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Light Quantized energy Quantum theory and the atom.

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Presentation on theme: "Light Quantized energy Quantum theory and the atom."— Presentation transcript:


2 Light Quantized energy Quantum theory and the atom

3 Electromagnetic radiation

4 Wave nature of light Electromagnetic spectrum Equations

5 All light is composed of an electric component and a magnetic component; thus the term electromagnetic radiation

6 There are a few terms that we use to characterize light waves – Wavelength (  Units are meters (m) – Frequency (  Units are Hertz (Hz = seconds -1 ) – Amplitude

7 There are a few terms that we use to characterize light waves

8 Although light waves can have different wavelengths and frequencies, they all travel at the same speed; the speed of light 3.0 x 10 8 m/s

9 Remember when we learned about radiation energy we looked at the electromagnetic spectrum Now we will apply the ideas of wavelength and frequency to the electromagnetic spectrum



12 Visible light

13 The frequency and wavelength of light are indirectly proportional As one increases, the other decreases

14 The frequency and wavelength of light are indirectly proportional When frequency and wavelength are multiplied together, they always equal the speed of light Speed of light (c) c = 

15 If we know either the wavelength or the frequency of light, we can calculate the other one by rearranging the equation Ex: If we know the wavelength, = c/ If we know the frequency,  = c/


17 Particle nature of light Emission and absorption spectra

18 Certain observations about light interacting with matter were not able to be described by the wave properties of light – Heated objects emit light at specific frequencies at a given temperature – When light of a certain frequency is shined on some metals, electrons are emitted

19 Max Planck, a German physicist, discovered that matter can only gain or lose energy in small specific amounts called quanta A quantum is the minimum amount of energy that can be gained or lost by an atom

20 The energy (E) that is emitted by hot objects is related to the frequency ( ) of the emitted radiation. They are related by a number called Planck’s constant (h) h = e -34 J x s E = h

21 h = e -34 J x s Energy is always released in multiples of h  h  h  h  ) E = h

22 When light of a certain frequency is shined on a metal surface, electrons are ejected from the metal. This phenomenon is known as the photoelectric effect

23 Photoelectric effect Packets of light energy called photons Energy of light is transferred to the electron increasing the electron’s kinetic energy

24 Photoelectric effect

25 Atomic emission spectra

26 – Electricity passing through the neon gas in the glass tube – Neon atoms absorb that energy and become “excited” – The “excited” atoms release energy as light as they return back to their “ground” state – The atomic emission spectra for an element is the set of frequencies of the light emitted by the atoms as they return to their “ground” state

27 Atomic emission spectra – If the light emitted by an element is passed through a prism, the frequencies of the emitted light can be determined

28 Atomic emission spectra

29 Continuous emission spectra Atomic emission spectra

30 Each element has its own unique emission spectra

31 Absorption spectra

32 Continuous vs. emission vs. absorption “Excited” gas “Ground state” gas

33 We don’t see the colors that are absorbed, only those that are reflected



36 Why do you think this pattern occurs?


38 Bohr’s model of the atom Electrons as waves Heisenberg uncertainty principle

39 Bohr used Planck’s idea of quantized energy and applied it to the atom He proposed that electrons orbit nuclei only at specific distances from the nucleus

40 Bohr atomic model

41 The orbital closest to the nucleus corresponds to the ground state The orbitals further away from the nucleus are excited states


43 Energy associated with electron orbital transitions

44  E = E f – E i If E is absorbed, E f > E i  E is positive If E is emitted, E f < E i  E is negative

45 Energy absorbed or emitted can be in frequencies other than just visible light

46 De Broglie proposed that electrons moving around the nucleus had wave-like behavior. The wavelength associated with an electron depends on the mass and velocity of the electron  = h / m v

47 The idea of particle – wave duality applies to all matter, not just light and electrons The mass of objects we can see are so large and the wavelengths are so small that we cannot see this effect = h / m v

48 Heisenberg uncertainty principle (for atomic scale particles) It is impossible to know both the position and the velocity of a particle at the same time

49 Anything we do to determine the location or velocity of an electron moves it from its original location and changes its velocity We can know one or the other but not both We talk about the probability for an electron to occupy a certain region around the nucleus (so the fixed orbital proposed in the Bohr model are impossible)

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