Untitled Flashcards Set

In Part A of this experiment, the visible emission spectrum of hydrogen will be examined. A spectroscope (Fig-

ure 3.1) will be calibrated using the known wavelengths emitted by helium, and then the same spectroscope will

be used to determine atomic hydrogen emission line positions. The principal quantum numbers and their cor-

responding energies will be determined and assigned to the hydrogen emission lines. The value of the Rydberg

constant will also be determined experimentally.

In Part B, flame tests will be made for several metal ions in the form of chloride salts. An unknown will be

identified, demonstrating the use of this simple test as an analytical tool.

Figure 3.1: Spectroscope.

Theory

Electrons in an atom can exist only in certain allowed energy levels. Electrons in a ground-state atom can be

excited to a high-energy state by absorption of energy, for example, from a high-voltage discharge. When the

electron returns to a lower-energy state, radiation in the form of light may be emitted. The energy emitted is

quantized, and corresponds to discrete frequencies (which are proportional to energies) and corresponding wave-

lengths in the spectrum. The emission spectrum then consists of discrete lines rather than a continuum in which

light is emitted at all frequencies. Many such electronic transitions take place with emission of light in the vis-

ible region of wavelength 370–700 nm; these transitions can be observed with relatively simple equipment and

the human eye. It is also possible to examine changes in energy states of atoms and molecules which do not

correspond to radiation in the visible region, but more elaborate instrumentation is required.

Laboratory Manual Prepared by Catalyst Education, LLC for the University of California at Riverside

Department of Chemistry.

3.2

Emission of Light from Hydrogen and Metal Atoms

Figure 3.2: Transitions in the hydrogen atom (shown schematically–not to scale.

We will examine hydrogen in detail in this experiment because it has a simple emission spectrum which has been

well characterized. Figure 3.2 shows transitions of a single electron between different principal quantum levels

(note that this is a schematic representation and is not to scale; the energy difference between each set of “orbits”

is not equal). The wavelengths of light in the electromagnetic spectrum corresponding to changes in electronic

states in hydrogen were described empirically by Rydberg according to the formula:

𝜆 = 𝑅𝐻( 1

𝑛2

𝑙−

𝑛2

ℎ

) (Equation 3.1)

where 𝜆 is the wavelength of light of the spectral line, 𝑛𝑙 is the lower (more stable) principal quantum number,

𝑛ℎ is the higher principal quantum number, and 𝑅𝐻 is the Rydberg constant with a value 𝑅𝐻 = 1.097 × 107𝑚−1

(Note: The Rydberg constant 𝑅𝐻 is not the same as the gas constant, which is given the symbol 𝑅.) Equation 3.1

may be described in terms of energy by using the equations 𝐸 = ℎ𝜈 and 𝑐 = 𝜆𝜈, where ℎ is Planck’s constant

(6.626 × 10−34𝐽 ⋅ 𝑠) and 𝑐 is the speed of light (3.00 × 108 m/s).

Transitions to 𝑛𝑙 = 1 correspond to a return to the most stable energy level (ground state). Because 𝑛𝑙 = 1 char-

acterizes the state closest to the positively charged nucleus, the energies associated with removal or return of an

electron involving this level are the highest. In fact, emission of energy for a transition to 𝑛𝑙 = 1 is found in the

Laboratory Manual Prepared by Catalyst Education, LLC for the University of California at Riverside

Department of Chemistry.

3.3

Emission of Light from Hydrogen and Metal Atoms

ultraviolet region for all values 𝑛ℎ. The name given this series of transitions is the Lyman series.

Similar series corresponding to transitions from higher energy levels to 𝑛𝑙 = 2, 3, 4, ... are also possible and are

found in Table 3.1. Note that as 𝑛𝑙 becomes larger the energy associated with the transition becomes smaller and

the wavelength becomes larger. We will be particularly interested in the transition to 𝑛𝑙 = 2, the Balmer series,

because the resulting emissions are in the visible region and can be easily studied with a Bunsen spectroscope.

Table 3.1: Hydrogen Spectral Series

Series 𝑛𝑙 Region of Spectrum

Lyman Balmer Paschen 1 ultraviolet

2 visible

3 infrared

Spectroscope

When light emitted from a discharge tube is passed through a prism, it is refracted into components of different

colors. Each of the lines in the spectrum thus produced is light of a particular wavelength. The spectroscope

(Figure 3.1) is an instrument containing a diffraction grating, a scale, and an eyepiece for viewing the spectrum.

It is difficult to manufacture a spectroscope containing a scale such that when light of 400 nm is viewed through

the eyepiece, it falls on a line on the scale marked 400 nm. The usual approach is to put approximate numbers

on the scale. Then the person who uses the instrument “calibrates” the scale. This is done by using a high-

voltage discharge tube containing a gas (for this experiment, helium) that produces lines whose wavelengths

are accurately known. The scale readings at which these lines appear correspond to the known wavelengths.

A “calibration curve” is then constructed by graphing the scale readings versus the known wavelength values.

Along the x-axis are the known wavelengths; along the y-axis are the scale readings. These points are connected

with a smooth curve.

Laboratory Manual Prepared by Catalyst Education, LLC for the University of California at Riverside

Department of Chemistry.

3.4

Emission of Light from Hydrogen and Metal Atoms

Flame Tests

A high-temperature flame (such as one from a Bunsen burner) can also be used to promote electrons in metal

atoms to excited states. When the electrons return to the ground-state configuration, energy in the form of light

of a wavelength (or color) characteristic of the element is emitted. This light appears to color the flame, thus

the term “flame test.” For ex- ample, when sodium atoms are heated in a flame, the outer 3𝑠 electrons may be

excited to the higher-energy 3𝑝 level. When the excited electron returns to the 3𝑠 level, the transition results in

the emission of yellow-orange light with a wavelength of about 590 nm, which is also the familiar color of sodium

vapor lamps.

Laboratory Manual Prepared by Catalyst Education, LLC for the University of California at Riverside

Department of Chemistry.