1. UNIT 10 LIGHT & ELECTRONS S.Fleck 2015-2016

2. Unit Objectives Calculate the wavelength, frequency, or energy of light, given two of these values Explain the origin of the atomic emission spectrum of an element Apply the aufbau principle., the Pauli exclusion principle, and Hund’s rule in writing the electron configurations of elements Explain why the electron configurations for some elements differ from those assigned using the aufbau principle

3. Light … is electromagnetic radiation that is visible to the eye can travel through vacuum can penetrate through transparent materials but cannot pass through opaque objects travels in a straight line bounces back, is reflective behaves like waves and particles

4. Atomic Spectra Vocabulary Amplitude Wavelength (λ) Frequency (ν) Hertz (Hz or s -1 ) – units of frequency

5. Electromagnetic Spectrum

6. Visible Spectrum

7. Other types of Electromagnetic Radiation Infrared Radiation – heat lamps Radio waves Microwaves – heat only the food transfer of energy to the moisture in the food Ultraviolet (UV) – sunburns & skin cancer X-Rays Gamma Rays – Solar Flares

8. Visual Examples

9. Planck’s Theory German physicist (1858-1947) Planck proposed that there is a restriction on the amounts of energy that an object emits or absorbs, and he called each of these pieces an energy quantum. Quantum = Fixed amount Energy of a photon of light (E) is directly proportional to the frequency of light E = h x ν Planck’s Constant (h) = 6.6262 x 10 -34 J x s

10. Quantum Concept Every time an electron has an increase or decrease in energy it can only do so in steps. A small energy change involves the emission or absorption of low-frequency radiation and vise versa.

11. Planck’s Constant = 6.6262 x 10 -34 J x s Why can’t we feel the changes in energy around us? Each quantum of energy is extremely small. This makes energy seem continuous. The quanta are too small to notice in the everyday world.

12. Photoelectric Effect When light is shined on a metal, electrons are ejected. For each metal a minimum frequency of light is needed to release an electron.

13. Einstein & the Photoelectric Effect Einstein used Planck’s theory to propose that light consists of quanta of energy that behave like tiny particles. PHOTONS So why does violet light free the electrons and red light does not?

14. Wavelength & Frequency Calculation Practice Speed of Light = Wavelength x Frequency c = λν c = 3.00 x 10 10 cm/s Calculate the wavelength of the yellow light emitted by a sodium lamp if the frequency of the radiation is 5.10 x 10 14 Hz (5.10 x 10 14 s -1 ).

15. Planck’s Constant Calculation Practice E = h x ν h = 6.6262 x 10 -34 J x s Calculate the energy (in J) of a quantum of radiant energy (the energy of a photon) with a frequency of 5.00 x 10 15 s - 1.

16. Equations, Units, & Constants c = λν E = h x ν E = (h c)/λ h = 6.6262 x 10 -34 J s c = 3.00 x 10 8 m/s Wavelength (λ) = m Frequency (ν) = Hz or s -1 Speed of Light (c) = m/s Planck’s Constant (h) = J  s

17. Recap Energy Level – region around the nucleus where the electron is likely moving Quantum – a discrete level Light is going to be releasing a specific amount of energy A packet of energy Not a continuous range (rungs on a ladder) Max Planck measured these quanta. Quantified the relationship between frequency and energy.

18. Quantum Mechanical Model Because electrons have different energies, they are found in different probable locations around the nucleus. An atomic orbital is a 3-d region around the nucleus of an atom where an electron with a given energy is likely to be found. Orbitals (not orbits) have characteristic shapes, sizes and energies.

19. Quantum Mechanical Model A principle quantum number (n) is assigned to indicate the relative SIZE & ENERGY of atomic orbitals. As the quantum (n) increases, the orbital becomes larger and is further away from the nucleus. An atom’s principal energy levels are specified by n.

20. Quantum Mechanical Model 4 principal energy levels, consisting of one or more sublevels... As n increases, the # of sublevels increases as does their distance from the nucleus.

21. Quantum Mechanical Model Sublevels are labeled s, p, d, or f, according to the shapes of their orbitals. For n=1, there is one sublevel. It is called “s”, specifically “1s” For n=2, there are 2 sublevels. They are called “s” and “p” (or 2s,2p). For n=3, there are 3 sublevels. They are called....?

22. Quantum Mechanical Model Each type of sublevel consists of one or more orbitals. There is 1 “s” orbital There are 3 “p” orbitals There are 5 “d” orbitals There are 7 “f” orbitals

23. Periodic Table & Orbitals

24. Aufbau Principle Electrons enter orbitals of lowest energy first. Energy Level = n Energy Levels Orbital Name # of OrbitalsHolds # Electrons 1s12 2p36 3d510 4f714

25. Examples – energy level, orbital, & number of electrons Hydrogen Helium Lithium Beryllium Boron

26. Two more rules. Pauli Exclusion Principle - An atomic orbital may describe at most two electrons. Hund’s Rule – When electrons occupy orbitals of equal energy, one electron enters each orbital until all the orbitals contain 1 electron with parallel spins.