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Where are the electrons ? Rutherford found the nucleus to be in the center. He determined that the atom was mostly empty space. So, how are the electrons.

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Presentation on theme: "Where are the electrons ? Rutherford found the nucleus to be in the center. He determined that the atom was mostly empty space. So, how are the electrons."— Presentation transcript:

1 Where are the electrons ? Rutherford found the nucleus to be in the center. He determined that the atom was mostly empty space. So, how are the electrons arranged in that space? This was the shortcoming, of Rutherford's model of the atom, the electron position could not be explained.

2 What is the electromagnetic spectrum? brainstorm

3 Electromagnetic radiation Is a form of energy that exhibits wavelike behavior as it travels through space. Electromagnetic spectrum – all the forms of electromagnetic radiation.

4 Electromagnetic spectrum

5 Each line on the spectrum represents a certain wave frequency of radiation. Each wave frequency is associated with a certain amount of energy.

6 Electromagnetic spectrum  The elecromagnetic spectrum includes all forms of radiation, one of which is visible light -- the radiation to which our eyes are sensitive.  Gamma-rays, X-rays, ultraviolet, infrared and radio waves are also forms of radiation. We divide the spectrum up according to the wavelength of the radiation.wavelength

7 Electromagnetic spectrum


9 A wave is described by Wavelength - the distance between corresponding points on adjacent waves Frequency- the number of waves that pass a given point in a specific time, usually one second. Unit is the hertz, (Hz). es_particles/ es_particles/

10 Wave description Mathematically – c= speed = 3 x10 8 m/s –Speed of light= frequency x wavelength –Since c is constant –Frequency is inversely proportional to wavelength –All three, speed, wavelength, frequency are related

11 example What is the frequency of a wave that has a wavelength of 200 nm? 3 x 10 8 / 200 nm =

12 1900’s Wave model of light was accepted

13 Early 1900’s Photoelectric experiments Experiment - Planck Photoelectric effect – refers to the emission of electrons from a metal when light shines on the metal. Shined light on a metal varying the frequency of the light. Below a certain frequency the electrons were not emitted. When the frequency was high enough there was enough energy to knock electrons loose from the metal.

14 Quantum Planck suggested that an object emits energy in little packets of energy called a quantum. Quantum – minimum quantity of energy that can be lost or gained by an electron. Energy that is given off is related to wave frequency.

15 Go to tumzone/lines2.html tumzone/lines2.html es_particles/

16 Planck in other words. Proposed that there is a fundamental restriction on the amount of energy that an object emits or absorbs; each piece of energy is a quantum. E= hv h = 6.6262 x 10 -34 Js v= the frequency of the wave

17 Dual wave-particle nature of light Einstein expanded on Planck’s theory and said that electromagnetic radiation has a dual wave-particle nature. While light exhibits wave like properties, it can also be thought of as a stream of particles. Each particle carries a quantum of energy – these particles are called photons.

18 Vocab Photon – particle of electromagnetic radiation having zero mass and carrying a quantum of energy. E photon = hv Quantum – the minimum quantity of energy that can be gained or lost by an electron.

19 Photoelectric effect explained Electromagnetic radiation is only absorbed in whole number of photons. In order for an electron to be bumped off, it must be struck by a photon of a certain minimal amount of energy. The minimal energy corresponds to a minimal frequency

20 continued Different metals require energy at different frequencies to exhibit a photoelectric effect. Each metal has a certain required minimal level of energy required for the electrons to be knocked loose.

21 Ground state The lowest possible energy level Close to the nucleus The lowest energy state of an atom

22 Excited state When an atom absorbs energy and the electrons move to a higher energy level. An atom has a higher potential energy than its ground state.

23 Like a ladder Energy levels are like the rungs on a ladder, you can not step between the rungs. Electrons must jump from level to level, they can not reside between the levels

24 electrons When an electron gains enough energy it jumps up an energy level ( rung on a ladder), and becomes excited. It then immediately returns to its ground state. The energy is released as a particular wavelength that corresponds to a particular color. A photon of radiation.

25 Line emission spectra The spectra given off by the electrons of a certain element. Each spectra is unique to the element or compound. Line spectra are used to identify elements and/or compounds

26 Line emission spectra The fact that specific frequencies are emitted indicates that the energy differences between an atoms energy states are fixed Each line on the spectra represents a photon of energy.

27 Bohr model of atom The electron circles the nucleus in an orbital. When the electron gains enough energy ( a certain photon) it jumps to a higher level orbital. When it returns to ground state it emits the energy (photon). The frequency of the photon is seen in the spectra.

28 Bohr successfully calculated the hydrogen line spectra Bohrs model of the hydrogen atom accounted mathematically for the energy of each of the transitions of the Lyman, Balmer and Paschen spectral series. Bottom line – Bohrs model of the H atom, with the line spectra, could be mathematically supported. summary and beyond

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