Chemistry Ch 4

1. Chemistry Ch 4 Arrangement of Electrons in Atoms

2. Rutherford’s Model • Gold Foil Experiment • Discovered the nucleus • Did not explain where the electrons were in an atom • Why were they not attracted to the protons in the nucleus?

3. Background Info • Electromagnetic Radiation are types of energy. • We describe these as waves. Only a portion of these waves are visible to us (the visible light waves). • Each type of wave has different wave characteristics: frequency, wavelength, and the amount of energy it contains.

4. Relationship between light and electrons • The electromagnetic spectrum includes all types of electromagnetic radiation that behaves as waves. • Light can behave as a wave (as in the spectrum) and as a particle of matter like a marble. • Speed of all electromagnetic radiation is 3.0 X 108 m/s.

5. The Electromagnetic Spectrum

7. Wave Characteristics • Wavelength (λ) -the distance between two points on a wave (measured in nm) • Frequency (v) –the number of waves that pass in a given point in one second • The speed of light (c) – is a constant • C= λ v • Wavelength and Frequency are inverse of each other (opposite).

8. Converting Wavelength Units • Wavelength is measured in nm. • Speed is measured in m/s. • They must both have the same unit so we must convert nm to m to use it in the equation. • You broke your big toe!  The x ray they take of  toe uses waves that have a length      4.0 X 10-7m. ( 1 meter = 1 X 109 nm) • What is the wavelength in nm? (l = 400 nm)

9. Frequency and Wavelength Problems • Calculate n for a l = 700 nm. • (Red 700nmn = 4.3 x 1014 /s    or  4.3 x 1014 Hz) • A purple light has a frequency of 7.42 x 1014 Hz.      What is its wavelength in nm? • (l = 404 nm)

10. Lets look at a video to see what we are going to be learning and what the scientists were investigating… • Video of electron behavior as waves and particles

11. Photoelectric Effect • Experiments by Einstein and others in the 1900s tried to explain the interactions between light and matter that were not explained with the wave theory • Their research led them to discover the dual wave particle nature. How electromagnetic radiation behaves as waves and as particles.

12. The evidence for this was the Photoelectric Effect experiment, which explained how light, that is usually thought of as a wave, can also behave like a particle of matter. • Lois de Broglie wondered if electrons (matter), normally thought of as a particle, maybe have some wave properties too.

13. Photo of Photoelectric Effect

14. The wave theory predicted that light of any frequency could supply enough energy to eject an electron from its position. • However, no electrons were emitted if a light’s frequency was below a certain minimum, regardless of how long the light was shone.

15. Max Planck suggested that an object emits energy in small, specific amounts, called quanta. • A quantum is the minimum quantity of energy that can be lost or gained by an atom. • E= hv • h = Planck’s constant 6.626 X 10-34 Js

16. Photons are the “particles of light” that carry a certain amount of energy. • The energy of a photon depends on the frequency of the wave.

17. In order for an electron to be ejected from a metal surface, the electron must be struck by a single photon with the minimum energy required to knock the electron lose. (supports the particle theory) • Because E = hv, the minimum energy needed corresponds to the frequency

18. Energy Problem • What is the energy of a photon whose frequency is 3.0 X 1012 Hz? E = hv Where h= 6.626 x 10-34J/Hz          E =[6.626 x 3.0] 10 (-34+12) J 1.99 X 10-21 J

19. An observation was that different metals required different amounts of energy or frequencies to exhibit the photoelectric effect. (different metals are different elements with different numbers of electrons) • So what did all this mean for where the electrons were in an atom?

20. It was concluded that electrons exist in specific energy levels in an atom: • Ground state = lowest energy state of an atom • Excited state = the highest energy state • When atoms are excited by energy (heat), they emit energy in the form of light.

21. Classical theory predicted that atoms would be excited by whatever amount of energy that was added to them. (there would be a continuous spectrum of frequencies given off-like a prism) • However, when current was passed through Hydrogen gas, a series of very specific frequencies were emitted and only certain colors were seen (line emission spectrum)

22. Hydrogen Line Emission Spectrum

23. This suggested that the electrons of an atom exist in very specific energy states. • So, Bohr put all this information together in his model of an atom.

24. Bohr Model • Orbits-electrons can only circle the nucleus in allowed paths • Each orbit has a fixed amount of energy • Closest to the nucleus has the least amount of energy (ground) • Farther from the nucleus has more energy (excited)

25. Or in other words, When an excited electron returns to its ground state, it gives off the energy (a photon) in the form of electromagnetic radiation (sometimes visible light).

26. From E2 to E1, the electron will gain or lose energy? • From E1 to E5, the electron will gain or lose energy?

27. Why do different atoms emit different light? Each atom is unique and contains its own unique electron structure in the different energy levels.

28. How does the emission of light relate to the electron structure? • Since each atom is unique in its electron structure with differing levels of energy, the transitions between those levels will be unique to each atom. • Electrons are in certain energy levels. When electrons give off light, they emit energy, and move to a lower level closer to the nucleus.

29. Balmer Series

30. Emission Spectrum

31. Some electron transitions result in energies and wavelengths within the visible light spectrum so we can see them (400-750 nm). • However, there are many transitions that we cannot see (radio waves, x-rays, gamma rays)

32. de Broglie concluded that since an electron is so small but its speed is so great, it could orbit a nucleus millions of times in 1 second! (He used algebraic methods and the equations of Einstein, Planck, and the speed of a wave to figure) • So, how could we possibly know where an electron is in an atom?

33. Schrödinger suggested that since electrons can be thought of like waves, they may be like standing waves outside the nucleus. • Only a certain number of waves can exist between the nucleus and a certain point. • This fits with Bohr’s idea of energy levels in an atom.

34. Ch. 4-2 The Quantum Model • Light can behave as waves and particles • Louis De Broglie investigated that electrons also behave like waves because: • They are confined to a specific frequency • Diffraction-bending of a wave as it passes by something • Interference-waves overlapping Fig. 4-10

35. Heisenberg Uncertainty Principle • Heisenberg: It is impossible to simultaneously determine the location and velocity of an electron or particle • Schrodinger’s theory (that there can only be so many wavelengths of energy in a certain level) led to the development of the quantum theory • Quantum theory-describes the wave patterns of electrons mathematically

36. As a result of the Schrodinger equation and Heisenberg’s Principle, the location of an electron is only its probable location around a nucleus • Orbital-3D region around the nucleus that describes its probable location • Fig. 4-11

37. Quantum Numbers • Specify the properties of atomic orbitals and properties of electrons in orbitals • There are 4 quantum numbers for each electron: • Principle Quantum Number • Angular Momentum Quantum Number • Magnetic Quantum Number • Spin Quantum Number

38. NO two electrons have the same 4 quantum numbers • Similar to a zip code-no 2 cities have the same zip code

39. Principal Quantum Number (n)-indicates the main energy level of the electron Ex: n= 1, 2, 3….. Also indicates how many sublevels there may be for a main energy level • Angular Momentum Quantum Number (l)-indicates the shape of the sublevel or oribital S-sphere P-dumbbell shaped D-3D shape F- Too complex Pg. 102

40. Magnetic Quantum Number (m) –orientation of an orbital around the nucleus S-1 P-3 D-5 F-7

41. Each orbital can hold 2 electrons • So, total electrons for each S = 2 (1 X 2) P = 6 (3 X 2) D = 10 (5 X 2) F = 14 (7 X 2)

42. How many electrons in each energy level? • Use 2n2 to figure out how many electrons can be in each energy level • Ex: for energy level 5, n=5 • So 2 (5)2 = 50 electrons

43. Spin Quantum Number – indicates the direction the electron will spin in orbit * has only 2 possible values for the spin, either +1/2 or -1/2 * The two electrons in each orbital have to have opposite spins

44. Ch. 4-3 Electron Configurations • The quantum model tells us more than the Bohr model of the atom because it tells us where the electrons are located • Electron configurations-the arrangement of electrons in an atom

45. Aufbau principle-an electron occupies the lowest-energy orbital that can receive it • Fig. 4-16 We always start at 1s and work up to 2s, 2p, etc.

46. Pauli exclusion principle-no two electrons in the same atom can have the same set of four quantum numbers • Ex: zip code

47. Hund’s Rule-orbitals of equal energy are each occupied by one electron before any orbital is occupied by a second electron. All the first electrons must all have the same spin

48. Electron Configurations • Electrons fill orbitals (s, p, d, f) and energy levels (1, 2, 3…) in a certain order according to energy:

49. Remember, s orbitals can hold 2 electrons p orbitals can hold 6 electrons d orbitals can hold 10 electrons f orbitals can hold 14 electrons

50. Electron Configurations 1. Electron Configuration Notation Ex: He- 1s2 Ex: Na-1s22s22p63s1 2. Orbital Notation Ex: He-_↑↓_ Ex: Na- (1s)_↑↓_ (2s)_↑↓_ (2p)_↑↓_ _↑↓_ _↑↓_ (3s) _↑_ 3. Noble Gas Configuration Ex: Li- [He] 2s1 Ex: Na-[Ne] 3s1