1. Chapter 4: Arrangement of Electrons in Atoms Coach Kelsoe Chemistry Pages 97-122

2. Section 4-1: The Development of a New Atomic Model Coach Kelsoe Chemistry Pages 97 – 103

3. The New Atomic Model  The previous models we studied did not explain where electrons were located.  A new atomic model revealed a relationship between light and an atom’s electrons.  We used to think of light as a wave. Light acts like a wave, but has particle- like characteristics as well.

4. The Wave Description of Light  Visible light is a type of electromagnetic radiation, a form of energy that exhibits wavelike behavior as it travels through space.  Gamma rays, x-rays, ultraviolet, visible light, infrared, microwaves, and radio waves make up the electromagnetic spectrum.

6. The Wave Description of Light  Wave motion involves wavelength and frequency.  Wavelength is the distance between corresponding points on adjacent waves.  Frequency is the number of waves that pass a given point in a specific time, usually one second. Frequency is measured in hertz. HIGH FREQUENCY LOW FREQUENCY

7. Frequency and Wavelength  Frequency and wavelength are related mathematically: c= λv, where c= speed of light, λ= wavelength, and v= frequency of the electromagnetic wave  Since c is a constant, the product of λv is constant too.

8. Photoelectric Effect  The photoelectric effect refers to the emission of electrons from a metal when light shines on the metal.  For a given metal, no electrons were emitted if the light’s frequency was below a certain minimum, no matter how long the light was shone. If light was a wave, it would be able to knock loose an electron from the metal. We couldn’t explain why there must be a minimum frequency in order for this effect to take place.

9. Planck’s Explanation  German physicist Max Planck suggested that a hot object emits energy in small, specific amounts called quanta (plural for quantum).  A quantum is the minimum quantity of energy that can be lost or gained by an atom.

10. The Particle Description of Light  Planck proposed the following relationship between a quantum of energy and the frequency of radiation.  Where E = energy (in Joules), v = frequency of the radiation emitted, and h is Planck’s constant, which is always 6.626 x 10 -34 J·s.

11. The Particle Description of Light  Einstein expanded this theory by saying that electromagnetic radiation has a dual wave-particle nature.  Light exhibits many wavelike features, but also acts like a stream of particles.  Each particle, or “photon,” of light carries a quantum of energy.  A photon is a particle of electromagnetic radiation having 0 mass and carrying a quantum of energy.

12. Einstein’s Photons  According to this equation, the minimum energy for an electron to be ejected from a metal surface depends on the minimum frequency.

13. The Hydrogen-Atom Line-Emission Spectrum  When current is passed through a gas at low pressure, the potential energy of some of the gas atoms increases.  The lowest energy state of an atom is its ground state.  A state in which an atom has a higher potential energy than it has in its ground state is an excited state.  When an excited atom returns to ground state, it gives off light.

14. What a Bright Idea!  Excited neon atoms produce light when they fall back to a lower energy excited state or the ground state.

15. Colors of Neon Signs  When electric current passes through noble gases, they emit different colors. HeliumPink NeonOrange-Red ArgonLavender KryptonWhite XenonBlue

16. Line-Emission Spectrum  When a narrow beam of the emitted light was shined through a prism, it was separated into a series of specific frequencies (and therefore specific wavelengths, λ = c/v) of visible light.  The line-emission spectrum of an element is a series of wavelengths of emitted light created when the visible portion of light from excited atoms is shined through a prism.

18. Line Spectra for Hydrogen, Mercury, and Neon H Hg Ne

19. Hydrogen’s Line-Emission Spectrum  Scientists expected to observe a continuous spectrum when they observed hydrogen through the spectrum, but they only saw a few select colors.  A continuous spectrum is the emission of a continuous range of frequencies of electromagnetic radiation. An example of this is a rainbow.

20. The Perplexing Electron of Hydrogen  Scientists were confused at why hydrogen only emits specific frequencies when current passes through it.  This suggested that the electron of a hydrogen atom exists only in very specific energy states.  If an atom at excited state (E 2 ) falls back to ground state (E 1 ), it releases a photon that has energy: E 2 -E 1 =E photon = hv

21. Bohr to the Rescue!  Danish physicist Niels Bohr proposed a model of the hydrogen atom that linked the atom’s electron with photon emission.

22. Atoms Aren’t Bohr-ing!  Bohr’s model of the atom says: The electron can circle the nucleus only in allowed paths, or orbits. When the electron is in one of these orbits, the atom has a definite, fixed energy. The electron, and the atom as well, is in its lowest energy state when the electron is in the orbit closest to the nucleus. The orbit is separated from the nucleus by a large empty space where the electron can not exist.

23. Here’s a Real-World Example  Let’s say I was climbing a ladder… On the first rung of that ladder, I don’t have much potential energy. It won’t hurt if I fall off. On the 20 th rung of the ladder, I do have a lot of potential energy. Welcome to broken- neck city!

24. And That’s Not Where It Ends!  And just as I can not stand where there is not a rung, an electron can not occupy where there is not an orbit.

25. How Does This Explain the Lines?  While an electron is in an orbit, it can neither lose or gain energy, but it can move to a higher energy orbit by gaining an amount of energy equal to the difference in energy between the higher- energy orbit and the initial lower-energy orbit.  When hydrogen is in an excited state, its electron is in a higher-energy orbit.

26. How Does This Explain the Lines?  When the atom falls back from the excited state, the electron drops down to a lower-energy orbit.  In the process, a photon is emitted that has an energy equal to the energy difference between E 2 -E 1.  Photons are absorbed when the electron moves to a higher energy orbit.  Photons are emitted when electrons move to a lower-energy orbit.