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Electromagnetic Spectrum and Bohr’s Model

Learn about the properties of light waves, line spectrum, Bohr’s model of the hydrogen atom, energy quantization, and how to draw Bohr diagrams for elements. Also, explore qualitative analysis methods including flame tests and Lewis symbols.

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Electromagnetic Spectrum and Bohr’s Model

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  1. UNIT 1 – MATTER AND QUALITATIVE ANALYSIS The Electromagnetic Spectrum, Bohr’s Model, Chemical Bonding

  2. OBJECTIVES

  3. The electromagnetic spectrum • Electromagnetic Energy: light energy that travels in the form of waves • Frequency: the number of cycles per second • Wavelength: he distance between successive crests or troughs in a wave • Nanometer: 10-9 m; unit nm

  4. WAVES

  5. LIGHT AND THE ELECTROMAGNETIC SPECTRUM • There are many different kinds of light: X-rays, ultraviolet, infrared, microwaves are all examples of light • Visible spectrum: the region of the electromagnetic spectrum tat the human eye can see • 400 to 700 nm • A rainbow contains all the colours of visible and near visible light • Continuous Spectrum: an uninterrupted pattern of colours that is observed when a narrow beam of white light passes through a prism

  6. PRISM

  7. VISIBLE SPECTRUM

  8. LINE SPECTRUM • Different types of matter emit different wavelengths of light: • Neon gas emits red light • Sodium ions emit yellow light • Hydrogen gas emits blue/violet light • Each element has a different ‘line spectrum’ • Discontinuous spectrum of light that is produced when light the light from the gas is passed through a type of prism called a spectroscope • You see coloured lines rather than a rainbow

  9. Continuous vs. Line

  10. DIFFERENT ELEMENT, DIFFERENT LINE SPECTRUM: THE ELEMENT ‘FINGERPRINT’

  11. QUESTIONS • Describe the difference between radio waves and X-rays using the concepts of wavelength, frequency, and energy • What range of wavelengths of electromagnetic radiation can the human eye detect • White light is composed of many different colours of light. Explain. • Distinguish between a continuous spectrum and a line spectrum

  12. The bohr-model of the hydrogen atom • Negatively charged particles orbiting positive nucleus according to Rutherford… what is the problem with that? • (+) and (–) attract… atom should collapse on itself • Also, laws of moving charges means that orbiting electrons should emit energy and eventually run out.... • SPIRAL DEATH!!! (whomp whomp… that’s a problem) • MATTER IS VERY STABLE

  13. HYDROGEN GAS LINE SPECTRUM

  14. Nielsbohr • To the rescue!!! • Looking at the discrete line spectrum of hydrogen… observed discrete lines, inferred fixed energy levels of electrons • Inferred electrons are restricted to quantized energy levels • Quantized: possessing specific values or amounts (a quantity) • That is, electron’s can only have specific energy and orbit around the nucleus in orbitals at specific distances, the further away the greater the energy

  15. Explanation • Electrons occupy specific energy states (orbitals) • Can only transfer levels by absorbing the EXACT amount of energy needed between them, to go up OR to give off (emit) the energy difference between lower energy levels.. • Hence, discrete energy levels • In doing so, it either absorbs or emits a particular amount of light…

  16. each atom has slightly different allowable energy differences and different numbers of electrons that can do the transitions, therefore different spectrums for each • Ground state: level the electron wants to be at • Excited state: level with added energy

  17. Each energy level can only have a certain number of electrons in it • 1st = 2 • 2nd = 8 • 3rd = 8

  18. example • Draw the Bohr diagrams for: • Neon • Helium • Carbon • Calcium • Oxygen

  19. THE NUMBER OF ELECTRONS IN THE VALENCE (OUTER) ORBITAL (SHELL) CORRESPONDS TO THE COLUMN (GROUP) THE ELEMENT IS N ON THE PERIODIC TABLE • Ex: Oxygen (Group 6) has 6 electrons in its valence shell; Sodium (group 1) has 1

  20. Homework questions • State the reasons why Rutherford’s model of the atom failed to describe the observed behaviour of matter • Describe Bohr’s model of the atom. How is it similar and how is it different to Rutherford’s Model of the Atom? • What role did Spectroscopy play in helping Bohr come up with his model? • Why do electrons emit light energy when they drop from a higher to lower energy level? • What does it mean by “energy levels are quantized? • Drawing Bohr diagrams

  21. FLAME TESTS https://www.youtube.com/watch?v=d8hpUtRnsYc

  22. Four types of qualitative analysis • Thermal Emission Spectroscopy (TES): substances are identified based on amount of heat they emit • Line spectroscope: substances are identified based on their line spectra • Flame tests: substances are identified based on the colours they emit when placed in a flame • Carbonation: the presence of metals is identified based on the emission of carbon dioxide when reacted with a metal

  23. Lewis symbols • The inner electrons are really just place holders; outer electrons determine reactivity • Chemists shorten Bohr Diagrams to Lewis Structures • Example: Draw the Lewis Structure For • Ca H • N Ne • Li Cl

  24. Polyatomic ions • Ion that is composed of two or more atoms (all of them end in –ate, you will be given a list of these on a test) • Examples: • Sulfate SO42- • Bromate BrO3- • Chlorate ClO3- • Nitrate NO3- • Phosphate PO43-

  25. Forming ionic compounds • Occur when metal and non-metal react with each other • Ionic crystals: solid that consist of a large number of cations and anions arranged in repeating 3D patterns • Ex: Show the formation of Sodium Chloride; Calcium Oxide

  26. Naming ionic compounds • USE IUPAC (International Union of Pure and Applied Chemistry) • METAL + NON-METAL (IDE) • Example: • CaO • NaCl • Al2O3

  27. Writing formulas • Use the criss-cross method • Example: • Sodium Oxide • Rubidium Fluoride • Strontium Nitride

  28. practice • Worksheets!!

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