Atomic Structure and the Periodic Table
The electronic structure of an atom determines its characteristics
Studying atoms by analyzing light emissions/absorptions Spectroscopy: analysis of light emitted or absorbed from a sample Instrument used = spectrometer Light passes through a slit to become a narrow beam Beam is separated into different colors using a prism (or other device) Individual colors are recorded as spectral lines
Electromagnetic radiation Light energy A wave of electric and magnetic fields Speed = 3.0 x 10 8 m/s Wavelength ( ) = distance between adjacent peaks Unit = any length unit Frequency ( ) = number of cycles per second Unit = hertz (Hz)
Relationship between properties of EM waves Wavelength x frequency = speed of light v = c Calculate the frequency of light that has a wavelength of 6.0 x 10 7 m. Calculate the wavelength of light that has a frequency of 3.7 x 10 14 s -1
Visible Light Wavelengths from 700 nm (red) to 400 nm (violet) No other wavelengths are visible to humans
Quanta and Photons Quanta: discrete amounts Energy is quantized restricted to discrete values Only quantum mechanics can explain electron behavior Analogy: Water flow
Another analogy for quanta A person walking up steps his potential energy increases in a quantized manner
Photons Packets of electromagnetic energy Travel in waves Brighter light = more photons passing a point per second Higher energy photons have a higher frequency of radiation Planck constant h = 6.63 x 10-34 J s E = hv The energy of a photon is directly proportional to its frequency
Deriving Planck s constant In a laboratory, the energy of a photon of blue light with a frequency of 6.4 x 10 14 Hz was measured to have an energy of 4.2 x 10 19 J. Use Planck s constant to show this: E = (6.63 x 10-34 J s) x (6.4 x 10 14 1/s) = 4.2 x 10 19 J
Evidence for photons Photoelectric effect the ejection of electrons from a metal when exposed to EM radiation Each substance has its own threshold frequency of light needed to eject electrons
Determining the energy of a photon Use Planck s constant! What is the energy of a photon of radiation with a frequency of 5.2 x 10 14 waves per second? E = hv
Another problem involving photon energy What is the energy of a photon of radiation with a wavelength of 486 nm?
Louis de Broglie proposed that matter and radiation have properties of both waves and particles (Nobel Prize 1929) = h m Calculate the wavelength of a hydrogen atom moving at 7.00 x 10 2 cm/sec m = mass h = Planck s constant = velocity
Hydrogen spectral lines Balmer series: n 1 = 2 and n 2 = 3, 4, Lyman series (UV lines): n 1 = 1 and n 2 = 2, 3,
Atomic Spectra and Energy Levels Johann Balmer noticed that the lines in the visible region of hydrogen s spectrum fit this expression: Observe the hydrogen gas tube, use the prism to see the frequencies of EM radiation emitted v= (3.29 x 10 15 Hz) x 1-1 4 n 2 n = 3, 4,
Rydberg equation: works for all lines in hydrogen s spectrum v= R H x 1-1 n 1 2 n 2 2 R H = 3.29 x 10 15 s -1 Rydberg Constant
Energy associated with electrons in each principal energy level Energy of an electron in a hydrogen atom E = -2.178 x 10-18 joule n 2 n = principal quantum number
Differences in Energy Levels of the hydrogen atom Use the Rydberg Equation OR Use the expression for each energy level s energy in the following equation: E = E final E initial
Niels Bohr s contribution Assumed e - move in circular orbits about the nucleus Only certain orbits of definite energies are permitted An electron in a specific orbit has a specific energy that keeps it from spiraling into the nucleus Energy is emitted or absorbed ONLY as the electron changes from one energy level to another this energy is emitted or absorbed as a photon
Summary of spectral lines When an e - makes a transition from one energy level to another, the difference in energy is carried away by a photon Different excited hydrogen atoms undergo different energy transitions and contribute to different spectral lines
The Uncertainty Principle Werner Heisenberg The dual nature of matter limits how precisely we can simultaneously measure location and momentum of small particles It is IMPOSSIBLE to know both the location and momentum at the same time
Atomic Orbitals more than just Erwin Schrodinger (Austrian) principal energy levels Calculated the shape of the wave associated with any particle Schrodinger equation found mathematical expressions for the shapes of the waves, called wavefunctions (psi)
Born s contribution Max Born (German) The probability of finding the electron in space is proportional to 2 Called the probability density or electron density
Atomic Orbital the wavefunction for an electron in an atom s high probability of e - being near or at nucleus ELECTRON IS NEVER AT THE NUCLEUS IN THE FOLLOWING ORBITALS: p 2 lobes separated by a nodal plane d clover shaped f flower shaped
More about orbitals Each orbital can hold 2 electrons Orbitals in the same subshell have equal energies
Quantum numbers like an address for an electron n = principal quantum number As n increases, orbitals become larger Electron is farther from nucleus more often higher in energy less tightly bound to nucleus
Quantum numbers l = angular momentum quantum number Values: 0 to n 1 Defines the shape of the orbital Value of l 0 1 2 3 Letter used s p d f
Quantum numbers m l = the magnetic quantum number Orientation of orbital in space (i.e. p x p y or p z ) Values: between l and l, including 0 Example: for d orbitals, m can be -2, -1, 0, 1, or 2 For p orbitals, m can be -1, 0, or 1
Quantum numbers m s = the spin number When looking at line spectra, scientists noticed that each line was really a closely-spaced pair of lines! Why? Each electron has a SPIN it behaves as if it were a tiny sphere spinning upon its own axis Spin can be + ½ or -1/2 Each represents the direction of the magnetic field the electron creates
Describe the electron that has the following quantum numbers: n = 4, l = 1, m l = -1, m s = +1/2 Principal level 4 4p orbital p orbital x spin up
Are these sets of quantum numbers valid? 3, 2, 0, -1/2 2, 2, 0, 1/2
Electron configuration: rules 1. Aufbau principle electrons fill lowest energy levels first 2. Pauli exclusion principle only 2 electrons may occupy each orbital 3. Hund s rule electrons spread out over orbitals of equal energy before doubling up
Special rules One electron can move from an s orbital to the d orbital that is closest in energy Only happens to create half or whole-filled d orbitals Examples: Cr, Cu
Noble Gas Configuration A shorter electron configuration Write the symbol for the noble gas BEFORE the element in brackets Write the remainder of the configuration Examples: Cl Cs
Energy level specifics s and d orbitals are close in energy Example 4s electrons have slightly lower energy than 3d electrons The s electrons can penetrate to get closer to the nucleus, giving them slightly lower energy 4s 3d