Waves

1. Waves What do light waves have to do with chemistry?

2. The Big Idea! • The atoms of each element have a unique arrangement of electrons

3. Light and Quantized Energy • Light, a form of electromagnetic radiation, has characteristics of both a wave and a particle

4. Electromagnetic Radiation • Light is just one form of electromagnetic radiation • Electromagnetic radiation travels in the form of waves • All waves have amplitude, wavelength, frequency and speed

5. Electromagnetic Spectrum

6. US Frequency Allocations

7. Wavelength • Symbol is λ • Defined as the distance between two crests or two troughs

8. Speed of Light • Is it really a constant? • It equals 3.00 x 108 m/s • Symbol is c

9. Speed of Light • The speed of light (3.00  108 m/s) is the product of it’s wavelength and frequency c = λν.

10. Frequency • Its symbol is  • The number of wave cycles that pass a given point in one second • Measured in hertz (Hz) (cycles/sec) • C = λ

11. Frequency vs Wavelength • Long wavelengths have low frequencies • Short wavelengths have high frequencies

12. Who is Max Planck and why do we care? • The year is 1895 • Physicist’s are trying to model the EM radiation of a black body • A black body absorbs all light and emits it as black body radiation • Max is working hard to model this

13. 1895 – Viola!! • Success! • Devised Planck’s constant: Equantum= h E = amount of energy h = Planck’s constant: 6.6262 x 10-34 J· s (J = joule)

14. What does this mean? Matter can emit or absorb energy only in whole-number multiples of h : 1 h, 2 h, etc.

15. 1905 – Einstein comes on the scene • Publishes a paper on light quanta (Won Nobel Prize for this in 1921) • Explained photoelectric effect using this particle based model • Proved Planck’s model correct • Millikan – Tried to disprove both Planck and Einstein for 10 years

16. Einstein’s explanation of his theory • “According to the assumption considered here, in the propagation of a light ray emitted from a point source, the energy is not distributed continuously over ever-increasing volumes of space, but consists of a finite number of energy quanta localized at points of space that move without dividing, and can be absorbed or generated only as complete units.”

17. 1920’s - Louis de Broglie • Scientist who used Planck’s constant to predict that electrons could act like waves, since waves could act like particles • This principle is used in electron microscopes, where electrons streams are diffracted the same way light waves are diffracted by lenses

18. A picture taken with an electon scanning microscope

19. Photons • Particles of light are called photons • Each photon carries with it a specific value of energy based on its frequency • High frequency photons carry large amounts of energy

20. Atomic Emission Spectra • Light in a neon sign is produced when electricity is passed through a tube filled with neon gas and excites the neon atoms. • The excited atoms emit light to release energy. • When the light passes through a slit and then a prism, a line spectrum is formed.

21. Line Spectra

22. Atomic Emission Spectra • A line spectrum is also called an atomic emission spectrum • It’s formed by the energy given off, in the form of light, when an electron moves from its excited state to its ground state • Origin of Spectral Lines

23. Atomic Emission Spectra • The atomic emission spectrumof an element is the set of frequencies of the electromagnetic waves emitted by the atoms of the element when it is excited. • Each element’s atomic emission spectrum is unique.

24. Photograph of a line spectrum from helium

25. Several examples of line spectra

26. Quantum Theory and the Atom • Wavelike properties of electrons help relate atomic emission spectra, energy states of atoms, and atomic orbitals

27. Bohr’s Model of the Atom • Bohr suggested that an electron moves around the nucleus only in certain allowed circular orbits.

28. Bohr’s Model of the Atom • When an electron drops from a higher energy orbit to a lower-energy orbit, a photon is emitted.

29. Electrons are Weird • Heisenberg Uncertainty Principle: • It is impossible to know both the location and velocity of an electron at the same time • An electron can only be located when a photon strikes it. The collision causes it to change direction and velocity.

30. Heisenberg Uncertainty Principle

31. How are electrons arranged around an atom? • Bohr determined that electrons gained or lost energy in quanta, or specific amounts. These he labeled as quantum numbers, or n. • The n numbers refer to the energy levels, which are the same as the period numbers on the periodic table.

32. What are sublevels? • Each energy level is divided into sublevels • There are four different sublevels: s, p, d and f

33. Probability and Orbitals • Orbital: region around the nucleus where an electron is likely to be found • Have characteristic shapes, sizes and energies • Each can only hold 2 electrons with opposite spins

34. How do orbitals relate to energy? • Each principle energy level is divided into sublevels: s,p,d,f • Each sublevel has an odd number of orbitals: s has 1, p has 3, d has 5, and f has 7

35. What do the orbitals look like? Orbital in the s sublevel Orbitals in the p sublevel

36. Orbitals in the d sublevel Orbitals in the f sublevel Each one of these orbitals holds only two electrons. The d sublevel holds 10 electrons total, and the f sublevel holds 14 electrons total.

37. Orbitals

38. Energy sublevels are contained within the principal energy levels. Energy levels can be thought of as rows of seats in a theater. The rows that are higher up and farther from the stage contain more seats.

39. David's Whizzy Periodic Table

40. Electron Configuration • A set of three rules can be used to determine electron arrangement in an atom

41. What is electron configuration? • Each electron has a characteristic energy and a characteristic spin • Each electron is located in a specific orbital with a specific quantum number; this is called its electron configuration • There are three rules to use to determine where each electron is located and what spin it has

42. What are the rules for assigning configurations? • Aufbau Principle: electrons are added to an atom one at a time starting with the lowest energy orbital • 1s will get the first 2 electrons (lowest energy) • 2s will get the second 2 electrons (next higher energy) • 2p will get the next 6 electrons (next higher energy)

43. What are some examples? • Hydrogen: has 1 electron • From the Periodic Table, it is in row 1, so n = 1 • It has a 1s orbital • Its configuration is 1s1

44. Another example • Helium has 2 electrons • According to the P. Table, it is in the 1st row, so n = 1 • It has an s orbital • Its configuration is 1s2

45. Another example • Lithium has 3 electrons • The first two electrons are the same as in He, so the first term is 1s2 • It is in the 2nd row, so n=2 • Its configuration is 1s2 2s1

46. What is sublevel notation? • The type of electron configuration we have just done is also called sublevel notation • The order in which the levels fill is shown on the next slide and on your handout

47. You can tell which sublevel the electrons of an element are in by looking at the element’s location in the Periodic Table. S = pink p = blue d = green f = yellow

48. What is the next rule for assigning configurations? • Pauli Exclusion Principle: Each orbital can hold up to two electrons; must have opposite spins • Hund’s Rule: When electrons occupy orbitals of equal energy, one electron enters each orbital until all the orbitals contain one electron with spins that are parallel

49. What are electron spins? • Electrons have their own spins, represented by up and down arrows • They can spin clockwise or counterclockwise • Each ortibal holds two electrons, and they have to be spinning in opposite directions

50. How do I represent this? • Electron spins and the order in which they fill orbitals can be represented using an orbital diagram • Orbital diagrams use arrows and lines or squares to represent electrons in their orbitals