Unit 4: The Bohr Model of the Atom Properties of light Before the 1900 s, light was thought to behave only as a wave. Light is a type of electromagnetic radiation - a form of energy that exhibits wave like behavior. electromagnetic spectrum - is a collection of many kinds of electromagnetic radiation. Wave motion is characterized by: amplitude crest or peak trough
wavelength ( lambda--λ ) the distance between corresponding peaks on adjacent waves. frequency ( nu--ν ) the number of wave peaks that occur in a unit of time ( @ 1 second ). It is expressed in waves / sec. called a hertz ( Hz ). amplitude is half of the vertical distance from the top to the bottom of a wave. frequency and wavelength are related by... C = λν C = speed of light it is the same for all forms of electromagnetic radiation. All forms of electromagnetic radiation move at a constant 3.0 x 10 8 m/s (speed of light in a vacuum). Photoelectric Effect - Electromagnetic radiation strikes the surface of the metal, ejecting electrons from the metal and causing an electric current. This could not be explained by the wave theory b/c for some of the metals, no electrons were emitted (if the light s frequency was below a certain minimum). Light was known to be able of knocking loose an electron from metal; So the problem with the wave theory was it predicted that light of any frequency could accomplish this.
1900, Max Planck - studied emissions of light by hot objects and suggested that the object emits energy in small specific amounts he called quanta. Quantum - is the minimum quantity of energy that can be lost or gained by an atom. The equation shows the relationship between quantum and frequency: E = released energy in joules h = Planck s constant (6.626 E-34 ) J./Hz ν = frequency of radiation in Hz E = hν 1905, Albert Einstein expanded this theory w/ the idea that electromagnetic radiation has a dual wave-partical nature. light exhibits wavelike particles, and can be thought of as a stream of particles. Each particle carries a quantum of energy. photon - is a particle of electromagnetic radiation having zero rest mass and carrying a quantum of energy. E photon = hν He explained the photoelectric effect by proposing that electromagnetic radiation is absorbed only in whole numbers of photons. If the frequency is below the minimum to bind the electron to a metal s surface, it will not move.
Line-emission Spectrums - Light contains all wavelengths (and energies) of visible light and gives a continuous spectrum of colors (if it is passed through a prism): Continuous (white light) spectra Bohr knew that when pure elements are excited, they gave off distinct colors rather than white light. When an electric current is passed through a gas, a distinct color is given off by the element. When light from an excited element is passed through a prism, only specific lines (or wavelengths) of light can be seen. These lines of light are called line spectra. For example, when hydrogen is heated and the light is passed through a prism, the following line spectra can be seen: Hydrogen line spectra
Spectrum - is the unique collection of lines, (absorbed or emitted), for any element. Visible and U.V. spectroscopy are used to study the wavelengths of light absorbed and emitted by electrons in atoms. Each element has its own distinct line spectra. For example: Helium line spectra Neon line spectra To Bohr, the line spectra phenomenon showed that atoms could not emit energy continuously, but only in very precise quantities (he described the energy emitted as quantized). Because the emitted light was due to the movement of electrons, Bohr suggested that electrons could not move continuously in the atom (as Rutherford had suggested) but only in precise steps. Bohr hypothesized that electrons occupy specific energy levels. When an atom is excited, such as during heating, electrons can jump to higher levels. When the electrons fall back to lower energy levels, precise quanta of energy are released as specific wavelengths (lines) of light. When an excited atom with energy E2 falls back to energy E1, it releases a photon that has energy E2 E1 = Ephoton = hν. The specific frequencies suggested that energy states were fixed.
Classical theory predicted that hydrogen would be excited by whatever amount of energy was added. so... Quantum Theory was developed to explain that electrons can only exist in atoms at specific energy levels. Bohr Model of Hydrogen - 1913, Niels Bohr proposed a model linking electrons w/ photon emission. the electron can circle the nucleus in allowed paths or orbits. It has a definite fixed amount of energy. it cannot exist between orbits or it will fall to the next lower energy orbit and emit the difference in energy. Ground state - the lowest energy state of an atom Excited state - an atom having a higher potential energy than its ground state. Energy is released as an excited atom returns to it s ground state.
This energy-state diagram for a hydrogen atom shows some of the energy transitions for the Lyman, Balmer, and Paschen spectral series. Bohr s model of the atom accounted mathematically for the energy of each of the transitions shown. A series of specific wavelengths of emitted light makes up hydrogen s line-emission spectrum.the letters below the lines label hydrogen s various energy-level transitions. Niels Bohr s model of the hydrogen atom provided an explanation for these transitions.