2. Ch. 5 Electrons in Atoms Objective: Understand where electrons are located within an atom, and the scientific progression to determine the location of the electrons

3. Ch. 5 Physics and Radiant Energy  Issac Newton tired to explain what was known about the behavior of light by assuming that light consists of particles  Around 1900, experimental evidence confirmed that light consists of waves  LIGHT TRAVELS IN WAVES SIMILAR TO THE WAVES CAUSED BY A MOVING BOAT  LIGHT IS A FORM OF ELECTROMAGNETIC RADIATION  X-RAYS, GAMMA RAYS, AND RADIO WAVES  CONSIST OF AN ELECTRIC AND MAGNETIC FIELDS OSCILLATING AT RIGHT ANGLES  Issac Newton tired to explain what was known about the behavior of light by assuming that light consists of particles  Around 1900, experimental evidence confirmed that light consists of waves  LIGHT TRAVELS IN WAVES SIMILAR TO THE WAVES CAUSED BY A MOVING BOAT  LIGHT IS A FORM OF ELECTROMAGNETIC RADIATION  X-RAYS, GAMMA RAYS, AND RADIO WAVES  CONSIST OF AN ELECTRIC AND MAGNETIC FIELDS OSCILLATING AT RIGHT ANGLES

4. RADIANT ENERGY Electromagnetic radiation consists of oscillating electric and magnetic fields that are perpendicular to the direction of propagation of the wave.

5. Electrons act like particles and waves  Properties:  Amplitude: the waves height from zero to crest  Wavelength (λ): the distance between crests  Frequency (γ): number of wave cycles to pass a given point per unit time  Units: cycles per second SI unit: Hertz  Wavelength and Frequency are inversely proportional Speed of light (c)- 3.0 x 10 8 m/s –light travels the fastest in a vacuum  Properties:  Amplitude: the waves height from zero to crest  Wavelength (λ): the distance between crests  Frequency (γ): number of wave cycles to pass a given point per unit time  Units: cycles per second SI unit: Hertz  Wavelength and Frequency are inversely proportional Speed of light (c)- 3.0 x 10 8 m/s –light travels the fastest in a vacuum

6. RADIANT ENERGY

7. Electromagnetic Spectrum

8. Bohr Model  Adapted Rutherford’s model to include discoveries of how energy of an atom changes when it absorbs or emits light  Studied simplest atom…..which is?  Niels Bohr discovered that H ’ s lone electron emits light in the discrete levels seen in the hydrogen spectrum - 5.3  He proposed:  1. Electrons are found only in specific circular paths, or orbits, around the nucleus  2. Each possible electron orbit has a specific energy  Adapted Rutherford’s model to include discoveries of how energy of an atom changes when it absorbs or emits light  Studied simplest atom…..which is?  Niels Bohr discovered that H ’ s lone electron emits light in the discrete levels seen in the hydrogen spectrum - 5.3  He proposed:  1. Electrons are found only in specific circular paths, or orbits, around the nucleus  2. Each possible electron orbit has a specific energy

9. Cont.  You can think of electrons existing in different energy levels just like you are climbing a ladder…  Ladder vs Bohr Model - drawing  Energy level- fixed energies an electron can have  Ground (n=1) vs excited state (n=2, 3, ect.)  Quantum- amt of energy required to move an electron from one energy level to another  Higher energy levels are closer together  Amt of energy needed to move between levels decreases the higher up you go  You can think of electrons existing in different energy levels just like you are climbing a ladder…  Ladder vs Bohr Model - drawing  Energy level- fixed energies an electron can have  Ground (n=1) vs excited state (n=2, 3, ect.)  Quantum- amt of energy required to move an electron from one energy level to another  Higher energy levels are closer together  Amt of energy needed to move between levels decreases the higher up you go

10. Atomic Emission Spectrum  Electric current energizes electrons, causing them to emit light  Ordinary light= mixture of all wavelengths  Light emitted by atoms consists of mixture of only specific frequencies  Each color corresponds to a specific frequency of visible light  We see these discrete lines as Atomic Emission Spectrum  Electric current energizes electrons, causing them to emit light  Ordinary light= mixture of all wavelengths  Light emitted by atoms consists of mixture of only specific frequencies  Each color corresponds to a specific frequency of visible light  We see these discrete lines as Atomic Emission Spectrum

11.  Emission Spectrum of an element is like its fingerprint  Each discrete line corresponds to one exact frequency of light emitted by the atom/sample  Absorption Spectrum – each discrete line corresponds to one exact frequency of light absorbed by the element/sample  Spectroscopy- analysis of electromagnetic radiation emitted by samples of materials  Allows Spectroscopist to analyze the concentrations of substances in a sample  Emission Spectrum of an element is like its fingerprint  Each discrete line corresponds to one exact frequency of light emitted by the atom/sample  Absorption Spectrum – each discrete line corresponds to one exact frequency of light absorbed by the element/sample  Spectroscopy- analysis of electromagnetic radiation emitted by samples of materials  Allows Spectroscopist to analyze the concentrations of substances in a sample

12. Explanation of Atomic Spectra  The light emitted by an electron moving from a higher to a lower energy level has a frequency directly proportional to the energy change of the electron  More energy emitted= higher frequency= different color light Violet = High energy Red = Low Energy **Remember the Balmer Seriers is the Visible Spectrum  The light emitted by an electron moving from a higher to a lower energy level has a frequency directly proportional to the energy change of the electron  More energy emitted= higher frequency= different color light Violet = High energy Red = Low Energy **Remember the Balmer Seriers is the Visible Spectrum

13. QUANTUM THEORY  * *Wavelength model of light could not explain why different elements glow different colors**  PLANCK’S THEORY  MAX PLANCK (1858-1947)  QUANTUM(A): (“Fixed Amount”)  THERE IS A FUNDAMENTAL RESTRICTION ON THE AMOUNTS OF ENERGY THAT AN OBJECT EMITS OR ABSORBS (PIECES OF ENERGY)  E=hv  * *Wavelength model of light could not explain why different elements glow different colors**  PLANCK’S THEORY  MAX PLANCK (1858-1947)  QUANTUM(A): (“Fixed Amount”)  THERE IS A FUNDAMENTAL RESTRICTION ON THE AMOUNTS OF ENERGY THAT AN OBJECT EMITS OR ABSORBS (PIECES OF ENERGY)  E=hv Energy Frequency Planck’s Constant = 6.6262 x 10 -34 J-s

14. CHAPTER 4LABORATORY CHEMISTRY13 QUANTUM THEORY  What is meant by energy quantization ?

15. QUANTUM THEORY  THE PHOTOELECTRIC EFFECT  ALBERT EINSTEIN (1879-1955)  ELECTRONS ARE EJECTED FROM THE SURFACE OF METAL WHEN LIGHT SHINES ON THE METAL  Used results and Planck ideas/research to say light has dual particle/ wavelength behavior  LIGHT BEHAVES LIKE TINY PARTICLES w/ specific QUANTA OF ENERGY (photon)  THE FREQUENCY OF THE PHOTON IS IMPORTANT, NOT THE NUMBER OF PHOTONS  THE PHOTOELECTRIC EFFECT  ALBERT EINSTEIN (1879-1955)  ELECTRONS ARE EJECTED FROM THE SURFACE OF METAL WHEN LIGHT SHINES ON THE METAL  Used results and Planck ideas/research to say light has dual particle/ wavelength behavior  LIGHT BEHAVES LIKE TINY PARTICLES w/ specific QUANTA OF ENERGY (photon)  THE FREQUENCY OF THE PHOTON IS IMPORTANT, NOT THE NUMBER OF PHOTONS

16. CHAPTER 4LABORATORY CHEMISTRY15 QUANTUM THEORY

17. Arthur Compton  Photon, traveling at the speed of light, hits electron like billiard balls  Light has property of both a wave and particle  Photon, traveling at the speed of light, hits electron like billiard balls  Light has property of both a wave and particle

18. CHAPTER 4LABORATORY CHEMISTRY17 ANOTHER LOOK AT THE ATOM  MATTER WAVES  LOUIS de BROGLIE (1892-1987)  PARTICLES OF MATTER BEHAVE LIKE WAVES AND EXHIBIT A WAVELENGTH, JUST AS WAVES OF LIGHT BEHAVE LIKE PARTICLES OF MATTER  PROVED BY 2 AMERICAN SCIENTISTS AT BELL LABS 3 YEARS LATER (C. DABVISSON AND L. GERMER)  ALL MOVING OBJECTS HAVE WAVELIKE BEHAVIOR  MATTER WAVES  LOUIS de BROGLIE (1892-1987)  PARTICLES OF MATTER BEHAVE LIKE WAVES AND EXHIBIT A WAVELENGTH, JUST AS WAVES OF LIGHT BEHAVE LIKE PARTICLES OF MATTER  PROVED BY 2 AMERICAN SCIENTISTS AT BELL LABS 3 YEARS LATER (C. DABVISSON AND L. GERMER)  ALL MOVING OBJECTS HAVE WAVELIKE BEHAVIOR

19. If all moving objects acts as waves, why cant I see these motions?  Wavelengths are very, very small (septillionth of a nm)  Classic mechanics describes the motions of bodies much larger that atoms, while quantum mechanics describes the motion of subatomic particles and atoms as waves  Heisenberg Uncertainty Principle- It is impossible to know both the velocity and position of a (small, electron-sized) particle at the same time  To locate an electron we strike it with a photon of light which in turn affects its motion in a way we cannot predict  Wavelengths are very, very small (septillionth of a nm)  Classic mechanics describes the motions of bodies much larger that atoms, while quantum mechanics describes the motion of subatomic particles and atoms as waves  Heisenberg Uncertainty Principle- It is impossible to know both the velocity and position of a (small, electron-sized) particle at the same time  To locate an electron we strike it with a photon of light which in turn affects its motion in a way we cannot predict

20. Quantum Mechanics reviewed …..  Bohr = Hydrogen Atom (emission spectrum)  Quantum mechanics  Light described as a quanta of energy (particles) = photon  Light acts as both waves and particles  Since electrons acts as waves they can be used to magnify objects  Electron microscope  Bohr = Hydrogen Atom (emission spectrum)  Quantum mechanics  Light described as a quanta of energy (particles) = photon  Light acts as both waves and particles  Since electrons acts as waves they can be used to magnify objects  Electron microscope

21. Problem with Bohr Model  Tested model on H atom  Failed to explain what happened with absorbed and emitted energies in atoms with more than one electron  So….  Tested model on H atom  Failed to explain what happened with absorbed and emitted energies in atoms with more than one electron  So….

22. The Quantum Mechanical Model  Erwin Schrodinger came up with a mathematical equation describing the behavior of an electron in a H atom.  Quantum Mechanical Model  Restricts energy of electrons to certain values like Bohr  Unlike Bohr, this model does not involve the exact path the electron takes around the nucleus  *** The Quantum Mech. Model determines the allowed energies an electron can have and how likely(probability) it is to find the electron in various locations***  Erwin Schrodinger came up with a mathematical equation describing the behavior of an electron in a H atom.  Quantum Mechanical Model  Restricts energy of electrons to certain values like Bohr  Unlike Bohr, this model does not involve the exact path the electron takes around the nucleus  *** The Quantum Mech. Model determines the allowed energies an electron can have and how likely(probability) it is to find the electron in various locations***

23. Cont.  The different probabilities of finding an electron are broken down into Principle Energy Levels (n)  n = 1,2,3,4, 5, 6, 7  Each energy level, n, may have energy sublevels, each corresponding to an orbital of a different shape which describes where the electron is likely to be found  ** the maximum # of electrons an orbital can hold is 2***  S= 1 orbitald= 5 orbitals  P= 3 orbitalsf= 7 orbitals  ** the maximum number of electrons that can be found in a certain principle energy level is calculated with the formula 2n^2  The different probabilities of finding an electron are broken down into Principle Energy Levels (n)  n = 1,2,3,4, 5, 6, 7  Each energy level, n, may have energy sublevels, each corresponding to an orbital of a different shape which describes where the electron is likely to be found  ** the maximum # of electrons an orbital can hold is 2***  S= 1 orbitald= 5 orbitals  P= 3 orbitalsf= 7 orbitals  ** the maximum number of electrons that can be found in a certain principle energy level is calculated with the formula 2n^2

24. Characteristics of the QMM  Atomic Orbital- A region of space in which there is a high probability of finding an electron  These areas have different energy levels  Due to the different densities of these areas the atomic orbitals can form certain shapes  Atomic Orbital- A region of space in which there is a high probability of finding an electron  These areas have different energy levels  Due to the different densities of these areas the atomic orbitals can form certain shapes

29. Electron Configuration  Arrangement of electrons in various orbital’s around the nucleus  Three rules used:  Aufbau Principle- Electrons occupy the orbitals of the lowest energies first  Pauli Exclusion Principle- An atomic orbital may describe at most two electrons. The electrons must have opposite spins  Hund’s Rule- The electrons occupy orbitals of the same energy in a way that makes the number of electrons with the same spin directions as large as possible  Arrangement of electrons in various orbital’s around the nucleus  Three rules used:  Aufbau Principle- Electrons occupy the orbitals of the lowest energies first  Pauli Exclusion Principle- An atomic orbital may describe at most two electrons. The electrons must have opposite spins  Hund’s Rule- The electrons occupy orbitals of the same energy in a way that makes the number of electrons with the same spin directions as large as possible

30. Two methods to writing electron configuration  Method 1: Diagonal Chart  1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, and 7p  Method 2: Periodic Table  Method 1: Diagonal Chart  1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, and 7p  Method 2: Periodic Table