What are Waves?

1. What are Waves? What do you think of? Types of Waves: Parts of waves:

2. Waves Light waves - vibration of photons Types of Waves: Sound waves - Vibration of air molecules and atoms Water waves - vibration of water molecules Parts of waves: Amplitude Trough Wavelength Frequency Crest

3. Parts of Waves? Amplitude Wavelength Trough Crest

4. What waves do: Reflection: Diffraction: Refraction

5. What waves do: Interference Destructive Constructive

6. Particles (Matter) What do you think of? What is Matter? What force defines how matter can behave?

7. Is Light A Wave or Particle ·This has been a debate for the same amount of time as the debate about what matter is made up of Ancient Greeks ·Aristotle: originally thought that light had to be a disturbance of the element AIR ·Democritus: originally thought that light had to be composed of small subatomic particles: Solar Atoms

8. ·Christian Huygens (1690): ·Predicted that light was a wave using the idea that light propagated as a disturbance in the air ·Sir Isaac Newton (1690) ·Light was made up of particles corpuscle ·Robert Hooke: Newton’s Rival ·Stated that Newton was wrong—light was a wave ·Newton did not publish his theory until after Hook passed away

9. ·Thomas Young (1773): ·First to actually use evidence that light was actually a wave ·Double slit experiment ·Same experiment was later used to prove that light was actually a particle

10. ·Do you think light is a wave or a particle? WHY?!?!? Lasers can pop balloons Rainbows (prisms) Shadows Lasers can be used in surgery to cut parts of the body

11. Duality of Light Video

12. Electrons Chapter 5.1

13. Light as a wave Originally scientist believed that light was a wave and that it only acted as such However, as science progressed this concept changed

14. Electromagnetic radiation: A form of energy that exhibits wavelike behavior as it travels through space Radio waves Microwaves Visible Light X-rays Gamma rays

15. Parts of Waves: Wavelength Crest Amplitude Trough Crest: Top of the wave Trough:Bottom of the wave Wavelength (λ): Is the shortest distance between equivalent points on a continuous wave (crest to crest or trough to trough) commonly measured in meters, centimeters or nanometers Amplitude: Is the wave's height from the origin to a crest, or from the origin to a trough. (wavelength and frequency do not affect amplitude) Frequency (υ): Is the number of waves that pass a given point per second. Hertz (Hz) is the SI unit for frequency

16. Wave Mathematical Relationships C = λ * υ λ - Wavelength υ - frequency C - Speed of light Speed of light is 3.00 x 108 m/s This speed is constant for all electromagnetic waves inside a vacuum (space) Note: As frequency increases, wavelength decreases (inverse relationship) Note: As frequency increases, energy of the wave increases Practice: What is the frequency of an X-ray with a wavelength of 1.15 x 10-10 m?

17. Electromagnetic Spectrum: A spectrum that includes all forms of electromagnetic radiation, with the only difference in the types of radiation being their frequencies and wavelengths ROYGBIV

18. Light as a particle: Light as a wave failed to explain: why heated objects emit only certain frequencies of light These colors correspond to different wavelength and frequencies Why some metals emit electrons when light at a given temperature shines on them (photoelectric effect)

19. Max Plank (1900) A German physicist was searching for an explanation of this phenomenon He found that matter could either gain or lose energy but only in small specific amounts called quanta Quantum-- is the minimum amount of energy that can be gained or lost by an atom (think of stepping stones)

20. Energy of a Quantum Plank came up with a relationship between the energy of a quantum and the frequency of a wave E = h * υ E -- Energy h -- Plank's Constant h = 6.626 x 10-34 J*s υ -- frequency This showed scientists that these quantum (abortions and emission of energy) were whole number multiples of hυ 2hυ 4hυ 6hυ hυ 7hυ 3hυ 5hυ 8hυ

21. Photoelectric Effect Some metals will eject electrons form their surface with light of a certain frequency (or higher) hits their surface

22. Albert Einstein (1905) Duality of light--light can be both a wave and a particle It is a beam of bundles of energy called photons Photons--a massless particle that carries a quantum of energy. This energy depends on the frequency of the photons EPhoton = hυ Ephoton = Energy of the photon h = Plank's Constant υ= frequency The blue color in some fireworks occurs when copper chloride is heated to approximately 1500K and emits blue light of wavelength 4.50x102 nm. how much energy does one photon of this light carry?

23. Atomic Emission Spectra Electrons around an atom's nucleus will absorb energy in quantums, the electrons will then jump up and fall back down to what is called ground state and release that same amount of energy. This energy has a specific frequency which can be seen as colors; meaning the frequency of the photons release is within the visible light of the electromagnetic spectrum Each element has a very specific range of colors that are emitted Atomic Emission Spectra: the set of frequencies of the electromagnetic waves emitted by atoms of the element These are used to identify elements and elements within compounds

24. Quantum Theory and the Atom Chap. 5.2

25. Niels Bohr (1913) Studied the hydrogen atom and, based off of Planck's and Einstein's concepts of quantized energy, determined that the atom only had certain allowable energy states ·Lowest possible energy state is called Ground State ·When the atom absorbs energy, it is said to be in an Excited State

26. Bohr's Model of the Atom Borh suggested that the electrons around the hydrogen atom could only be allowed in certain circular orbits around the nucleus ·The smaller the electron's orbit, the lower the atom's energy state or energy level ·The larger the electron's orbit, the higher the atom's energy level

27. Quantum Number: The number Borh gave to each orbital around the atom

28. Energy State Ground State for Hydrogen is when hydrogen's single electron is in the first energy level or the first quantum level ·The hydrogen atom does not give off energy in the ground state Once energy is added the single electron moves up to a higher energy orbit (such as n = 2) making the atom in an excited state ·The electron will fall back into its original quantum level (ground state) and release the energy gained as a photon Since these quantums are set only a set energies can be absorbed and emitted by the atom, therefore only specific frequencies are emitted by the atom

29. Louis de Broglie (1924) Thought that if Light can be have both wave and particlelike characteristics, then can matter (electrons) behave as both Predicted that all moving particles have wave-like characteristics (including cars or baseballs) λ = h/mυ λ = Wavelength h = planck's constant υ = frequency m = mass

30. Werner Heisenberg (1901-1976) ·Stated that it is impossible to take any measurement of an object without disturbing the object Heisenberg uncertainty principle: states that it is fundamentally impossible to know precisely both the velocity and position of a particle at the same time ·Meaning Bohr's defined orbits were not accurate

31. Erwin Schrodinger (1926) Austrian Physicist Quantum Mechanical model of the atom: the atomic model in which electrons are treated as waves ·This allowed for scientist to determine particular volumes of space around the nucleus in which the probability of finding an electron is very high Atomic Orbital: the probable location of an electron within an atom

32. Hydrogen's Atomic Orbitals Principal quantum number (n)--number of the atomic orbitals. (also called Principal energy level) ·As n increases, the orbitals become larger and have more energy These levels contain what are called energy sublevels ·The first energy level contains 1 sublevel; the second energy level contains 2 sublevels,the third energy level contains 3 sublevels, ect. There are a total of 4 sublevels labeled s, p, d and f ·These sublevels are then broken down further into orbitals ·A single orbital can only hold 2 electrons total ·Meaning if there are 3 orbitals there can be a total of 6 electrons in that sublevel

33. Periodic breakdown of levels and sublevels

34. S-Sublevel The s-sublevel has a spherical shape The s-sublevel only has one orbital and therefore s-sublevel can only hold two electrons

35. P-Sublevel The p-sublevel is a dumb-bell shape The p-sublevel contains 3 orbitals and therefore contains 6 electrons total

36. d-sublevel The d-sublevel has two shapes one that like two dumb-bells put together and the other is like a single dumb-bell The d-sublevel contains 5 orbitals and therefore can hold up to 10 electrons

37. f-sublevel The f-sublevel has a very complex shape The f-sublevel contains 7 orbitals and therefore contains 14 electrons

38. Periodic breakdown of levels and sublevels Notice that the energy level 1 contains only 1 sublevel (s), energy level 2 contains 2 sublevels (s and p), energy level 3 contains 3 sublevels (s, p, and d), and energy level 4 contains 4 sublevels (s, p, d, and f) ·Which is the order of the orbitals: s, p, d, and f

39. Electron Configuration Chap. 5.3 1s 2s2p 3s3p4s3d4p 5s4d5p6s4f 5d6p7s5f6d7p

40. Review Energy levels can be found by looking at what on the periodic table?

41. Basic goal: Break the energy levels down into sub-levels and orbitals! Energy levels Sub-levels (s, p, d, and f) Orbitals Only 2 electrons

42. Electron Configuration This is the arrangement of electrons in an atom Review How many electrons can the s-sublevel hold? How many electrons can the p-sublevel hold? How many electrons can the d-sublevel hold? How many electrons can the f-sublevel hold? Where can each of the sublevels be found on the periodic table?

43. Rules for Electron Configuration There are three rules that must be followed while writing electron configurations Rule 1: Aufbau Principle ·Each electron occupies the lowest energy orbital available Rule 2: Pauli Exclusion Principle ·A maximum of two electrons can occupy a single orbital, but only if the electrons have opposite spins Rule 3: Hund's rule ·Single electrons with the same spin must occupy each equal energy orbital before additional electrons with opposite spins can occupy the same orbital

44. Rule 1: Aufbau Principle ·Each electron occupies the lowest energy orbital available This means that you must first be able to determine the order of energy levels, sub-levels and the orbitals within those sub-levels Energy levels move in the order of periods down the periodic table ·Order: 1, 2, 3, 4, 5, 6 and 7 Sub-levels are in the order of s, p, d, and f and how they appear on the periodic table 1s 2p 2s 2p 3s 3p 3d 4p 4s 5p 4d 5s 6s 5d 6p 7s 6d 7p 4f 5f ·Notice that the d-level and f-level are "behind" the s and p-level

45. Rule 2: Pauli Exclusion Principle ·A maximum of two electrons can occupy a single orbital, but only if the electrons have opposite spins Remember that the sub-levels contain orbitals. S-sublevel contains 1 orbital, p-sublevel contains 3 orbitals, d-sublevel contains 5 orbitals, and f-sublevel contains 7 orbitals ·s-sublevel can only hold 2 electrons ·p-sublevel can only hold 6 electrons ·d-sublevel can only hold 10 electrons ·f-sublevel can only hold 14 electrons This means that an energy level that contains ALL sublevels can only contain 32 electrons all together

46. 6 1 2 2 1 3 4 5 p 6 10 1 2 4 5 7 9 3 8 S d 2 1 8 3 13 14 4 6 9 10 11 12 5 7 f

47. Lets say we want to write the electron config for Nitrogen 6 1 2 2 1 3 4 5 p 6 10 1 2 4 5 7 9 3 8 S d 1 2 8 3 13 14 4 6 9 10 11 12 5 7 f

48. Practice: Write the electron diagram for the following elements: Cobalt Phosphorus Oxygen Argon Calcium Silver

49. Rule 3: Hund's rule ·Single electrons with the same spin must occupy each equal energy orbital before additional electrons with opposite spins can occupy the same orbital Lets look at our electrons configuration for Nitrogen 1s22s22p3 1s 2s 2p This is called an orbital diagram

50. Practice: Write the orbital diagram for the following elements: Manganese Potassium Gallium Neon