Atomic and Nuclear Physics

1. Atomic and Nuclear Physics The Atom Light Spectrum for Hydrogen Modern Model The Photon

2. The Atom • Matter is made of particles called atoms. • The atom was made of a positive and negative side like a magnet. This model continued until JJ Thompson discovered electrons in 1897. This let to the plum pudding model.

3. Rutherford Model: the atom consists of a small dense positive nucleus, surrounded by electrons that orbit the nucleus (as planets orbit the sun) as result of electrostatic attraction between the electrons and the nucleus. • In 1909 under the direction of Ernest Rutherford, Geiger and Marsden performed an experiment which proved that Thompson’s model was not correct either.

4. Marsden and Geiger fired very small fast moving particles called alpha particles at a very thin gold foil. Some of the particles went through, and others bounced back. The ones that bounced back implied that the alpha particles had a head on collision with a heavier particle; whatever this was had to be very small since few hit it head on. This small object must also be positive causing the positive alpha particle to be pushed away.

5. Rutherford Experiment

6. Bohr’s Model • The model created by Rutherford caught on faster although somewhat inaccurate than the Bohr model that followed. • In 1913, Neils Bohr developed a model that had the electrons orbiting the nucleus of the atom like planets orbiting the Sun. • In order for a body to rotate around another body there must be a force of attraction. For the atom this is an electrostatic attraction between the electrons and the nucleus.

7. Limitations to Rutherford’s Model • Did not explain why electrons surrounding the nucleus were not drawn into the nucleus by strong electrostatic attractions to the protons of the nucleus. • Did not specify composition of nucleus. • How did protons (positively charged particles) in the nucleus stay closely bound when electrostatic forces should have forced them apart?

8. Energy Levels • All elements will emit light in characteristic colors when heated. • The patterns of light from heated gases differ from those of sunlight, consisting of a series of bright lights separated by dark gaps—line spectrum. • Emission spectrum: a spectrum of light emitted by an element; a series of bright lines, with dark gaps between the lines where no light is emitted. • Absorption spectrum: a bright continuous spectrum covering the full range of visible colors, with dark lines where the element absorbs light. • Spectroscopy: the study of emission and absorption spectra.

9. When an electron falls between two energy levels it will emit a photon equal in energy to the difference in energy levels. The energy of a photon is dependent on its frequency. Thus, the existence of discrete wavelengths in the spectrum is evidence that energy levels are discrete.

10. Relationship between the atom and light • Bohr’s model did explain the connection between the atom and light. According to this model, the electrons could only exist in certain orbitals. The radius of the orbit was dependent on the energy of the electrons; high energy electrons had larger orbits. • When the electrons changed from a high to a low energy orbit, energy was emitted in the form of light.

11. Frequency of light • The frequency of the emitted light was found to be proportional to the change of energy of the electron. hf= E2 – E1 Where h= Planck’s constant = 6.6 x 10-34 m2 kg s-1 Since blue light has the highest frequency, it must come from the greatest energy changes.

12. Line Spectrum for hydrogen

13. The Modern Model Computer generated model of a Neon atom with electron energy levels. www.scienceclarified.com/As-Bi/Atom.html

14. The photon • Light undergoes reflection, refraction, diffraction, and interference. • If light is produced each time an electron changes from high energy to low energy then light is not a continuum but made of units or packets. These packets are called photons. • Each photon has a frequency that is proportional to the change in energy of the electron.

15. Law of Conservation of Energy • It states that the total amount of energy in an isolated system remains constant over time (is said to be conserved). • According to the law of conservation of energy, this means that each photon has an energy hf. ΔE=hf • A beam of light is made up of a large number of photons emitted each time an electron changes energy levels. The brighter the light, the photons in the beam.

16. The energy of an electron • Using the formula ΔE=hf the energy change required to produce a photon of red light of wavelength 700nm can be calculated by: f =c/λ = 3.0 x 108 /700 x 10-9 = 429 x1012 Hz ΔE=hf = (6.6 x 10-34 )(429 x1012) = 2.8 x 10-19 J

17. Electron voltage • We found that the energy of the electron was 2.8 x 10-19J so to find voltage we divide Joules by 1.6 x10-19 J and you get eV. 2.8 x 10-19J/ 1.6 x10-19 J = 1.75eV

18. Practice Problems Refer to the figure at left to help with questions. 1. How many possible energy transitions are there in this atom?