1. Honors Chem Chapter 4 The Tiny but Mighty Electron

2. Electron: What do we already know? It has a negative charge (Thomson) It is small in mass: approximately 0.0005 amu or 9.11x10 -28 grams (Milikan) Previously, it was thought of traveling in orbits around the nucleus (Bohr) Currently, it is thought of as moving rapidly outside of the nucleus in orbitals

3. Orbital A probable space outside of the nucleus where an electron is likely to be found Electrons are organized in their orbitals according to relative energy Low energy: Closest to the nucleus High energy: Farthest from the nucleus

4. What do orbitals look like? “In a science that we cannot see, a lot is left to the imagination…” –Correspondent from NOVA NOW

5. Level 1 In energy level 1, there is one orbital shape. Its called an s orbital and looks like a sphere.

6. Level 2 In energy level 2, there are two orbital shapes. First, an s orbital and then a p orbital. The p orbital is shaped like dumbbell. There are 3 of these shapes. Each one is a subshell.

7. Level 3 In energy level 3, there are three orbital shapes. One s, Three p’s, and Five d’s.

8. Level 4 In energy level 4, there are four orbital shapes. One s, Three p’s, Five d’s, and Seven f’s

9. Heisenberg Uncertainty Principle We cannot know the exact location and speed of an electron at the exact time. We can know one or the other precisely, but never both at the same time. WHY????

10. Finding An Electron: Quantum Numbers We can assign an electron a series of 4 numbers to “guesstimate” where it can be found in an atom. 4 parts: n, l, m, s

11. Quantum Numbers: Principle Energy Level Represented by the letter “n” Whole number: 1, 2, 3, 4, etc… Represents the size of the orbital. The bigger the number, the larger the orbital and also the further it is from the nucleus

12. Quantum Numbers: Angular Represented by the letter “l” Whole number: 0, 1, 2, and 3 Represents the shape of the orbital – For an “s” shape: l = 0 – For a “p” shape: l = 1 – For a “d” shape: l = 2 – For an “f” shape: l = 3

13. Quantum Numbers: Magnetic Represented by the letter “m” Can be a range of numbers from the negative integer to the positive integer Represents the orientation (or number of subshells) – For “s” shape: m = 0 (only one orientation) – For “p” shape: m= -1, 0, 1 (3 orientations) – For “d” shape: m = -2, -1, 0, 1, 2 ( 5 orientations) – For “f” shape: m = -3, -2, -1, 0, 1, 2, 3 (7 orientations)

14. Quantum Numbers: Spin Represented by the letter “s” Can be either +1/2 or -1/2 Represents the spin direction of the electron

15. Why can’t electrons stay in one place? Electrons are “hit” by ambient radiation sources and when they are given more energy they are “promoted” to a higher energy level This is borrowed energy so what goes up must come down. This release of energy comes in the form of light and/or heat!

16. Dual Nature: Particles and Waves Light is packets or bundles of energy, known as photons, that travel in waves (Einstein)

17. Electrons and Energy Photoelectric effect: electrons are emitted from samples of matter when they are exposed to radiation energy—LIGHT! Electrons have energy, absorb energy, and release energy Electrons and Light exhibit properties of both particles and waves

19. Why can’t electrons stay in one place? Electrons are “hit” by ambient radiation sources and when they are given more energy they are “promoted” to a higher energy level This is borrowed energy so what goes up must come down. This release of energy comes in the form of light and/or heat!

20. Sources of Energy Light/Heat Electricity Chemical Reaction Nuclear Reaction

21. Ground State The ground state of an electron is its lowest-energy state.

22. Excited State An excited state is when an electron has been promoted to a higher energy level. As the electron returns to its ground state, it releases the specific gained energy in the form of light

23. Bright Line Spectrum

25. Understanding Waves: Wavelength Wave = repetitive transfer of energy Wavelength (λ) = distance over which the wave’s shape repeats Generally measured in nanometers (1 x 10 9 nm = 1 m)

26. Understanding Waves: Frequency Frequency (ν) = number of occurrences of a repeating event per unit time Measured in Hertz (Hz) OR “waves per second” (1/s = s -1 )

27. Frequency and Wavelength Related Frequency and Wavelength are inversely related. One variable increases while the other decreases.

28. The Light Equation c = λν The speed of light (c) is equal to 3.0 x10 8 m/s The wavelength (λ) is in meters The frequency (ν) is in Hertz

29. Wave Practice An orange light has a wavelength of 492nm. What is the frequency of this light?

30. Energy and Frequency High frequency (short wavelength) is a high energy wave. Frequency and energy are directly related Both variables increase together

31. The Electromagnetic Spectrum

32. The Visible Spectrum ROY G BIV: Red, Orange, Yellow, Green, Blue, Indigo, and Violet Increasing energy Increasing frequency decreasing wavelength

33. Bright Line Spectrum Light is a combination of multiple colors Each atom has its own unique pattern of electrons that absorb energy differently No two atoms have the same bright line spectrum: used for identification Each line in the spectrum represents electrons releasing a specific amount of visible energy— translating to a specific color.