Atomic Spectra Evidence for electron arrangement Types of spectra The entire range of colours or wavelengths/frequencies are present. Only specific colours/lines or wavelengths/frequencies are present. Specific colours or wavelengths/frequencies are missing. Flame tests Comment on the colours of the flames when different substances are heated in a Bunsen flame. Do different substances burn with different coloured flames? What is the species in the substances that is responsible for the colours? Are the flames characteristic of these species? What causes the flames to have these colours? What inference can you make regarding atomic structure from the observations through the spectroscope? http://webmineral.com/help/FlameTest.shtml Dalton’s atom Thomson’s atom Rutherford’s atom Bohr’s atom How did Bohr figure it out? Based on observations from atomic spectroscopy. White light consists of electromagnetic radiation of a specific range of frequencies. When an atomic sample is vaporised and heated, it emits light of a certain colour that can be analysed for the frequencies it contains using an atomic spectrometer. This emitted light contains only specific frequencies which appear as bright lines in an emission line spectrum. What do we know about light? Light is a form of electromagnetic radiation. Electromagnetic radiation is energy travelling in waves. A wave is described by its frequency (υ) and its wavelength (λ). They are related by λ = c/υ where c is the speed of light (3.00 x 108 m s-1). Wavelength and frequency C = λν problems…. Work out υ: If λ is 3.00 m If λ is 30.0 cm If λ is 3.00 mm Work out λ: If ν is 2.1 s-1 c = 3.00x108 m s-1 Electromagnetic Spectrum Energy of a wave We can consider light to be not just a wave, but also a particle. Hence, light can be emitted or absorbed in packets or quanta of energy (called photons). The energy of a light photon is given by E = hυ where h is the Planck constant (6.63 x 10-34 J s). From ground to excited state Hydrogen atom spectra decreasing energy E = hυ = hc/λ Bohr’s atomic model Niels Bohr, in 1913, proposed a model of an atom that helped to explain the observations. As spectra of large multi-electron atoms are very complex and difficult to interpret, Bohr used the spectra of hydrogen to develop his model. He stated that Electrons can only occupy fixed energy levels; each energy level being assigned a principal quantum number. The energy level closest to the nucleus being n = 1 and of lowest energy. The energy of an electron is quantised; the electron may only have certain energies. Bohr’s atomic model Electrons can only absorb or emit energy of specific frequencies. When the electron occupies the energy level of lowest energy the atom is said to be in its ground state. An atom can have only one ground state. If the electron occupies one of the higher energy levels then the atom is in an excited state. An atom has many excited states. spectral lines get closer at higher energies/higher frequencies/shorter wavelengths because energy gaps between successive higher energy levels get closer. increasing frequency / energy The Electronic Configuration of Atoms The structure of the atom is pictured as consisting of a positively charged nucleus surrounded by negatively charged electrons. The protons and neutrons are tightly packed together in the tiny nucleus. The Electronic Configuration of Atoms The electrons move rapidly and orbit round the nucleus, held by electrostatic forces of attraction. The strength of the attractive force of the nucleus on the electrons is known as the nuclear charge and depends on the number of protons in the nucleus. The electrons are found at relatively great distances away from the nucleus in clearly defined regions called electron shells that have specific energies (energy levels).
