Atomic Spectra for Atoms and Ions What will you be doing in lab next week? Recording the line spectra of several different substances in discharge tubes. Recording the line spectra of several ions from salts present in a flame. Understanding the correlation between emission spectra & atomic structure. Identifying unknowns by their atomic fingerprints. Light is made up of different wavelengths White light is the combination of all wavelengths of visible radiation resulting in the colors of the rainbow 1
Line Spectra Not all radiation is a continuous spectrum of energies (colors). Neon sign it gives off the equivalent of orange light to our eye So, why do line spectra occur? Max Planck (1900) proposed that atoms in a heated solid could absorb or emit electromagnetic energy He observed that a quantum of energy is related to the frequency of that radiation: 2
Einstein's Description of Events Einstein (1905) considered energy as packets he called photons. He observed that when photons of light hit a substance, only photons of a certain energy were absorbed by the atom. Electrons may only become excited by input of a specific amount of energy. The energized electrons overcome their attraction to the nucleus and escape from their normal ground state. Bohr Model of the Atom Niels Bohr (1913) described the structure of the atom based on Planck s and Einstein s theories. Bohr proposed: The H atom has only certain allowable energy levels, called stationary states. The atom does not radiate energy while in one of its stationary states. The atom changes to another stationary state (one electron moves to another orbit) only when absorbing or emitting a photon whose energy equals the difference of energy between the two states. 3
Bohr s Theory Bohr proposed that the H atom had only specific energy states E = B 2 n What happens when the electron is far away? Bohr explains line spectra http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/flash.mhtml 4
Bohr s Equation predicts H atom lines B Ef = 2 nf Final State minus B Ei = 2 ni Initial State Transitions are Quantized on the H atom 5
Atoms emit various kinds of radiation Visible light emission n = 6 n = 5 n = 4 n = 3 n = 2 n = 1 Ultraviolet light emission Infrared light emission Ultraviolet light = electron relaxing from n > 1 to an n = 1 orbital. Visible light = electron relaxing from n > 2 to an n = 2 orbital. Infrared light = electron relaxing from n > 3 to an n = 3 orbital. Terminology When an atom has its electrons in their lowest possible energy levels, it is in its ground state. When an electron has been promoted to a higher level, it is in an excited state. Electrons are promoted through an electric discharge, heat, or some other source of energy An atom in an excited state can emit photons as the electron relaxes to the ground state. 6
The hydrogen atom All other atoms Multi-electron atoms In the hydrogen atom, energy levels with the same integer number are of the same energy In multi-electron atoms, ALL of the orbitals have different energy levels. 7
Discharge tubes & emission spectra Line Spectra 8
All elements show emission spectra Every element has a unique emission spectrum. The spectral lines identify the element like an atomic fingerprint. (e.g., atomic absorption) Flame ionization also provides enough energy, but heat is used to excite the elements instead of electricity. http://www.loc.gov/bicentennial/propage/ MA/ma_s_kennedy3.html What you will do in experiment 2? You will use discharge tubes in the lab to evaluate line spectra of various gases. You will also perform flame ionization to evaluate the emission spectra of ions once they are excited in a flame. 9
Q: Calculate the energy of a photon with wavelength 540 nm (green light). You are performing Spectroscopy! Spectroscopy Quantitative measurement of absorption & emission of energy Atomic spectroscopy information about the structure of the atom based on energy information Emission Spectroscopy analysis of light emitted from a energized atom or ion, giving an emission spectrum 10
Let s wrap up what we know about electrons 1. An electron in an atom occupies an orbital. 2. Each orbital has a specific energy. 3. The energies available to electrons are quantized. 4. Absorption of energy by an atom causes an electron to move from a lower-energy orbital to a higher energy orbital (the atom is now in an excited state). 5. An excited atom will relax to the ground state. 6. Electrons may return to a lower-energy level by emission of an amount of energy equal to the energy difference between the higher and lower energy levels. 7. Since the energies of the orbitals are quantized, the amount of energy emitted is quantized. 8. One can calculate the energy differences between orbitals by measuring the frequency, wavelength, or energy of light emitted. 11
Safety & Waste Disposal Absolutely NO substances found in a chemistry lab should be ingested. Wear gloves while using the metal salts. Prevent contamination of burners by using a small amount of water and metal salt. They will drip!! Turn off the Bunsen burners when finished and ensure the gas is turned off. Do not touch any part of the power supply unless it is unplugged. Do not look directly at the discharge lamps while they are illuminated; look only off to the sides of the lamp to observe the line spectra. Before next week s lab Print the prelab for Atomic Spectra for Atoms and Ions from the website and complete the problems. Prepare final draft of the Introduction section of the formal lab report. 12