1	Periodicity and Atomic Structure Chapter 5
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3	Light and the Electromagnetic Spectrum
4	Wave - A vibrational disturbance which transmits energy.
5	Definitions Wavelength (l, Greek lambda) - The distance between identical points on successive waves. Frequency (n, Greek nu) - The number of peaks that pass a given point in a second frequency = cycles/sec = hertz = Hz Speed of light, c = ln - 3.00 x 10⁸ m/sec = 3.00 x 10¹⁰ cm/sec
6	"Sodium vapor lamps - the..." Sodium vapor lamps - the yellow street lights - emit light with l = 589.2 nm. What is its frequency?
7	"KPBS has a frequency of..." KPBS has a frequency of 89.5 MHz (MHz = 10⁶ cycles/sec). What is the wavelength of this radiation in meters?
8	Planck’s Quantum Theory Max Planck Blackbody radiation Intensity varies with wavelength (red-orange-white) Classical physics doesn’t explain
9	Try some more experiments
10	Experiment 1 Add an elemental gas to a cathode ray tube and get ----- colors Hydrogen (H₂) purple blue Neon (Ne) red orange Helium (He) yellow pink Argon (Ar) lavender Xenon (Xe) blue
11	Experiment 2 Shine white light through a prism -- rainbow A prism separates light of different wavelength, each color represents a different wavelength.
12	Experiment 3 Shine the colored light from our gas discharge tubes through a prism ¾® get distinct bands of color (light).
13	Quantization of energy Energies in atoms are quantized, not continuous. Quantized means only certain energies allowed.
14	Bohr model of the atom Electrons orbit the nucleus like little planets (planetary model) each with its own energy. Electrons can move from one energy level to another by absorbing or releasing energy. Energy is released as radiant energy or light.
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16	Quantum of energy the smallest quantity of energy that can be emitted (or absorbed) in the form of electromagnetic radiation. Energy (1 quantum) = hn or energy = n h n n = number of quanta of energy (must be a whole number) h = Planck’s constant = 6.626 x 10^-³⁴ J sec n = frequency
17	"What is the minimum energy..." What is the minimum energy of a sodium lamp (with l = 5.892 x 10^-⁷ m and n = 5.09 x 10¹⁴/sec)?
18	"Calculate the energy of a..." Calculate the energy of a quantum of blue light with wavelength = 410 nm.
19	Photoelectric Effect Observation - Electrons can be ejected from some metals when they are exposed to light. Is light behaving like a particle which can bounce electrons out of atoms? Light can behave as both a wave and a particle and energy is quantized the same either way.
20	"If a light with a..." If a light with a wavelength of 200 nm shines on sodium atoms with an ionization energy of 496 kJ/mol, what will be the speed of the electrons emitted?
21	deBroglie
22	deBroglie Wavelength Calculate the wavelength in nanometers associated with a 0.072 kg golf ball moving at 30 m/sec?
23	Quantized Energy
24	Energy Levels for H where n in an integer.
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29	What Next? Light behaves like waves --- and particles. Particles can behave like waves. Energy is quantized. ???????
30	Heisenberg Uncertainty Principle The first thing we would like to learn about electrons is where they are and how they travel. Heisenberg Uncertainty principle says this is impossible. (Dx)(Dmv) ³ h/4p (»10^-³⁴ kg m²/sec)
31	Schrodinger’s quantum mechanical model of the atom Ey = Hy y is the wave function or orbital y² (probability function) represents the probability of finding an electron at any given position in an atom.
32	Quantum Numbers The behavior of an electron is described mathematically by Schrodinger’s wave equation and each orbital contains as set of three variables called quantum numbers.
33	The principal quantum number (n) -- · an integer · determines energy level of orbital
34	Angular momentum quantum number (l)-- equal to (n-1) to 0 so for n = 1, l = 0 for n = 2, l = 0, or 1 for n = 3, l = 0, 1, or 2 · determines type of subshell of an electron quantum number subshell type 0 s 1 p 2 d 3 f
35	Magnetic quantum number (m_l) · equal to -l to +l in integer increments · identifies number of orbitals within a sublevel describes spatial orientation orbitals within a sublevel
36	Spin quantum number (m_s) · equal to +1/2 or -1/2 · necessary because each orbital contains 2 electrons and each electron needs its own space.
37	s orbitals · spherical in shape · one spatial orientation (m_l = 0) · contain nodes as move to higher quantum levels (nodes are places probability of finding an electron goes to zero) · makes sense if we look at electrons as waves, waves have nodes.
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40	p orbitals · dumbbell shaped · three different spatial orientations (m_l =1, 0, -1,)
41	d orbitals · cloverleaf shaped + one dumbbell in a doughnut · five different spatial orientations (m_l = 2, 1, 0, -1, -2)
42	f orbitals · complex shape (8 lobes) · seven different spatial orientations (m_l = 3, 2, 1, 0, -1, -2, -3)
43	Electronic Energy Levels
44	Electronic configuration of the atoms 1. Lowest energy orbitals are filled first. 2. Only 2 electrons (of different spin) allowed in each orbital. 3. When sublevels are filling, fill each orbital with 1 electron of same spin and then pair openly when necessary.
45	Pauli Exclusion Principle No more than two electrons can be assigned the same four quantum numbers. This means that no more than 2 electrons may occupy the same orbital.
46	Hund’s Rule Electrons pair only after each orbital in a subshell is occupied by a single electron.
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48	Electronic Configuration of the Ions Cations - Electrons are removed from the highest energy occupied orbital Anions - Electrons are added to the lowest energy unoccupied orbital For transition metals -- The highest ns electrons are removed first (even though they are not the last added)
49	Effective nuclear charge
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51	Atomic Radii Radius decreases as we move across the periodic table to the right.
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54	Ionic Radii Cations -- radius decreases due to an increase in Zeffective Anions -- radius increases due to crowding of more electrons into a shell
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