Understanding Light Interactions: A Lab Exploration of Spectra and Emission Lines, Summaries of Chemistry

Download Understanding Light Interactions: A Lab Exploration of Spectra and Emission Lines and more Summaries Chemistry in PDF only on Docsity!

lOMoARcPSD|

S. Tauhid,

Chemical principles

Sadaf Tauhid

Chem 117: Section 2 Work Station 5, TA Daniela Graf Stillfried

Lab Partners: Janet Bangura, Josseline Aupla

Lab 3: Spectra

Introduction:

Electromagnetic radiation exists as a form of energy, that also exists as a photon. Photons

are energy particles that are released during atomic processes, and through these particles, they

act as waves that are wavelengths and frequencies. Light is put into categories based on its

wavelength and frequencies. Each wavelength then has its own nm (nanometer) that can be used

to classify it into visible and UV light. Frequency and wavelength are inverses because when

frequency decreases then wavelength increases, and vice versa. Since electromagnetic radiation

has different frequencies and wavelengths, they have different energies as well, that can be

calculated using Planck’s constant.

When gas-phase ions and atoms absorb light, it’s possible for the outermost valence

electrons to be upgraded to a higher energy level. The emissions from an excited gas-phase atom

can be seen through a spectroscope. This is used to diffract incoming light to the different

wavelengths, which then distributes photons of the different energies across the two-dimensional

plane. It’s also used to measure the energies of the emission lines.

Data:

Table 1 shows the different flame ionization test colors that were seen during the first part

of the experiment, by naming the color seen for each of the different metals.

Table 1. Flame Ionization Colors

Metal Barium Calcium Cesium Lithium Potassium Sodium

Color Green Orange Violet Red Purple Orange

Unknown Flame Color Green Unknown Metal Barium

Table 2.1 shows the wavelength center of the spectrum emitted by the incandescent light,

given in nm.

Table 2.1. Composition of Fluorescent Lamp

Red Light Green Light Blue Light

650 nm 520 nm 470 nm

Table 2.2. shows the wavelength center of the spectrum emitted by the mercury vapor

tube, given in nm.

Table 2.2. Composition of Mercury Tube

Spectroscope Calibration Line for Helium

60

50

40

30

20

10

0 350 400 450 500 550 600 650 700 750

Spectroscop Reading (r), nm

f(x) = 0 x² − 1.15 x + 357.

Table 4 shows the emission spectrum that was shown by sodium. It displays the

wavelength reading, alongside the estimated spectroscope error and the corrected wavelength.

All of these are given in nm.

Table 4. Emission Spectrum of Sodium

Emission Line of Sodium—Wavelength 578 nm

Estimated Spectroscope Error 10 nm

Corrected Wavelength 588 nm

Table 5 shows the transition of the sodium emission spectrum, alongside the wavelengths,

frequencies, photon energies, and the spectral region as well.

Table 5. Transitions and Calculation of Sodium Emission Spectrum

Transition Wavelength

(nm)

Spectral

Region

Frequency

(s

• 1 )

Photon

Energy

(J/photon)

Photon

Energy

(kJ/mol)

A. [Na]4p

1

[Na]3s

1

330 UV 9.09 * 10

14 6.02 * 10

• 19 362

B. [Na]4s

1

[Na]3p

1

1140 x-ray 2.63 * 10

14 1.74 * 10

• 19

C. [Na]3d

1

[Na]3p

1

819 UV 3.66 * 10

14 2.43 * 10

• 19

D. [Na]3p

1

[Na]3s

1

588 visible 5.10 * 10

14 3.38 * 10

• 19

Table 6 shows the transitions of the sodium energy in an energy level diagram.

Spectroscope

Error

(e), nm

Table 6. Sodium Energy Levels

[Na]4p

1 362 kJ/mol (^) [Na]4p^1 362 ↓_

[Na]3d

1 349.86 ↓

[Na]4s

1 308.3_↓C ↓

[Na]3p

1 203.5↓B_↓ ↓

[Na]3s

1 ↓D 0 A↓

[Na]3d

1 146.36 kJ/mol

[Na]4s

1 104.80 kJ/mol

[Na]3p

1 203.5 kJ/mol

[Na]3s

1 (ground state) zero

Discussion Questions:

1. In this lab, there was an attempt to understand how light interacts with matter, and its

properties. First, the four examples of the emission light spectrum were provided. There

was an observation of hydrogen, helium, nitrogen and mercury’s spectral lines in order to

determine and measure their wavelength. This first experiment showed the relationship

between the wavelength, energy released and the spectral lines. The shorter that the

wavelength was, the greater the energy released by the element’s atom was based on the

formula provided. It was found that the vapor that was being emitted was most likely a

mixture of all three—the incandescent lamp, the helium and the mercury tubes—because

of the wavelength that was being shown. The next part of the experiment was based on

the spectroscopic error of helium, and it was shown that the reading of the helium tube

turned out to be less in terms of nm than the known amount—meaning there was an error

of calculation. The final experiment proved that sodium isn’t able to absorb proper light

because it dropped a transition state each time, until it hit zero completely.

2. The spectroscopic error in Part 3 was most likely random due to the precision limitation

of the spectroscope. It’s not possible for the reading of each color to be accurate because

it is difficult to read the numbers, which leads to an error each time. There was a

spectroscope error bound to happen with each color due to the fact that it wasn’t possible

to accurately read the wavelength of each color exactly. It’s random because the data only

relied on the precision of the reading through the spectroscope.

3. The wavelength of the visible emission line in the spectrum for sodium is 588.9950 nm,

and the reading was 578 nm. The experimental error turned out to be 1.86%. The

experimental error calculation is shown below.

%error =

theoretical value − experimental value ∗ 100 theoretical value

%error =

%error =1.86 %