1. Light, Quantitized Energy & Quantum Theory CVHS Chemistry Ch 5.1 & 5.2

2. Problems w/ Rutherford’s model Didn’t explain how electrons were occupying space outside the nucleus Didn’t explain why the electrons didn’t fall into the positively charged nucleus Didn’t explain differences in chemical behavior of different elements

3. Observation Early 1900’s Certain atoms emitted light when heated in a flame Analysis of emitted light indicated element’s chemical behavior is related to the arrangement of its electrons

4. Flame Tests

5. Wave Nature of Light Amplitude: wave height from origin (vertical center) to crest or trough Wavelength ( λ) – Measured from crest to crest or trough to trough – The wavelength of light is usually expressed in nanometers (1 nm = 1 x 10 –9 m).

6. Wave Nature of Light Frequency (  : nu – # of waves that pass a given point per second – One hertz (Hz), the SI unit of frequency, equals one wave per second. In calculations, frequency is expressed with units of “waves per second,” or (s –1 ) where the term “waves” is understood

7. Speed of Light (c) All light travels at 3.00 x 10 8 m/s in a vacuum Speed of light is the product of it’s frequency and wavelength Microwaves are used to transmit information. What is the wavelength of a microwave having a frequency of 3.44 x 10 9 Hz? Solve the equation relating the speed, frequency, and wavelength of an electromagnetic wave for wavelength (λ). Substitute c and the microwave’s frequency,, into the equation. Note that hertz is equivalent to 1/s or s –1. Divide the values to determine wavelength, λ, and cancel units as required.

8. Particle Nature of Light The wave model of light cannot explain why heated objects emit only certain frequencies of light at a given temperature, or why some metals emit electrons when colored light of a specific frequency shines on them

9. Planck & The Quantum Concept Matter can gain or lose energy only in small, specific amounts called quanta. – That is, a quantum is the minimum amount of energy that can be gained or lost by an atom. – It appears as though water is absorbing energy constantly in a microwave but in reality, it is being absorbed in little packets called quanta

10. Planck’s Equation Energy of a quantum is related to its frequency Planck’s constant has a value of 6.626 x 10 –34 J · s, where J is the symbol for the joule, the SI unit of energy According to Planck’s theory, for a given frequency,, matter can emit or absorb energy only in whole- number multiples of h ; that is, 1h, 2h, 3h, and so on Notice in the above equation, waves have mass: The wave/particle duality of light! Light is a Wavicle

11. Practice A helium-neon laser emits light with a wavelength of 633 nm. What is the frequency of this light? What is the wavelength of X rays having a frequency of 4.80 x 10 17 Hz?

12. Photoelectric effect In the photoelectric effect, electrons, called photoelectrons, are emitted from a metal’s surface when light of a certain frequency shines on the surface. That is, while a beam of light has many wavelike characteristics, it also can be thought of as a stream of tiny particles, or bundles of energy, called photons Thus, a photon is a particle of electromagnetic radiation with no mass that carries a quantum of energy.

13. Photoelectric Effect Further, Einstein proposed that the energy of a photon of light must have a certain minimum, or threshold, value to cause the ejection of a photoelectron. – That is: The photon of light must have enough energy to bump the electron off the metal

14. Atomic Emission Spectra The atomic emission spectrum of an element is the set of frequencies of the electromagnetic waves emitted by atoms of the element Each emission spectra is unique, like a fingerprint

15. Bohr’s Model of the Atom Since Energy is quantitized (Plank & Einstein) atoms only have 1 allowable energy state – Ground State: Lowest Energy state – Excited State: Energy added above ground state – Energy state is related to location of electron around nucleus Ground state: Close to nucleus Excited state: Away from nucleus – Father away = More Energy

16. Bohr’s Quantified Model Bohr assigned a quantum number, n, to each orbit and even calculated the orbit’s radius. For the first orbit, the one closest to the nucleus, n = 1 and the orbit radius is 0.0529 nm; for the second orbit, n = 2 and the orbit radius is 0.212 nm; and so on. When electron is in ground state n=1 When atom absorbs energy, electron moves up and n=2 When electron moves up, the atom emits a photon w/ the energy difference between the two levels

17. Hydrogen’s Line spectrum The four electron transitions that account for visible lines in hydrogen’s atomic emission spectrum are shown. Note that electron transitions from higher- energy orbits to the second orbit account for all of hydrogen’s visible lines

18. Louis de Broglie’s Atom Convinced Bohr was incorrect Treated electrons orbit as a wave Developing his idea, de Broglie derived an equation for the wavelength (λ) of a particle of mass (m) moving at velocity (ν). NOTICE: Waves have mass!

19. Heisenberg Uncertainty Principle Measuring the location or velocity of an electron changes the location and velocity of the electron The Heisenberg uncertainty principle states that it is fundamentally impossible to know precisely both the velocity and position of a particle at the same time.

20. Physicist Erwin Schrödinger (1926) & Wave Equations Treated the electron as a wave – Aka: quantum mechanical model Seemed to apply to all atoms, not just hydrogen like Bohr’s model Puts electrons in distinct energy levels but doesn’t describe its path around the nucleus Atomic Orbital: 3-D region that describes an electrons most probable location around the nucleus – Sort of a fuzzy cloud – The circle represents the 90% probability area for finding an electron

21. Quantum Numbers Principal Quantum Number (n) – Distance from the nucleus – The # of sublevels found in an energy level is equal to the number of the energy level Energy Sublevels: describe by the shape of the atom’s orbitals – s, spherical (1) – p, peanut or dumbell (3) – d, daisy (variable) (5) – f, flower (variable) (7)