Chapter 4- Understanding the Line Spectra

Importance of the Electron It is the number of e- and the type of electrons in the outer shell of an atom that determines the chemical properties of the atom. Since chemistry is the study of the behaviour of matter, it is important that we understand the concepts surrounding the electron, especially the latest theory regarding it.

Electromagnetic Spectrum

Back to the Spectral Analysis Activity… Hydrogen’s Line Spectra

Back to the Spectral Analysis Activity…

Back to the Spectral Analysis Activity… Neon’s Line Spectra

Mercury’s Line spectrum…

So what do atoms have to do with Light? Much of the understanding of how electrons behave in atoms comes from the studies of how light interacts with matter. When an object (metal ) is heated, it first gives off infrared radiation ( heat energy), then if it is still heated, it starts to glow red, then yellow, and finally white if you get it hot enough.

This baffled scientists, they couldn’t understand what was causing the different colors at the different temperatures.

History In the 1860's, Robert Bunsen and Gustav Kirchoff determined that each element, when heated to a vapour (gas), produced its own unique bright-line spectra. (See page 135/6 of your textbook) Using this process of spectral analysis, many new elements were discovered such as: Cs Rb He (on the sun)  

They first burned each element to see the flame color produced. How does it work??? They first burned each element to see the flame color produced.

Hydrogen Emission spectral Analysis How does it work?? Hydrogen Emission spectral Analysis

Neither scientist, however, could explain what caused the bright-line spectra or why it was different for each element.

Explanation for the bright-line spectra?? It came in stages… Ernest Rutherford, through the results of his famous Alpha scattering “gold foil" experiment, theorized that the atom consisted of A dense, very small positive nucleus surrounded by e- in motion like bees around a hive. He also was not able to explain the occurrence of the bright-line spectra from his model...but it got the ball rolling!

What was the experiment? He shot alpha particles (positive charge) at pure gold ( very thin) and noticed that many of the alpha particles were deflected by something in the atoms… Only positives repel positives – so he discovered that there were positive charges present in the atom!

Rutherford built upon the idea that when + charges meet + charges … they repel… Scientists had thought the atom looked like this… Negative charges distributed evenly throughout the atom (the “plum pudding model” ) But after the gold foil expt., they found that there was something Positive was blocking some of the + particles they shot at it, and came up with the fact that there must be a positive nucleus!

Niels Bohr (1885-1962), a Danish scientist, did come up with a model to explain the spectral lines for the elements using what he called the "staircase model".

In 1913, Bohr stated that e- around an atom could only exist in definite energy levels (shells/orbits) which possessed a definite amount of energy. These energy levels (shells/orbits) are found at certain distances from the nucleus. These distances determine the energy of each particular energy level The further the energy level is from the nucleus… the more potential energy (PE) the e- possesses in that energy level (EL).

He also stated that e- can "jump" from energy level to energy level if they gain or lose just the right amount of energy.(a quantum of energy). They gain energy by Absorbing energy from high energy Electromagnetic Radiation (UV and up) High temperature Electrical energy

After an atom has been heated or electrocuted…it’s electrons absorb the perfect amount of energy…become EXCITED… and jump to a higher energy level!!! BUT…Electrons want low energy (more stable), so they quickly release the energy (as light), and fall back down to their original level What happens…

5. The “Staircase Model” (p. 136-137) This is the model that Bohr proposed in order to explain the spectral lines for the elements. Each line on the diagram represents a different energy level with a different amount of energy.

e- want to possess as little energy as possible so try to stay in the lowest possible energy level (Ground State = Stable!!!) close to p+ = happy! By gaining just the correct amount of energy, electrons will "jump“ to a higher energy level (Excited State) In order to get back to a lower energy level, which is where they want to be, the electrons release energy in the form of Light.

The different colours of light correspond to different amounts of energy that are released by the e-…according to Plank’s equation:   where E = energy   h = Planck’s constant (6.26X10-34J ∙s)   ν = frequency (Hz) (how many wavelengths in a given time) Which has a greater frequency…Blue light or red light? Since the frequency of blue light is greater than that of red light, blue light is higher in energy than red light E=h∙ν

The colour of the lines in the bright-line spectra indicate how much energy the electrons have given off as they drop to a lower energy level. Blue light in the spectra indicates that the electrons have released more energy than, say, red light. Light is the energy released (photons) by the electrons jumping back down from a higher energy level to a much lower energy level. (see pp. 137 for different series) A Photon is a quantum (a specific packet) of energy… light is made of photons!! Photons behave as both a particle and a wave = Wave/Particle duality of light!

8. Because each electron can jump to many different energy levels 8. Because each electron can jump to many different energy levels and drop back to many more, the bright-line spectrum for any element usually consists of many different coloured lines. (See pg.136 and chart) 9. Why are the bright-line spectra different for each element?   a. Each element has its own unique energy levels due to its different # of p and e!  10. Bohr was able to accurately predict, using his model, the bright-line spectra for the hydrogen atom. a. BUT could not predict the bright-line for any other element (this is clearly a problem!)

Take a look…

Another look… The drop from energy level 4  1 releases lots of energy (small wavelength) = blue light band The drop from energy level 3  1 releases moderate energy (medium wavelength) = green light band The drop from energy level 2  1 releases little energy (long wavelength) = red light

Do we Get why the bands of the line spectra are there now… It's all about the electrons in the atom dropping down from different orbits back to ground state and releasing a certain amount of energy as they drop! The energy released depends on which level you dropped from (shows up as a colour band on the spectrum) The greater the drop = they higher the energy released = the shorter the wavelength (closer to the blue light side)!

To do… Do the Chapter 4 Review and Enforcement Sheet ( don’t do first 5 questions!) Review Hebden pages 139-144