If you are missing a notebook, come find it on the front desk. • Extra help: Tuesdays 8:15-9:00 pm • Thursdays 7:00-8:00 pm • Chem is try!!
Atomic Emission Spectra • Read p. 54-59
What is light? • White light: reflection of all colors • Black light: absorption of all colors • Colors are each a different wavelength (λ: lamda) of light
Colors • Different wavelengths of light are seen as different colors.
Question of the Day • Where are the electrons around the nucleus?
Study the light emitted (produced) by atoms and ions to deduce (find out) the structure of atoms. • When an atom is “excited” its electrons gain energy and move to a higher energy level. To return to a lower energy level, electrons must lose energy. They do this by giving off light.
Continuous spectrum: all wavelengths of visible light contained in white light. • Light emitted by an atom can be separated into a line spectrum that shows exactly what frequencies of light are present.
Because the light emitted from atoms is a line spectrum (not a continuous spectrum) we determine that: • There are “discrete” (separate) energy levels for each atom that can only produce light of certain wavelengths (this is NOT ordinary white light!).
c=fλ (velocity of light = frequency x wavelength) • the greater the frequency the shorter the wavelength • ΔE = hf • (energy lost by the electron = h(constant) x frequency • Frequency (and thus, color) of the light depends on the amount of energy lost by the electron.
ΔE = hf c=fλ Frequency (f) increases Wavelength(λ) increases
When atoms are “exited” (energy is added) they produce light. • Not white or all-colored light, but one color at a time. • Different colors indicate (show) different energy levels.
Hydrogen Emission Spectrum • Only certain energy levels can occur (not a continuous spectrum)
Energy Level Diagram Increasing frequency • The larger the difference in energy, the greater the frequency (thus, the more purple the light). Visible
Increasing frequency Visible
Increasing frequency Further from nucleus Closer to nucleus
convergence:the lines in a spectrum converge (get closer together) as frequency increases. • related to how much energy is required to remove the electron from the atom (ionize)
If you are missing a notebook (lab or class/homework), please come to my desk at 5:10 pm today. • Extra help: Tuesdays 8:15-9:00 pm • Thursdays 7:00-8:00 pm • Chem is try!!
When electrons are excited, they emit colors in a line spectrum. Only certain wavelengths of light are produced. Since c=flamda only certain frequencies are produced. Since ChangeE=hf, only certain changes in energy occur. Thus, electrons can not be making all changes in Energy, but only changes between discrete, separate energy levels.
Electronic Structure Energy Levels Shells
Electronic Structure and the Periodic Table • 1st energy level = 2 • 2nd energy level = 8 • Electronic structure: number of electrons in each energy level • After second level – more complicated • they don’t fill in order – more later
2 2 2 8 2 8 5
H=1 • O=2,6 (two electrons in the first energy level, six in the second) • Al=2,8,3 • S= • Cl= • Ar= • Different isotopes have the same electronic structure and the same chemical properties!
Electron Behavior • Valence shell: outer energy level of an atom • determine the physical and chemical properties of an atom Valence Shell
1 8 2 3 4 5 6 7 Valence electrons Group Period
How many electrons in valence shell? • Al • Ne • Li • Ca
Electronic Structure of Atoms Zumdahl2: p. 307-312
Electronic Structure Energy levels Sub-levels Orbital Spin (electrons)
Energy Levels • Major shells (layers) around the nucleus • filled before higher levels are filled • 1st: 2 electrons • 2nd: 8 electrons
Sub-levels • Different shapes • s – sphere • one orbital • p – figure eight • three orbitals • d – • five orbitals • f – • seven orbitals
p Sub-level • p sub-level has three orbitals • px, py, pz
Orbitals • Each orbital can hold two electrons. • Electrons spin in opposite directions