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The Rutherford model of the atom was an improvement over previous models, but it was incomplete. J. J. Thomson’s “plum pudding” model, in which electrons.

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Presentation on theme: "The Rutherford model of the atom was an improvement over previous models, but it was incomplete. J. J. Thomson’s “plum pudding” model, in which electrons."— Presentation transcript:

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2 The Rutherford model of the atom was an improvement over previous models, but it was incomplete. J. J. Thomson’s “plum pudding” model, in which electrons are surrounded by a soup of positive charge to balance the electron’s negative charge, like negatively-charged “plums” surrounded by positively charged “pudding”. It did not explain how the atom’s negatively charged electrons are distributed in the space surrounding its positively charged nucleus.

3 In the early twentieth century, a new atomic model evolved as a result of investigations into the absorption and emission of light by matter. The studies revealed a relationship between light and an atom’s electrons. Before 1900, scientists thought light behaved solely as a wave. This belief changed when it was later discovered that light also has particle- like characteristics. A quick review of these wavelike properties follows.

4 Properties of Light Visible light is a kind of electromagnetic radiation (form of energy that exhibits wavelike behavior as it travels through space). Other examples include X-rays, ultraviolet and infrared light, microwaves, and radio waves. The electromagnetic spectrum consists of all types of electromagnetic radiation.

5 Electromagnetic Spectrum

6 All forms of electromagnetic radiation move at a speed of 3.0 x 10 8 m/s through air. Wavelength – λ (m, cm, or nm) and frequency – ν (wave/second) are measureable properties of wave motion. One wave/second is called a hertz (Hz). The relationship between wavelength (λ) and frequency (ν) is c = λν where c = speed of light.

7 Wavelength and Frequency

8 In the early 1900s, scientists conducted two experiments involving interactions of light and matter that could not be explained by the wave theory of light. One experiment involved a phenomenon known as the photoelectric effect.

9 Photoelectric Effect The photoelectric effect refers to the emission of electrons from a metal when light shines on the metal. Light may cause electrons to be emitted from an electrode in a photocell. Long wavelength light does not have enough energy to cause the electron to escape, regardless of its intensity. When light of a shorter wavelength (higher energy) light strikes the electrode, electrons are released. The amount of current produced depends on the intensity of the light and the energy of the escaping electrons depends on the wavelength of the light. Show video clip.

10 The wave theory of light predicted that light of any frequency could supply enough energy to eject an electron. Scientists couldn’t explain why the light had to be of a certain frequency in order for the photoelectric effect to occur. The German physicist Max Planck proposed an explanation for the photoelectric effect. He proposed that a hot object does not emit electromagnetic radiation continuously, as would be expected if the energy emitted were in the form of waves.

11 Max Planck proposed that objects emit energy in small, specific amounts called quanta (1900). Quantum is the minimum quantity of energy that can be lost or gained by an atom. The relationship between a quantum of energy and the frequency of radiation is illustrated by the following equation: E = hν E is the energy, in joules, of a quantum of radiation, ν is the frequency in s -1 of the radiation emitted, and h is a physical constant now known as Planck’s constant.

12 Einstein proposed that electromagnetic radiation has dual wave-particle nature (1905). These particles are called photons. A photon is a particle of electromagnetic radiation having zero mass and carrying a quantum of energy. In order for an electron to be ejected from a metal surface, the electron must be struck by a single photon possessing the minimum energy and frequency to knock it loose. The energy of a particular photon depends on the frequency of the radiation. E photon = hν Show video clip.

13 Hydrogen Atom Line-Emission Spectrum The ground state is the lowest energy level of an atom. When it has higher potential energy an atom is in its excited state. When an excited atom returns to its ground state it gives off energy in the form of colored light. (Example: Neon lights.)

14 When doing experiments with hydrogen gas, it was found that hydrogen atoms emit only specific frequencies of light. The fact that hydrogen atoms emit only specific frequencies of light indicated that the energy differences between the atom’s energy states were fixed. This suggested that the electron of a hydrogen atom exists only in very specific energy states (led to quantum theory).

15 Hydrogen’s Line Emission Spectrum Show video clip.

16 Bohr Model of the Hydrogen Atom Niels Bohr proposed a model of the hydrogen atom that showed that the electron can circle the nucleus only in allowed paths (orbits) (1913).


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