1. Recap – Last Lecture To balance a nuclear equation check the sum of the charge (subscript) and mass (superscript) of reactants equals those of the products. eg  1

2. 2 Structure of Atoms – History 1808 J DaltonAtomic Theory All matter consists of atoms that cannot be created or destroyed. Atoms of one element cannot be converted into atoms of another element. Atoms of an element are identical and are different from atoms of any other element. http://www.humantouchofchemistry.com/john-dalton.htm

3. 3 Structure of Atoms - History 1897 J J Thomson Cathode Rays Negatively charged particles. All metals produced the same particles. ~ 1000 times lighter than a hydrogen atom. Atoms are divisible! Cathode rays were later renamed electrons.

4. 4 Structure of Atoms – History 1909 E Rutherford Nucleus of atom Atoms are mostly empty space occupied by electrons. All the positive charge and essentially all the mass lies in the nucleus.

5. 5 Electrons in atoms can only occupy certain energy levels (orbits). When an electron moves from one energy level to another, energy is absorbed or emitted. This energy corresponds to light of a specific energy/frequency. commons.wikimedia.org/wiki/File:Niels_Bohr.jpg 1909 N Bohr Electrons in orbits Structure of Atoms – History

6. 6 Electromagnetic Radiation Silberberg fig 7.3

7. 7 Electromagnetic Radiation Wavelength,, lambda The distance between two adjacent identical points of the wave. Frequency,, nu The number of wave crests passing a given point per unit time. Blackman fig 4.1

8. 8 Wavelength and frequency are related to the speed of light. c = Electromagnetic Radiation All light waves travel at exactly the same speed (in a vacuum) – the speed of light, c, is a constant. C = 2.998  10 8 ms -1

9. 9 Electromagnetic Radiation All radiation may have the same speed but the energy can vary. The higher the frequency, the more rapidly the wave is oscillating and the higher the energy. Energy = Planck’s constant  frequency E = h h = 6.626 x 10 -34 Js

10. 10 Large wavelength Low frequency Low energy Short wavelength High frequency High energy Electromagnetic Radiation

11. 11 Converting wavelength  energy eg: A radio station transmits at a wavelength of 2.84 m. Calculate the frequency. = c so = c / = 3.00 x 10 8 ms -1 / 2.84 m = 1.056 x 10 8 s -1 or 106 MHz (this is the radio station 2JJJ) Calculate the energy associated with this radiation. E = h = 6.626 x 10 -34 J s x 1.056 x 10 8 s -1 = 7.00 x 10 -26 J and for one mole of radiation E = 7.00 x 10 -26 J x 6.022 x 10 23 mol -1 = 0.0421 J mol -1

12. 12 Atomic Emission Spectra Our eye sees the combination of wavelengths: Pass the light through a prism to see the lines: commons.wikimedia.org/wiki/File:Flametest--Na.swn.jpgcommons.wikimedia.org/wiki/File:Flametest--Cu.swn.jpg commons.wikimedia.org/wiki/File:Flametest-Co-.swn.jpg Co Cu Na http://chemistry.bd.psu.edu/jircitano/periodic4.html Co Cu Na

13. 13 Only light of certain energies is emitted. The pattern of lines is unique to hydrogen. Suggests the process of emitting light from the atom is quantised (comes in discrete amounts). Atomic Emission Spectrum

14. 14 The Bohr Model Silberberg Fig 7.11 Energy n = 1 n = 2 n = 3 n = 4 n = 5 n = 6

15. 15 Theory & Experiment agree for H Energy of the hydrogen atom orbits is inversely proportional to the square of the orbit number: E = - E R (1 /n 2 )Z 2 E R = 2.18 x 10 -18 J Z = atomic number As  E = E final - E initial then  E= - 2.18 x 10 -18 J (1/n 2 final - 1/n 2 initial ) Z 2 n=3→n=2 n=4→n=2n=5→n=2n=6→n=2

16. 16 eg: Calculate the wavelength of light emitted when an electron moves from the n = 3 to the n = 2 orbit of a hydrogen atom.  E = - 2.18 x 10 -18 J (1 / 2 2 - 1 / 3 2 ) (1) 2 =- 3.03 x 10 -19 J (minus indicates light emitted) Now E = h and E = hc / So = hc / E = (6.626 x 10 -34 Js) (3.00 x 10 8 ms -1 ) / (3.03 x 10 -19 J) = 6.56 x 10 -7 m or 656 nm (red light) n=3→n=2 n=4→n=2n=5→n=2n=6→n=2 Theory & Experiment agree for H

17. 17 Applications Na + is the major ion in extracellular fluid. The body requires 1-2 mmol/day of Na + while a typical daily diet contains 100 times this amount. The excess is excreted by the kidneys. Commonly, atomic absorption spectroscopy is used to determine sodium ion concentration in which the line in the spectra used for this measurement is at 589 nm

18. By the end of this lecture, you should: −Recognise the historical context of the Bohr model of the atom. −Be able convert between the wavelength, frequency and energy of light. −Be able to calculate the energy of a hydrogen orbit. −Be able to calculate the atomic emission spectrum of a hydrogen atom. −be able to complete the worksheet (if you haven’t already done so…). 18 Learning Outcomes:

19. 19 Questions to complete for next lecture: 1.Calculate the energy of the light absorbed when an electron in a hydrogen atom moves from the n=1 to n=3 orbit. 2.What is the wavelength of the light calculated in Q1? 3.To what region of the electromagnetic spectrum does the light in Q1 belong? 4.Is the light in Q1 more or less energetic than the light associated with the transition between n=2 to n=3? Type of LightWavelength range Ultra-violet10 – 400 nm Visible400 – 700 nm Infra-red700 nm – 1 mm