PreAP Chemistry Chapter 4 Notes

1. PreAP Chemistry Chapter 4 Notes

2. Section 4.1 The Development of a New Atomic Model Previously, Rutherford reshaped our thoughts of the atom by showing the protons were located in the nucleus of the atom, but he could not model for us where the electrons were, other than outside the nucleus somewhere. Fortunately, studies into the properties of light and the effects of light on matter soon gave clues to where electrons actually are.

3. Light is a small part of all the radiation (something that spreads from a source) called electromagnetic radiation. Electromagnetic radiation is energy in the form of waves (of electric and magnetic fields). Electromagnetic radiation includes radio waves, microwaves, infrared, visible light, X-rays, and Gamma rays. All these together are considered the Electromagnetic Spectrum.

5. As all the forms of electromagnetic radiation are waves, they all have similar properties. • All electromagnetic radiation travels at the speedoflight (c), 299,792,458 m/s (3 x 108) in a vacuum

7. •The crest is the top of the waves, the trough is the bottom of the waves, and the amplitude is a measurement from the rest or zero line to a crest or trough

9. •The wavelength (λ – lambda) is the distance between successive crests/troughs and is measured in meters (often nm = 10-9 m)

10. •The wavelength (λ – lambda) is the distance between successive crests/troughs and is measured in meters (often nm = 10-9 m) •The frequency (ν – nu) is the number of waves that pass a point in one second and is measured in (per second – can be written as s-1) or Hz (Hertz)

11. Wavelength is the distance between two crests/troughs λ (lamda) is the symbol of wavelength and m is the unit Frequency is the number of crests passing through a point per second

12. How many hertz is the first wave? How many hertz is the second wave?

13. The speed of a wave is directly proportional to the wavelength and the frequency; c = λν is the formula c ν λ

14. Example. A certain violet light has a wavelength of 413 nm. What is the frequency of the light?

15. Example. A certain violet light has a wavelength of 413 nm. What is the frequency of the light?

16. Example. A certain violet light has a wavelength of 413 nm. What is the frequency of the light?

17. Example. A certain violet light has a wavelength of 413 nm. What is the frequency of the light?

18. Unfortunately, thinking of light as waves lead to a problem. It was noticed that if light strikes a metal, then sometimes it could cause electrons to be emitted (leave the atoms entirely – like in a solar panel); called the photoelectric effect. If light was a wave, then all amounts of light energy should cause this to happen, but this was not the case. It always took some minimum amount of energy to get the electrons to be emitted.

20. This lead Max Planck to theorize that light must carry energy in basic minimum amounts that he called quanta. Like a delivery person cannot correctly deliver half a box, the electrons in atoms cannot gain a fraction of a quantum of energy (it has to be in whole numbers).

21. E h ν He proposed that this energy was directly proportional to the frequency of the electromagnetic radiation and a constant, now called Planck’s constant. E = h ν E = energy in Joules (J) h = Planck’s constant = 6.63 × 10-34 Js ν = frequency in Hz or 1/s

22. Example. What is the energy content of one quantum of the light with a wavelength of 413 nm?

23. Example. What is the energy content of one quantum of the light with a wavelength of 413 nm? Note: wavelength is not in the energy equation, but frequency is. So first, you must solve for the frequency. As seen in the earlier example, a wavelength of 413 nm gives a ν = 7.26 × 1014 Hz.

24. Example. What is the energy content of one quantum of the light with a wavelength of 413 nm? ν = 7.26 × 1014 Hz E = h × ν E = 6.63 × 10-34 Js × 7.26 × 1014 1/s

25. Example. What is the energy content of one quantum of the light with a wavelength of 413 nm? ν = 7.26 × 1014 Hz E = h × ν E = 6.63 × 10-34 Js × 7.26 × 1014 1/s

26. Example. What is the energy content of one quantum of the light with a wavelength of 413 nm? ν = 7.26 × 1014 Hz E = h × ν E = 6.63 × 10-34 Js × 7.26 × 1014 1/s E = 4.81 × 10-19 J

27. In 1905 Einstein used Plancks work to propose that electromagnetic radiation had a dual wave-particle nature. As a particle, electromagnetic radiation carries a quantum of energy of energy, has no mass, and is called a photon.

28. So to get an electron to emit from a metal, it must be struck with a photon having quantum energy big enough, or nothing will happen. Each metal requires a different quantum energy, thus each metal can be identified by the frequency of light needed to emit electron.

29. This idea was expanded upon to develop an idea of where the electrons were in an atom. It was found that low pressure gases could be trapped in a tube and electrified, and would then glow a color particular to the gas inside.

30. Furthermore this light could be passed into a prism, and instead of getting the entire spectrum (rainbow) of colors, only certain wavelengths of light would be seen as small bars of color, called a line-emission spectrum.

31. This would indicated that the electrons in an atom were only absorbing specific amounts of energy from the electricity, causing the electrons to move from their ground state (normal position close to the nucleus) to an excited state (higher energy position further away from the nucleus). The electrons do not stay in the excited state for long and fall back to their ground state, losing the energy equal to what they gained.

32. Quanta Video

33. Niels Bohr used this to develop a model of the atom where the electrons could only be in certain, specific energy level (n) orbits around the nucleus. Just as you cannot go up half a rung on a ladder, the electron could not go up a partial energy level. The electrons gained or lost enough energy to move a whole number amount of energy levels (n) away from or closer to the nucleus, or it did not move.

34. He calculated the amount of energy needed for an electron of hydrogen to move between each energy level (n) (which was not constant) and his calculations agreed with experimental results.

35. The Balmer series of hydrogen spectral lines refer to the four lines seen in the visible light region (the four colored bars). If the electron was excited to energy level (n) 6, 5, 4, or 3 and fell to energy level (n) 2, the resulting energy given off would have a frequency in the visible region of electromagnetic radiation. (One line for dropping from 6 to 2, one for 5 to 2, one for 4 to 2, and one for 3 to 2).

36. Line Spectra Video

37. However, there are other possibilities. If the electrons drop from n=6, 5, or 4 to n=3, then the energy given off is not big enough to be seen as it is in the infrared region. These three lines in the infrared region are referred to as the Paschen series. If the electrons drop to n=1, then the five lines given off are too high in energy to be seen, as they are in the ultraviolet region. These lines are referred to as the Lyman series.

40. Model of Atom Review:

41. Thomson’s Plum Pudding Model – the atom is a ball of evenly spread positive stuff with random negative particles (electrons).

42. Thomson’s Plum Pudding Model – the atom is a ball of evenly spread positive stuff with random negative particles (electrons). • Rutherford’s Nuclear Model – the atom has a central nucleus containing the positive particles (protons) with the electrons outside.

43. Thomson’s Plum Pudding Model – the atom is a ball of evenly spread positive stuff with random negative particles (electrons). • Rutherford’s Nuclear Model – the atom has a central nucleus containing the positive particles (protons) with the electrons outside. • Bohr’s Orbital Model – The electrons circle the nucleus in specific energy orbits, like the planets orbit the sun. Unfortunately this only works for atoms with one electron…

44. Quantum Mechanical Model – electrons are found in specific regions around the nucleus, but the exact location of the electrons inside the regions cannot be determined

45. Atom Models Video

46. The quantum mechanical model of the atom is built on the ideas and calculations of several scientists. • Louis deBroglie suggested a way to show that a particle could have wave like behavior with the equation: h = Planck’s constant m = mass of particle v = velocity of particle

47. • The Heisenberg Uncertainty principle states that it is impossible to know both the velocity and the location of an electron at the same time. If the position was known, then there is no way to know where it has moved to, and if the velocity is known, then there is no way to know where it was.

48. • Schrödinger developed wave-based equations that form the basis of the current Quantum theory, which mathematically describe the probably location of electrons, often referred to as an electron cloud. The electrons clouds describe areas around the nucleus with a 90% chance of finding the electron inside. Solving the equations has lead to QuantumNumbers, which will be studied later.

49. The quantum mechanical model starts with a Principal Quantum Number (n), which is the basic energy level of an electron, and often matches the period number. Possible values (currently) are 1-7.

50. The quantum mechanical model starts with a Principal Quantum Number (n), which is the basic energy level of an electron, and often matches the period number. Possible values (currently) are 1-7. Inside the principal quantum energy level are sublevels that correspond to different cloud shapes. The sublevels are designated as s (sharp), p (principal), d (diffuse), and f (fundamental).