleus in a well-defined orbit analogous to the moon circling the earth was easy to grasp and Bohrs theory gained wide acceptance. However it soon was realized that this simple theory was not complete description of the atom. Phenomenon is known of as a Zeeman Effect and it cannot be explained by the simple Bohr Theory but it explants by the combined Bohr-Sommerfild theory

The Uncertainty Principle One of the most important consequence of the dual nature of matter is the uncertainty principle, which was derived in 1927 by Werner Heisenberg [The Heisenberg Uncertainty principle] X . mV h/n X mV Uncertainty Uncertainty In the position In the momentum of e Of the particle of the particle Both the momentum and position of particle are uncertainty

Limitations of Bohrs Model of an Atom 1. Bohr s model of an atom failed to explain the Zeeman Effect (effect of magnetic field on the spectra of atoms). 2. It also failed to explain the Stark effect (effect of electric field on the spectra of atoms). 3. It violates the Heisenberg Uncertainty Principle. 4. It could not explain the spectra obtained from larger atoms