![]() It is possible to summarise Bohr’s atomic model by stating that the energy levels of the electrons are mostly determined by the size of their orbital planes. Similarly, to the motion of planets around the sun, Niels Bohr’s atomic theory has electrons with fixed sizes and energies travelling in orbits around a positively charged nucleus, analogous to the motion of planets around the sun. According to Bohr, his model was supported by the classical laws of physics as well as the quantum theory of radiation. Due to the emission of electromagnetic radiation by the charged particles, this model was able to alleviate the instability problem associated with the previous Rutherford model of the atom, which displayed a motion in which electrons would lose their energy and subsequently spiral into the nucleus as a result of the loss of energy. Watch the following video (Start at 32:27 minutes) for a discussion on the Bohr model and line spectra.In 1913, Niels Bohr proposed a quantized shell model of an atom in order to provide an explanation for how electrons can maintain a stable orbit around the nucleus. Although, in a multi-electron atom, if we know the wavelength of the emitted or the absorbed light, it can be related to the frequency and then finally to the energy differences between energy levels. In multi-electron species, there are electron repulsions, and the energy levels cannot be determined by one simple equation. Bohr’s postulates do not work well for atoms that have more than one electron. Bohr was able to explain both emission and absorption for an electron in the hydrogen atom. A transition from n = 4 → n = 2, would be emission of energy. #HYDROGEN BOHR MODEL SERIES#Below is an absorption spectral series for the electron in a hydrogen atom.Ī transition for an electron in the hydrogen atom from n = 1 → n = 3 would be absorption of energy. Eventually the electron will release the energy, as light, as it goes from a higher energy level to a lower energy level. ![]() When the electron gains energy it goes to a higher energy level, n, in a process called absorption. The electron in a hydrogen atom usually resides in the lowest energy state, n = 1. A gas has a very high kinetic energy, and an electron can be excited to a higher state. If a collision is forceful enough, the electron in an atom can gain enough energy to transition to a higher energy level. One way this can happen is through collisions between hydrogen atoms. The electron must gain energy to get from a lower to a higher energy level. We now ask how the electron in a hydrogen atom gets to a higher energy level. If we change n f to equal 1, we get the lines in the ultraviolet series and when n f = 3, we get the wavelengths in the infrared series for the hydrogen atom. This gives the four wavelengths in the visible region. This is the same constant as in the Balmer formula where n f = 2. Balmer, in 1885, showed the wavelengths in the visible region of the hydrogen atom could be determined by the following equation: The stability of the atom could not be explained by this theory. The electron would spiral into the nucleus within about 10 -10 s. According to classical physics, if an electron was orbiting a nucleus, it would continuously lose energy in the form of electromagnetic radiation. Rutherford’s model stated the atom was made up of a dense positively charged nucleus with negatively charged electrons in the extranuclear region. We will discuss this in more detail later in the study guide. This is due to the electrons in the sodium atoms gaining energy and then emitting the energy as yellow light. ![]() If you have ever let food boil over when cooking, you will notice the gas flame will turn a bright yellow-orange. The line spectrum is like a fingerprint and is specific for each element. Each element has its own line spectrum and a line spectrum can be used to identify different elements. Fireworks are made up of metal salts that when heated will produce light of different colors. When heated, elements will produce line spectra. We see that sodium has two yellow lines that are typical of some street lights. The line spectra below are for lithium and sodium. The colors in the hydrogen emission spectrum correspond to wavelengths of 410.1, 434.1, 486.1, and 656.3 nm. The spectrum is not continuous and only has four different colors. This is called a line spectrum which consists of radiation of specific wavelengths. If hydrogen gas is placed into a discharge tube, and a high voltage is applied, the result is 4 different colored lines. “Europa Rainbow” by Robert Couse-Baker is licensed with CC BY 2.0. ![]()
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