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1 Unit 5 Chapter 13 Electrons in the Atom Electrons in the Atom (Chapter 13) & The Periodic Table/Trends (Chapter 14) Niels Bohr s Model Recall the Evolution of the Atom He had a question: Why don t the electrons fall into the nucleus? The Electrons move like planets around the sun. In circular orbits at different levels. Amounts of energy separate one level from another. Planetary model Bohr s planetary model Energy level of an electron Levels are analogous to the rungs of a ladder An electron cannot exist between energy levels, just like you can t stand between rungs on ladder A Quantum of energy is required to move an electron to the next highest level Bohr s Absorbance & Emission of this Quantum of Energy 1

2 Section 13.3 Physics & Nature of Light, and the Quantum Mechanical Model What s in a Flicker of Light? OBJECTIVES: Calculate the wavelength, frequency, or energy of light, given two of these values. Explain the origin of the atomic emission spectrum of an element. What is Light? The study of light led to the development of the quantum mechanical model. Light is a wave phenomenon, or a kind of electromagnetic radiation. Electromagnetic radiation includes many kinds of waves All waves travel in a vacuum at 3.00 x 10 8 m/s = c Crest Origin Parts of a wave Trough Wavelength Amplitude Origin - the base line of the energy. Crest - high point on a wave Trough - Low point on a wave Amplitude - distance from origin to crest Wavelength - distance from crest to crest Wavelength is abbreviated by the Greek letter lambda = λ Frequency The number of waves that pass a given point per second. Units: cycles/sec or hertz (hz or sec - 1 ) Abbreviated by Greek letter nu = ν c = λν Frequency and wavelength Are inversely related; λ α 1/υ Are inversely related; As one goes up the other goes down. Which is the Stronger Wave? 2

3 Frequency and wavelength Different wavelengths (frequencies) of light are different colors of light. There is a wide variety of frequencies The whole range is called the Electromagnetic spectrum The Electromagnetic Spectrum Gamma Rays X Rays UV IR Microwave TV Radio nm cm meters VISIBLE UVB UVA Near IR nm R - Red O - Orange Y - Yellow G - Green B - Blue I - Indigo V - Violet The Electromagnetic Spectrum UV The Visible Region IR nm Red; 630 nm Orange; 610 nm Yellow; 590 nm Green; 525 nm Blue; 470 nm Indigo; 405 nm Violet; 380 nm The Electromagnetic Spectrum What are the wavelengths of the Visible Colors? According to the Bohr Model Electrons can change orbits, accompanied by the absorption (electrons on the way up) or emission (electrons on the way down) of a photon of a specific color of light. Features of Light Why is the Cup Red? Red λ is reflected All other wavelengths are absorbed. Similarly UV IR nm 3

4 Features of Light When does a sample Absorb a given color? A sample absorbs a given color when that color s frequency is a match to cause the excitement of an electron to a higher state. White Light through a Prism White light is made up of all the colors of the visible spectrum. Passing it through a prism separates it. Visible light represents a continuous spectrum. If the light is not white By heating a gas at low pressure with electricity we can get it to EMIT LIGHT. Passing this light through a prism does something different. Each element gives off its own characteristic colors. Can be used to identify the atom. These are Atomic Emission Spectrum Heating a gas at low pressure These are called discontinuous spectra, or line spectra unique to each element. Each line corresponds to a specific amount of energy being emitted. How are Continuous & Emission Spectra Different? Vs. Continuous spectra are produced by solids, liquids, and dense gases under high pressure Emission spectra are produced by heating a gas at low pressure Recapping the Wave: Frequency, Wavelength, and Energy of Light (c = λυ) Types of Spectra: High Pressure Vs. Low Pressure 4

5 Exciting the Electron of Hydrogen Heat or electricity or light can move the electron up energy levels ( excited ) Exciting the Electron of Hydrogen Heat or electricity or light can move the electron up energy levels ( excited ) Exciting the Electron of Hydrogen As the electron falls back to ground state, it gives the energy back as light Exciting the Electron of Hydrogen The electron can make many allowable transitions to excited states, and each transition up will result in different steps downward, each with a different energy. The Emission Spectrum of Hydrogen using the Bohr Model 1. Lyman Series: UV Emissions down to n = 1; The Longest Transition Exciting the Electron of Hydrogen The electron may fall down in steps, each with a different energy 2. Balmer Series: Visible Emissions down to n = 2 3. Paschen Series: Infrared Emissions down to n = 3; The Shortest Transition 5

6 What The Electrons Do This model is simplified The orbitals also have different energies inside energy levels All the electrons can move around. This is where the Bohr Model falls short. The Bohr Model is insufficient for more complex atoms with more electrons. T 2 T 1 Triplet State Phosphorescence Intersystem Crossing S 0 S 2 S 1 Absorption Singlet State Internal Conversion Fluorescence Revisiting Atomic Emission Spectrum Let s Look at the Atomic Emission Spectra of Various Gases using the Spectrometer Each line of color represents an emission of light from an excited state ground state e - transition. The more electrons in an atom, the more transitions are allowed, and the more complex the emission line spectrum can be. How Can the Wave Theory of Light Explain these line spectra? It Can t!! Wave Theory suggests that the energy change for a wave can be teenie- weenie, or infinitesimally small. How do the lines separate? They should be continuous. Is Light a Wave or a Particle? Max Planck pointed out that Energy is quantized. Light is energy Light must be quantized. These smallest pieces of light are called photons. Light has a dual wave-particle behavior. 6

7 The Dual Nature of Light The Particle (Photon) Nature of light helps explain the Previously Mysterious Photoelectric Effect. In the Photoelectric Effect, metals eject electrons when light shines on them. The Dual Nature of Light Einstein discovers that light can be Quantized via E = hυ In order for the photoelectric effect to occur, the incident light had to reach a threshold frequency, which corresponds to a threshold photon of Energy, via E=hυ. The Dual Nature of Light Einstein discovers that light can be Quantized via E = hυ The threshold frequency provides the electrons with enough energy to escape from the metal. The number of photons with sufficient energy = the number of electrons ejected. Energy and frequency E = h ν E is the energy of the photon ν is the frequency h is Planck s constant h = x Joules x sec. Joule is the metric unit of Energy Frequency, Wavelength, and Energy of Light Problems (E = hυ and c = λυ) What is the wavelength of blue light with a frequency of 8.3 x Hz? λ = c/υ = (3.0x10 8 m/s)/8.3x10 15 sec -1 λ = 3.6x10-8 m What is the frequency of red light with a wavelength of 4.2 x 10-5 m? υ = c/λ = (3.0x10 8 m/s)/4.2x10-5 m υ = 7.14x10 12 sec -1 Frequency, Wavelength, and Energy of Light Problems (E = hυ and c = λυ) What is the energy of a photon of each of the above? 1. Frequency = 8.3x10 15 Hz E = hυ = 6.63x10-34 Jsec x 8.3x10 15 Hz E = 5.5x10-18 J 2. υ = 7.14x10 12 sec -1 E = hυ = 6.63x10-34 Jsec x 7.14x10 12 Hz E = 4.7x10-21 J 7

8 Frequency, Wavelength, and Energy of Light Problems (E = hυ and c = λυ) The Energy for a photon of light is 3.55x10-19 Joules. What is the wavelength of this light, and what is it s color? υ = E/h = 3.55x10-19 J / 6.63x10-34 J-s υ = 5.35x10 14 sec -1 then: λ = c/υ = (3.0x10 8 m/s)/5.35x10 14 sec -1 λ = 5.61x10-7 m = 561 nm (Color Yellow) Frequency, Wavelength, and Energy of Electron Transitions (E = hυ and c = λυ) Calculate the Energy (E), Frequency (υ), and Wavelength (λ) for each of the first three Balmer Series Transitions Wave & Light Worksheets 8

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