{"id":10150,"date":"2016-09-16T08:03:13","date_gmt":"2016-09-16T12:03:13","guid":{"rendered":"https:\/\/www.masterorganicchemistry.com\/?p=10150"},"modified":"2025-09-24T10:39:03","modified_gmt":"2025-09-24T15:39:03","slug":"introduction-to-uv-vis-spectroscopy","status":"publish","type":"post","link":"https:\/\/www.masterorganicchemistry.com\/2016\/09\/16\/introduction-to-uv-vis-spectroscopy\/","title":{"rendered":"Introduction To UV-Vis Spectroscopy"},"content":{"rendered":"

Understanding UV-Vis Spectroscopy Will Make You More Fun At Parties<\/strong><\/p>\n

In today’s post we’ll discuss why most molecules are colourless, introduce the useful technique of UV-visible spectroscopy, and finally explain\u00a0why<\/em> molecules like chlorophyll and \u03b2-carotene are coloured.<\/p>\n

We’ll even finish by showing you how to use a UV-Vis spectrum to predict the color of a molecule, which is a great trick to roll out at parties. No, seriously, it’s incredible. We’re talking about understanding the chemistry of colour, everyone! It doesn’t get better than this.<\/p>\n

First,\u00a0the summary. We’ll walk through the\u00a0details below.<\/p>\n

\"summary<\/p>\n

Table of Contents<\/strong><\/p>\n

    \n
  1. Converting Frequency Units Into Energy\u00a0 Units<\/a><\/li>\n
  2. Ground\u00a0 State Electrons Can Be Promoted To Excited\u00a0 States Through The Absorption\u00a0 Of Light<\/a><\/li>\n
  3. Case Study: The Molecular Orbital\u00a0 Diagram Of H2<\/a><\/li>\n
  4. Why Most Molecules Containing Only Single Bonds Are\u00a0 Colorless<\/a><\/li>\n
  5. Pi Bonds Absorb At Longer, More Energetically Accessible Wavelengths<\/a><\/li>\n
  6. The UV-Vis Spectrometer<\/a><\/li>\n
  7. How Does\u00a0 Conjugation Of Pi Bonds Affect Lambda Max?<\/a><\/li>\n
  8. How Does Lambda Max Relate To The Color We Perceive?<\/a><\/li>\n
  9. Conclusion: UV-Vis Spectroscopy<\/a><\/li>\n
  10. Notes<\/a><\/li>\n
  11. Quiz Yourself!<\/a><\/li>\n<\/ol>\n
    \n

    <\/a>1. Converting The Frequency Of Light Into Energy Units<\/strong><\/h2>\n

    If you want someone to read something you’ve wrote, starting with an equation is generally a bad idea. However, we’re going to start with one of the most beautiful, amazing, and just plain\u00a0useful\u00a0<\/em>equations in all of science. So if you drop off after reading this, really, that’s your problem.<\/p>\n

    From general chemistry, you may recall the immortal equation<\/p>\n

    E = h\u03bd <\/em><\/p>\n

    where E<\/strong> is energy, h<\/em><\/strong> is Planck’s constant (6.626 \u00d7 10-34<\/sup>\u00a0m2<\/sup>\u00a0kg \/ s\u00a0) and\u00a0\u03bd<\/strong>\u00a0<\/em>is frequency \u00a0(in m-1<\/sup>) .<\/p>\n

    Why is this equation so frickin’ useful? Because it\u00a0relates energy to the frequency of light.\u00a0<\/strong><\/p>\n

    To be more specific, when I say “light” I mean, “photon”, as in a carrier of electromagnetic radiation. For the purposes of today’s post, here’s the part of the electromagnetic spectrum<\/a> we’ll be discussing today: the UV and visible frequencies.<\/p>\n

    \"electromagnetic<\/p>\n

    <\/a>2. Ground State Electrons Can Be Promoted To Excited States Through The Absorption of Light<\/h2>\n

    Way back in general chemistry (Bohr model of the hydrogen atom<\/a>, anyone?) you saw how an electron can be promoted from the ground state<\/strong> orbital to an excited-state<\/strong> orbital through the absorption of a photon of frequency:<\/p>\n

    \u03bd\u00a0=\u00a0\u0394E \/\u00a0h <\/em><\/p>\n

    \u00a0<\/em>[where\u00a0\u0394E is the difference in energy between the ground and excited states].<\/p>\n

    Since it is an electron that is being promoted from one energy level to another, we call these “electronic transitions<\/strong>“. Generally, the frequency of radiation required for electronic transitions is in the\u00a0ultraviolet and visible portion\u00a0<\/strong>of the electromagnetic spectrum.<\/p>\n

    This has practical importance in a great number of ways, but for our purposes, we’ll see that the most prominent is that light can promote electrons from\u00a0bonding<\/strong> orbitals to\u00a0anti bonding\u00a0<\/strong>orbitals, and therefore potentially lead to the breaking of chemical bonds.<\/p>\n

    This\u00a0will help\u00a0us to understand:<\/p>\n