Emission Spectra and Quantum Leaps

Introduction Rutherford’s Gold Foil Scattering experiments published in 1911, established that the atom consisted of a small, dense positively charged nucleus surrounded by an electron cloud. The Rutherford model offered no details as to the behavior of the electrons and could not explain why the negatively charged electrons were not pulled into the positive nucleus. Scientists trying to design experiments to determine how electrons behaved inside the atom faced a daunting challenge: how do you “see” an electron that is so tiny that it fits inside an atom and is moving near the speed of light? The key insight was made by a Danish physicist Niels Bohr who realized the emission of light from elements that had absorbed energy in the form of electric current, light or heat could be explained by quantum leaps of electrons. In this investigation, we will look at patterns of light released from different elements and understand how quantum mechanics can explain the observations.

Purpose: To observe the spectra of white light and the line emission spectra of several elements including hydrogen and to understand how your observations can be explained by quantum mechanics. Safety: No concerns although a giddy euphoria can be produced because the experiments are so cool! Materials: Special spectral glasses fitted with prism lens that separate light into its individual wavelengths. High voltage power source fitted with spectral tubes filled with gas samples of different elements.

Part 1: Spectrum of White light Procedure: Put on your spectral glasses, look at a source of white light and sketch a diagram of your observations in a box (see below).

White Light:

What colors do you see? Do the colors appear as distinct lines or do they appear to blend? (In reality, although the colors appear to blend, they actual consists of thousands of individual colors line so close together that your brain interprets the colors as blended) Part 2: Emission Spectrum of Element Procedure: Put on your spectral glasses, look at the light source for each element and sketch a diagram of your observations in a box (see below). What colors do you see? Do the colors appear as distinct lines or do they appear to blend? Mark lines in the appropriate color area

Helium H ydrogen Violet blue green yellow orange Violet blue green yellow orange red red

QNeonuestions: Violet blue green yellow orange 1) How are the patterns of light emitted from the element samples differ from the pattern of red white light?

2) Are the line spectra patterns of the different elements the same or different? What does it mean to describe the line spectra patterns of an element as “fingerprints”?

3) How do astronomers know that our sun is composed mainly of hydrogen and helium if no one has ever been to the sun?

4) The fact that only particular individual wavelengths and frequencies of light are emitted is taken as evidence that the light energies are quantized. What does the term quantized mean?

5) Provide an example of quantized system and an example of a non-quantized system.

6) Bohr proposed the Bohr Model in 1913 which introduced the concept that electron energies are quantized. A key feature of this model is the concept of the quantum leap to explain the emission line spectra. A) Importance of energy absorbance: No light emission is observed until the atom absorbs energy in the form of heat, light or electricity. According to the Bohr model, what part of the atom absorbs the light energy?

B) Which represents a lower energy, more stable state? (Does nature prefer oppositely charged particles close together or far apart?)

Choose the appropriate word to complete each statement. (From POGIL, edited by Laura Trout, Flinn Scientific). C) Electrons and protons (attract/repel) each other.

D) As an electron get closer to the nucleus the (attraction/repulsion) to the nucleus gets (stronger/weaker).

E) For an electron to move from an energy level close to the nucleus to an energy level far from the nucleus it would need to (gain/lose) energy.

F) For an electron to move from an energy level far from the nucleus to an energy level close to the nucleus it would need to (gain/lose) energy.

7).Which picture shows a quantum leap in which an electron is absorbing energy? Which picture shows a quantum leap which releases or emit light energy? Label each picture below as either ABSORPTION OR EMISSION.

8) Use the energy level diagrams below for questions. The y-axis represents energy.

8A) Consider the energy diagram for the hydrogen atom on the left. Is the energy distance between all of the energy levels exactly the same? In other words, is the amount of energy released when electron jumps from level 4 down to level 3, the same amount of energy that would be released when an electron jumps from level 2 down to level 1?

8B) Are the energy gaps the same in both elements? In other words, would an electron jumping from level 2 to level 1 in hydrogen, release the same amount of energy as an electron jumping from level 2 to level 1 in a helium atom?

8C) The energy that an electron has in an energy level is determined by the attraction of the electron from the positive protons and the repulsion the electrons experience from other electrons. How would these interactions be different in a hydrogen atom vs a helium atom? (Hint: How many protons and electrons in each type of atom?)

9) In the hydrogen spectra, the deep blue/violet line at 434 nm, the teal line at 486 nm and the red line at 656 nm corresponds to quantum leaps from a higher energy excited state of either 3,4 or 5 down to level 2. Which color line is produced by the 3 → 2 transition? Which color line is produced by the 4 → 2 transition? Which color line is produced by the 5 → 2 transition? EXPLAIN.

n = 5

n = 4 n = 3

n = 2

10) Each of the lines of color of the hydrogen emission is the result of a different quantum leap. A hydrogen atom has one electron. How can all of these different leaps be happening at once? How can all of these different colors of light be emitted at the same time?

11) Why can we see only 3 lines of color emitted from the hydrogen spectra? Each quantum leap from an excited state to a ground state should release energy in the form of light (e.g. 2 → 1, 3 → 1, 4 → 1, etc.), yet we see only 3 colors of light from the 5 → 2, 4 → 2, 3 → 2 transitions. Why don’t we see more colors of light?