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Atomic Structure and Periodicity Chapter 7. Electromagnetic Radiation. The Nature of Light. Wavelength ( ) is the distance between two consecutive peaks.

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Presentation on theme: "Atomic Structure and Periodicity Chapter 7. Electromagnetic Radiation. The Nature of Light. Wavelength ( ) is the distance between two consecutive peaks."— Presentation transcript:

1 Atomic Structure and Periodicity Chapter 7

2 Electromagnetic Radiation. The Nature of Light. Wavelength ( ) is the distance between two consecutive peaks or troughs in a wave. Frequency ( ) is the number of waves (cycles) per second that pass a given point in space.

3 Wavelength and frequency are inversely proportional

4 EM rays All types of electromagnetic radiation (waves), travel at the speed of light (c). The speed of light (c) = x m/s.

5 The EM spectrum:

6 . Wavelength and frequency are inversely related. = c where is the wavelength in meters, n is the frequency in cycles per second, and c is the speed of light. In the SI system, frequency has a unit of per second, 1/s or s -1, which is called hertz (Hz).

7 Question Frequency of Electromagnetic Radiation. The brilliant red colors seen in fireworks are because of the emission of light with wavelengths around 650 nm when strontium salts such as Sr(NO 3 ) 2 and SrCO 3 are heated. Calculate the frequency of red light of wavelength 6.50x10 2 nm

8 B EM Waves Some Distinguishing Properties of Wave. Refraction is the bending of a light wave (a change in angle) when it strikes a boundary. Diffraction is the bending of a light wave around a boundary. Light must be a wave because it possesses these properties.

9 The Nature of Matter Light vs. Matter Blackbody Radiation. When a solid (specifically a particle) is heated, it emits light waves. Example: Stove top burner glowing when heated. German Max Planck ( ) said that a hot, glowing object emits or absorbs a specific amount of energy to produce this.

10 Pickle Light

11 Photons The energy that comes out is in photons E photon = h where h is Planck’s constant (6.626 x J s), is the frequency of the electromagnetic spectrum absorbed or emitted (Hz)

12 The Energy of a Photon (a packet of energy) The blue color in fireworks is often achieved by heating copper(I) chloride (CuCl) to about 1200 o C. Then the compound emits blue light having a wavelength of 450 nm. What is the photon energy that is emitted at 4.50 x 10 2 nm by CuCl?

13 This total energy leaving an atom is quantized, or lost or gained only in integer multiples of h (the energy of a photon).  E = n h  E is the change in energy for a system, and n is a whole-number integer (1, 2, 3,...). This implies that light energy of matter is not continuous (like a rainbow), but that it is absorbed or emitted only at specific quantized energy states.

14 The Photoelectric Effect. When light strikes a monochromatic plate, an electrical current flows (like a solar calculator). Light must transfer momentum to matter, like particles do; light is behaving like a particle!

15 Copyright © Cengage Learning. All rights reserved15 The Photoelectric Effect

16 Confusing feature of the Photoelectric Effect. A threshold (minimum) frequency is required to knock an electron free from a metal. Wave theory associates the light’s energy with the wave amplitude (intensity), not its frequency (color), so if light is a wave then an electron would be knocked free when the metal absorbs enough energy from any color of light. However, that is not the case.

17 Photoelectric effect Current flows the moment that light of high enough frequency shines on the metal, regardless of its intensity. The wave theory predicted that in dim light there should be a time lag before current flowed, while the electrons absorbed enough energy to break free. However, that doesn’t happen.

18 Albert Einstein’s Photon Theory Einstein proposed that radiation is particulate (made of particles not waves), occurring as quanta of electromagnetic energy (packets), later called photons.

19 Einstein solved the mysteries of the Photoelectric Effect A beam of light is composed of large numbers of photons. Light intensity (brightness) is related to the number of photons emitted per unit of time, not the energy of the individual photon. One electron is freed from the metal when one photon of a certain minimum energy (frequency) is absorbed.

20 Notice this graphic includes increasing energy

21 Expulsion of electron An electron is freed the moment it absorbs a photon of enough energy (frequency), not when it gradually accumulates energy from many photons of lower energy. The amount is different from atom to atom.

22 Arthur Compton In 1922, Arthur Compton performed experiments involving collisions of X rays and electrons that showed that photons do exhibit an apparent mass, a property of matter! Which would make like a particle. However, that is only if moving at relativistic speeds (near the speed of light). A photon would have no rest mass.

23 Summary Energy is quantized. It can occur only in discrete units called quanta. Electromagnetic radiation, which was previously thought to exhibit only wave properties, seems to show certain characteristics of particulate matter as well. This is wave-particle duality or the dual nature of light.

24 Wave Particle Duality This means sometimes light or electrons is looked at as waves, sometimes it is looked at as particles. It depends on the situation which is better suited. There is not an exact answer of what it actually is. This is uncertainty and quantum mechanics

25 Wave-Particle Duality of Matter and Energy. From Einstein, we have the following E = mc 2 This equation relates the energy mass. Instead of looking at energy and matter as different things, Einstein saw them as two sides of the same coin. Energy, such as light, can “condense” into matter, and matter can convert into energy.

26 Louis de Broglie French physicist Louis de Broglie ( ) proposed in 1923 that if waves of energy have some properties of particles, perhaps particles of matter have some properties of waves. This is similar to a guitar string (a particle) producing sound (a wave.)

27 de Broglie wave equation λ = h/mv where λ is wavelength, h is Planck's constant, m is the mass of a particle, moving at a velocity v. de Broglie suggested that particles can exhibit properties of waves.

28 Calculations of Wavelengths Compare the wavelength for an electron (mass = 9.11 x kg) traveling at a speed of 1.0 x 10 7 m/s with that of a baseball (mass = 0.14 kg) traveling 45 m/s (101 mph).

29 Compare the wavelength’s of the baseball and electron The electron’s wavelength is 7.3x m. The baseball’s wavelength is 1.1x m. The wavelength is the distance for it to complete one wave. For the baseball it is incredibly tiny. So much so that the waving isn’t detectable. For the electron the wavelength is tiny, but on an atomic scale it is measurable. It is similar to the distance between atoms in a crystal structure

30 Cont. Waves diffract (spread out and exhibit interference patterns) through openings similar to their wavelength. In 1927, Davisson and Germer verified De Broglie’s concept when an electron (particle) was seen to exhibit wave properties of interference and diffraction.

31 The Spectrum of light The spectrum (light emitted) was thought to be continuous, similar to a rainbow

32 The Atomic Spectrum of Hydrogen. When the emission spectrum of hydrogen in the visible region is passed through a prism, only a few lines are seen.

33 Line Spectrum These lines correspond to discrete wavelengths of specific (quantized) energy. Only certain energies are allowed for the electron in the hydrogen atom. The hydrogen emission spectrum is called a line spectrum.

34 Bohr model Niels Bohr ( ) proposed in 1913 that the electron in a hydrogen atom moves around the nucleus only in certain allowed circular orbits.

35 Energy levels The atoms has stationary states, called energy levels, of specific energy around the nucleus. Electrons can move to other energy levels by absorbing (jumping to higher energy levels) or emitting (jumping to lower energy levels) photons of specific (quantized) energy. Energy levels farther from the nucleus more “unstable” and therefore more “energetic.”

36 As we will see, the closer to the nucleus an energy level is, the more stable it is and the less energetic it is. The lowest (first) energy level of an atom is called ground state. This model only works for one-electron atoms!

37 The most important equation to come from Bohr’s model is the expression for the energy levels available to the electron in the hydrogen atom: E = x J (Z 2 /n 2 ) in which n is an integer (the larger the value of n, the larger is the orbit radius) and Z is the nuclear charge.

38 Bohr was able to calculate hydrogen atom energy levels that exactly matched the values obtained by experiment. The negative sign means that the energy of the electron bound to the nucleus is lower than it would be if the electron were at an infinite distance from the nucleus (n = infinity), where there is no interaction and the energy is zero.

39 Again, the closer to the nucleus, the more stable, and the less energetic (less meaning a negative energy value). Electrons moving from one energy level to another.

40 Electronic Transitions in the Bohr Model for the Hydrogen Atom b) An Orbit-Transition Diagram, Which Accounts for the Experimental Spectrum

41 Equation For a single electron transition from one energy level to another: ΔE = change in energy of the atom (energy of the emitted photon) n final = integer; final distance from the nucleus n initial = integer; initial distance from the nucleus 41

42 Energy Quantized in Hydrogen Calculate the energy required to excite the hydrogen electron from level n = 1 to level n = 2. Calculate the wavelength of light that must be absorbed by a hydrogen atom in its ground state to reach this excited state.

43 Electron Energies Calculate the energy required to remove the electron from a hydrogen atom in its ground state.

44 Although Bohr’s model fits the energy levels for hydrogen, it is a fundamentally incorrect model for the hydrogen atom, mainly because electrons do not travel in circular orbits.

45 The Quantum-Mechanical Model of the Atom The Schrodinger equation is an extremely complex equation used to describe the 3-D quantum mechanical model of the hydrogen atom. The hydrogen electron is visualized as a standing wave (a stationary wave as one found on a musical instrument) around the nucleus. The circumference of a particular orbit would have to correspond to a whole number of wavelengths.

46 This is consistent with the fact that only certain electron energies are shown to exist.

47 The Heisenberg Uncertainty Principle The Heisenberg uncertainty principle states that it is impossible to know simultaneously the exact position and velocity of a particle. Δx = uncertainty in a particle’s position Δ(mν) = uncertainty in a particle’s momentum h = Planck’s constant

48 The Physical Meaning of a Wave Function The square of the function indicates the probability of finding an electron near a particular point in space. The probability distribution, or electron density diagram, can be plotted to give the most probable region where an electron may be located.

49 orbitals The definition most often used by chemists to describe the size of the hydrogen 1s orbital is the radius of the sphere that encloses 90% of the total electron distribution.

50 Quantum numbers A quantum number describes various properties of an atomic orbital or probable electron location. Types. The principle quantum number (n) refers to the energy level or integral values: 1, 2, 3, etc. The angular momentum (azimuthal) quantum number ( l ) is related to the shape of atomic orbitals. The integral values are from 0 to n - 1.

51 Cont. l = 0 is called s; l = 1 is called p; l = 2 is called d; l = 2 is called f. The magnetic quantum number (m l ) is related to the orientation of the orbital in space. The integral values are between l and - l, including zero.

52 Quantum Numbers for the First Four Levels of Orbitals in the Hydrogen Atom.

53 Electron Spin Although the first three quantum numbers describe atomic orbitals, one more is needed to describe a specific electron. This is called spin, and is not a property of the orbital. The spin quantum number (m s ) indicates the “direction” an electron spins and can have one of two possible values, + ½ or - ½

54 Pauli Exclusion Principle American Wolfgang Pauli stated that no two electrons in the same atom can have the same set of four quantum numbers. This is the Pauli exclusion principle An atomic orbital can hold a maximum of two electrons and they must have opposing spins. The electrons of an atom in its ground state occupy the orbitals of lowest energy.

55 Electrostatic Effects In one electron systems, the only electrostatic force is the nucleus- electron attractive force. Many-electron systems (polyelectronic atoms). Nucleus-electron attractions. Electron-electron repulsions.

56 Polyelectronic Atoms For polyelectronic atoms, one major consequence of having attractions and repulsions is the splitting of energy levels into sublevels of differing energies. The energy of an orbital in a many- electron atom depends on its n value (size) and secondly on its l value (shape).

57 This means that energy levels may split into s, p, d, and f sublevels, depending on the n value.

58 Copyright © Cengage Learning. All rights reserved58 Orbital Energies

59 The Effects of Electrostatic Interactions on Orbital Energies. Recall that, by definition, an atom’s energy has a negative value. A more stable orbital has a larger negative energy than a less stable one. Stronger attractions make the orbital lower in energy (larger negative number). Repulsions make the system higher in energy (smaller negative number).

60 Conclusions A greater nuclear charge (Z) lowers orbital energies. Why? Electron-electron repulsions raise orbital energy. Why? Electrons in outer orbitals (higher n) are higher in energy. Why?

61 Aufbau principle As protons are added to the nucleus to build up a proton, electron are added to the lowest energy level hydrogen-like orbital. Hence we fill up 1s then 2s, 2p, 3s, 3p, 4s, 3d, 4p… Why did we you fill 4s before 3d? These aren’t hydrogen orbitals (unless you are talking about hydrogen). 4s is simply a lower energy state than 3d.

62 Aufbau exceptions You see the first exceptions to the Aufbau principle in the 3d orbital. Chromium is [Ar]3d 5 4s 1 As opposed to the expected [Ar]3d 4 4s 2 It then goes back to “normal”, until copper. Copper is [Ar]3d 10 4s 1 There are more examples is the d orbitals below, and in the f orbitals

63 Effective Nuclear Charge (Z eff ) Z eff is the net positive charge experienced by a specific electron in a multi electron atom. As electrons get further away, they will be “shielded” from the nucleus by other electrons. What is the period trend, with explanation?

64 Ionization Energy The energy required to remove an electron. What type of ion is produced? Is the process endothermic or exothermic? What is the group trend, with explanation? What is the period trend, with explanation?

65 Exceptions Normally moving left to right across a period will increase in ionization energy because it is being held in by a stronger nucleus. However there is an observed dip in ionization energy going from Be to B. Why? It is harder to remove an electron from 2p than 3s, these aren’t hydrogen orbitals. This trend continues for Mg-Al, then you see the trend for Zn-Ga, and Cd-In, and Hg-Tl

66 Another exception There is another between group 15 and 16 (nitrogen group to oxygen group). This exception is because in group 15 you have one electron in each of the 3 p orbitals, group 16 has a 4 th electron that is easier to remove.

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68 Successive Ionization Energies Ionization energy doesn’t have to stop at one electron, you can remove multiple electrons, to get a 2 nd ionization energy I 2, 3 rd ionization energy I 3 … Why are second ionization energies more than first ionization energies? Electrons are more easily removed until what occurs?

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70 Trends in Ionization Energy The first ionization energy for phosphorus is 1060 kJ/mol, and that for sulfur is 1005 kJ/mol. Why?

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72 Consider atoms with the following electron configurations: 1s 2 2s 2 2p 6 1s 2 2s 2 2p 6 3s 1 1s 2 2s 2 2p 6 3s 2 Which atom has the largest first ionization energy, and which one has the smallest second ionization energy? Explain your choices.

73 Atomic Radius ~relative size of an atom. As you move down a period you are increasing the number of energy levels so the atom is getting larger. As you move left to right across a period, a more positive nucleus pulls everything in closer, making the atom smaller. It increases down and to the left.

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75 Electron Affinity Electron affinity is the change in energy when an atom captures a bonding electron. Energy is normally released, so the number is normally reported as a negative. What is the group trend, with explanation? What is the period trend, with explanation?


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