Electrons And The Bright-Line Spectrum

Submitted By emlang22
Words: 928
Pages: 4

Emma Lang
Electrons and the Bright-Line Spectrum
The purpose of this lab is to observe the bright-line spectrum of the elements Mercury and Helium and calculate the frequency and energy of each photon of light given off by the specific element.
In 1909 a scientist named Rutherford fired alpha particles which were positively charged through gold foil. He discovered 99.99% of the particles went straight through the gold foil and only .01% of the particles were deflected back. Rutherford then developed a model of the atom which he believed consisted of mostly empty space. Rutherford believed the atom contained a positively charged nucleus with electrons which orbited it like the planets around the Sun. Rutherford’s model had one flaw; the movement of the electrons seemed to be losing energy continuously and the electrons were falling into the nucleus. This shortcoming led into another scientists study. In 1913 Bohr set out to improve Rutherford’s model. In Bohr’s model the electrons stayed in orbit and did not lose energy. Bohr discovered that electrons could only orbit in a certain level. Each level had a distinct amount of energy. The farther away an electron was the more energy the electron had. Bohr’s model shows that each electron has a definite amount of energy while it’s in orbit and can travel around the nucleus without giving energy off. Bohr discovered the farther away an electron is the more energy is needed to keep it in orbit. These bands of energy which keep electrons in orbit are known as energy levels. From here, Bohr determined how many electrons could occupy every energy level inside an atom. The lowest energy level that electrons can occupy is known as the ground state. This is where the electron is the most stable. If an atom absorbs enough energy then the electron will jump to another energy level. When an electron jumps from ground state to another state it is the electrons excited state in which it is very unstable. In order to become more stable, the electron will return to a lower energy level when it releases energy. The energy released from the electron is known as a photon of light. These photons of light are visible and can be examined using the bright-line spectrum. Each element has its own bright-line spectra in which it can be identified.

Procedure: * Observe the elements Mercury and Helium through a spectroscope * Record each color and its corresponding wavelength * Calculate the frequency and energy of each wavelength
Mercury Color | Wavelength (x 10-5cm) | Frequency (1/s) | Energy (J) | Blue | 4.5 x 10-5cm | 6.7 x 1014 1/s | 4.4 x 10-19 | Yellow | 6.0 x 10-5cm | 5.0 x 1014 1/s | 3.3 x 10-19 | Green | 5.6 x 10-5cm | 5.4 x 1014 1/s | 3.6 x 10-19 |


Color | Wavelength (x 10-5cm) | Frequency (1/s) | Energy (J) | Red 1 | 7.0 x 10-5cm | 4.3 x 1014 1/s | 2.9 x 10-19 | Yellow | 6.0 x 10-5cm | 5.0 x 1014 1/s | 3.3 x 10-19 | Dark Green | 5.0 x 10-5cm | 6.0 x 1014 1/s | 4.0 x 10-19 | Dark Blue | 4.7 x 10-5cm | 6.4 x 1014 1/s | 4.2 x 10-19 | Blue | 4.5 x 10-5cm | 6.7 x 1014 1/s | 4.4 x 10-19 | Red 2 | 7.8 x 10-5cm | 3.8 x 1014 1/s | 2.5 x 10-19 | Blue-Green | 4.9 x 10-5cm | 6.1 x