Spectrum and its types-Jayam chemistry learners

 What is spectrum? Explain its types.

Since childhood, we might have experienced the beauty of a rainbow in the sky at least once. When we turn the pages of high school science books, we might find somewhere that sunlight passing through raindrops portrays this colorful arrangement of visible light in the sky known as a rainbow. Until we read the atomic spectrum, we might not realize it as a continuous solar spectrum. Let us dive into today's blog topic- "Spectrum and its types."

What is spectrum?

A spectrum is a patterned arrangement of various light wavelengths obtained by passing polychromatic radiation through the prism.

For example: When we pass white light or sunlight through the prism, it splits into seven colored radiations of visible light is nothing but the visible spectrum or solar spectrum.

Similarly, every chemical element of the universe exhibits a unique atomic spectrum. It helps to determine their deposits in a specific terrestrial or astronomical object. Hence, spectral studies play a vital role in the composition analysis of the universal matter.

Table of contents:

1. What is spectrum?

2. What is light?

2.1. Monochromatic light

2.2. Polychromatic light

3.Discovery of the spectrum

4.Types of the spectrum

4.1.Emission spectrum

4.2.Absorption spectrum

4.3.Continuous spectrum

5.Types of the spectrum of light

5.1.Solar emission spectrum

5.2. Solar absorption spectrum

6. Atomic spectrum

6.1. Numerical problems on atomic spectrum

6.2. Atomic emission spectrum

6.3. Atomic absorption spectrum

6.4. Difference between emission and absorption spectrum

7.Atomic spectra of hydrogen and sodium atoms

7.1.Hydrogen atomic spectrum

7.2.Sodium atomic spectrum

8.Uses of spectral studies

9.What is spectroscopy?

10.MCQ on the spectrum and its types

What is light? What is the wavelength classification of light radiation?

The modern quantum theory explains that light is a flow of photons. It is a propagating wave of electric and magnetic fields couple. So, wave characteristics such as wavelength and frequency express the photon energies of light radiations.

Light radiation is of two types when classified according to its wavelength. They are

1. Monochromatic radiation

2. Polychromatic radiation

Monochromatic light:

• It consists of light radiations of a single wavelength.
• For example, Lyman-alpha radiation occurs at 121.5 nm.
• Monochromatic radiations cannot exhibit a spectrum.
• Instead, each spectral line of atomic spectra denotes a single radiation wavelength, i.e., monochromatic radiation.

Polychromatic light:

• It consists of light radiation of multiple wavelengths.
• For example, white light is a mixture of seven different colored light radiations with altering wavelengths.
• Polychromatic light can exhibit a spectrum.
• Polychromatic light that passes through the prism splits into its constituent wavelengths. And it is a dispersion process. When the dispersed light proceeds through the spectrometer exhibit its spectrum on a photographic plate fixed over there.
• Alternatively, a spectrum is an arrangement of dispersed radiations of polychromatic light.

Discovery of the spectrum:

In ancient times, people identified the composition of ores by their imparted colors to the flame.

In 1666, Isaac Newton predicted that white light is a mixture of seven colored radiations of sun rays. He proved it by passing white light through the prism. He obtained a continuous rainbow of sunlight which was named spectrum. So, Sir Isaac Newton originated spectral studies by inventing the word spectrum.

In 1802, William Hyde Wollaston discovered the first crude spectrometer based on Newton’s model.

In 1859, Gustav Kirchhoff and Bunsen developed a refined spectrometer to understand the relationship of chemical elements with their spectral traits. They explained the chemical composition of compounds from spectral data.

Types of the spectrum:

It is of three kinds. They are;

1. Emission spectrum
2. Absorption spectrum
3. Continuous spectrum

Emission spectrum:

It is a specimen of bright lines or bands obtained by analyzing the emitted photons of an excited atom. The expelled light radiations of excited atoms pass through the spectrometer to produce their emission spectra. The emission of light radiations by the atoms in an excited state gives its emission spectrum.

The emission spectrum of an atom is of two kinds.

1. Continuous emission spectrum
2. Discontinuous emission spectrum

Continuous emission spectrum:

A continuous emission spectrum is a pattern of diffused glowing bands merged into each other. Here, one color band consolidates into the other without any separation. There is no beginning or ending to the color bands of the continuous emission spectrum. For example- The solar spectrum of sunlight

The continuous emission spectrum corresponds to the undivided diffused bright bands.

Discontinuous emission spectrum

It consists of a series of sharp lines separated by dark bands. Here, the spectral lines are separate and unmerged. And each spectral line corresponds to a particular wavelength or frequency. For example -The atomic spectra of chemical elements.

Its other name is the line spectrum which corresponds to the sharp lines of the spectrum.

Absorption spectrum:

It is a pattern of dark lines or bands corresponding to the light radiations absorbed by the substance. The transmitting light beam acquired from the absorbing medium produces the absorption spectrum when passed through the spectrometer. Absorbed light radiations of ground-state atoms generate their absorption spectrum.

Continuous spectrum:

It consists of a definite range of all wavelengths of the electromagnetic spectrum. The source of continuous spectra is hot celestial matter and stars. The range of wavelengths that occur in the continuous spectrum depends upon the temperature of the emitted light source.

Types of the spectrum of light:

Sunlight exhibits both absorption and emission spectrums.

Solar emission spectrum:

Sunlight that passes through the prism shows a series of seven-colored diffused bands called the solar spectrum. It is an example of a continuous emission spectrum. The seven colors of the solar spectrum are Violet, Indigo, Blue, Green, Yellow, Orange, and Red, called VIBGYOR.

Solar absorption spectrum:

The temperature of the outer layers of the sun's atmosphere is lower than that in its center. Hence, the absorption of light radiations of different wavelengths in the outer solar atmosphere gives a continuum of dark absorption lines called Fraunhofer lines. The German physicist Joseph von Fraunhofer studied the dark absorption lines of the solar spectrum. Therefore, the solar absorption spectrum is also known as the Fraunhofer spectrum.

Atomic spectrum:

1. What is an atomic spectrum?

Answer:

It is the structured arrangement of electromagnetic radiation sorted by increasing wavelengths, derived from the analysis of light emitted or absorbed by an atom of an element.

This spectrum comprises a series of spectral lines, with each line denoting the energy of electromagnetic radiation in terms of wavelengths or frequencies

2. What are the different categories of atomic spectra?

Answer:

There are two classifications of atomic spectra:

•  Atomic emission spectrum
•  Atomic absorption spectrum

The atomic emission spectrum features series of spectral lines representing the electromagnetic radiation emitted by the substance. Conversely, the atomic absorption spectrum displays a dispersed pattern of electromagnetic radiation absorbed by the substance.

3. What energy level must electromagnetic radiation possess to generate the atomic spectra of an atom?

Answer:

Neil Bohr proposed that atoms possess infinite discrete stationary orbits, each associated with a specific amount of energy, termed energy levels.

These energy levels vary depending on the atom's structure and electronic configuration. Bound electrons within the atom absorb electromagnetic radiation of specific frequencies, unique to each electron transition and atom, varying with temperature and pressure.

Atomic spectroscopy utilizes these principles to identify known and unknown elements within compounds.

4. Why are spectral line series observed instead of a single line in the atomic spectrum of an element?

Answer:

We understand that electron transitions between distinct atomic energy levels are quantized. Each spectral line in the atomic spectrum represents a specific electron transition. When electromagnetic radiation passes through a gaseous sample, different atoms absorb radiation of different frequencies corresponding to their electron transitions. Consequently, we observe spectral line series indicating multiple electron excitations.

5. Is it possible to observe a rainbow when sunlight passes through an atom?

Answer:

No, a rainbow isn't observed when sunlight is passed through an atom under normal conditions. To generate an atomic emission spectrum, the atom's electrons need to be excited either by intense heating in a flame or by passing an electric discharge at low pressure.

The bound electrons in the atom's ground state cannot absorb energy continuously to produce a rainbow-like spectrum. Instead, they absorb discrete energy specific to electron transitions, resulting in bright, sharp lines against a dark background, demonstrating the quantized nature of light.

6. Why is the term "line spectrum" used interchangeably with "atomic spectrum"?

Answer:

The atomic spectrum of an atom is characterized by discrete, sharply defined spectral lines instead of continuous, uninterrupted spectral bands.

Each line in the atomic spectrum corresponds to a specific wavelength or frequency of photon associated with a particular electron transition. This distinct feature of the spectrum, with its clearly separated lines, is why it is often referred to as a line spectrum.

7. What makes the atomic spectrum akin to a fingerprint for atoms?

Answer:

Just as each individual possesses a unique fingerprint, each atom exhibits a distinct atomic spectrum determined by its structure and electron arrangement. This characteristic spectrum serves as a signature for identifying elements.

Across the periodic table's 118 known elements, each possesses its own specific atomic spectrum.

For instance, sodium's emission spectrum prominently features a peak at 589 nm, while potassium's peak is observed at 766 nm.

Consequently, atomic spectroscopy plays a crucial role in various fields, including celestial matter analysis, metallurgy, biological research, and medicinal studies.

8. What are the defining features of an atomic spectrum?

Answer:

1. Each atom has its characteristic emission and absorption spectra. 
2. The wavelengths of bright lines in the emission spectrum are the same as that of the dark lines in the absorption spectrum. 
3. The spectroscopic data of known substances serves as a reference to reveal spectroscopic data of unknown bodies. Hence it helps in the identification of new chemical elements in the universe.

9. In what way does the atomic spectrum provide proof of the quantization of light?

Answer:

The particle nature of light, demonstrated through quantization, implies that if an atom's ground state electron absorbed light across all frequencies continuously, the resulting atomic spectrum would be continuous, not discontinuous.

This discontinuity in the spectrum highlights that the absorption or emission of photons by electrons occurs in quantized steps, aligning with the energy difference between specific stationary orbits involved in electron transitions.

Consequently, the atomic spectrum displays distinct, sharp lines, providing evidence of discrete energy transitions within the atom.

10. What method do you use to derive the energy of light from the wavelength of a spectral line in the atomic spectrum?

Answer:

Based on the corpuscular nature of light demonstrated by the atomic spectrum, we can employ Planck's quantum theory to ascertain the energies of photons engaged in different electron transitions. According to this theory,

ΔE = nhν

Where, ΔE = Energy difference between initial and final states

n = number of photons

h = Planck’s constant

ν = Frequency of light

To determine the energy of light based on the wavelength of a spectral line in the atomic spectrum, you can use the equation:

energy = (Planck's constant * speed of light) / wavelength.

This equation relates the energy of a photon to its wavelength through Planck's constant (h) and the speed of light.

By plugging in the wavelength of the spectral line, you can calculate the corresponding energy of the light.

To grasp the concept more effectively, here are a few numerical exercises.

Numerical problems on Atomic spectrum

1. Neon gas is used in the sign boards. If it emits strongly at 616 nm, calculate (a) frequency of emission (b) distance travelled by this radiation in 30 s (c) energy of quan tum and (d) number of quanta present if it produces 2J of energy

Answer:

To calculate frequency of emission, we have;

ν =  ^c⁄_λ

 = ^{3 x 108}⁄_{616 x 10^-9}

 = 4.87 x 10¹⁴ s^-1

Distance travelled by the radiation in one second = 3 x 10⁸ m

Distance travelled by radiation in 30 seconds = 30 x 3 x 10⁸ = 9 x 10⁹ m

Energy of one quantum of radiation = hν

  = 6.62 x 10^-34 x 4.87 x 10¹⁴ = 32.27 x 10^-20J

Total energy produced = 2J

Energy of one quantum of radiation = 32.27 x 10^-20 J

Number of quanta = ^{Total energy produced} ⁄ _{Energy of one quantum}

 = ² ⁄ _{32.27 x 10^-20}

 =6.2 x 10¹⁸

2. In astronomical observations, signals observed from the distant stars are generally weak. If the photon detector receives a total of 3.15 x 10^-18 J from the radiations of 600 nm, calculate the number of photons received by the detector

Answer:

Total energy received by photon detector = 3.15 x 10^-18 J

Wavelength of radiation 600 nm = 600 x 10^-9 m

Energy of one photon = hν = ^hc ⁄ _λ

 = ^{6.626 x 10-34 x 3 x 108} ⁄ _{600 x 10^-9}

 =3.313 x 10^-19 J

Number of photons received by the detector = ^{Total energy received}  ⁄ _{Energy of one photon}

 = ^{3.15 x 10-18} ⁄ _{3.313 x 10^-19}

 = 10

3. Lifetimes of the molecules in the excied states are often measured by using pulsed radiation sources of duration nearly in the nano second range. If the radiation source has the duration of 2 ns and the number of photons emitted during the pulse source is 2.5 x 10¹⁵, calculate the energy of the source

Answer:

Duration of radiation source = 2ns = 2 x 10^-9 s

Frequency = ¹ ⁄ _{2 x 10^-9}

  = 0.5 x 10⁹ s^-1

Energy of the source = Nhν = (2.5 x 10¹⁵) x 6.626 x 10³⁴ Js (0.5 x 10⁹ s^-1) = 8.25 x 10^-10 J

4. The longest wavelength doublet absorption transition is observed at 589 nm and 589.6 nm. Calculate the frequency of each transition and the energy difference between the two excited states.

Answer:

Given that λ₂ = 589nm = 589 x 10^-9 m

λ₁ = 589.6 nm = 589.6 x 10^-9 m

ν₂ = ^c ⁄ _λ2

ν₂ = ^{3 x 108} ⁄ _{589 x 10^-9}

ν₂ = 5.093 x 10¹⁴ s^-1

ν₁ = ^c ⁄ _λ1

ν₁ = ^{3 x 108} ⁄ _{589.6 x 10^-9}

ν₂ = 5.088 x 10¹⁴ s^-1

5. Similar to electron diffraction, neutron diffraction microscope is also used for the determination of the structure of molecule. If the wavelength used here is 800 pm, calculate the characteristic velocity associated with the neutron

Answer;

Mass of neutron = 1.675 x 10^-27 kg

Wavelength = 800 pm = 800 x 10^-12 m

According to de-Broglie equation, λ = ^h ⁄ _mv

800 x 10^-12 = ^{6.626 x 10-34} ⁄ _{1.675 x 10^-27 x ν}

ν = ^{6.626 x 10-34} ⁄ _{1.675 x 10^-27 x 800 x 10^-12}

 = 494 ms^-1

Atomic emission spectrum

It is a discontinuous spectrum of emitted light radiations given off by an atom's excited electron. It consists of a series of sharp lines separated by dark bands.

Each atom has a specific electron arrangement that keeps its energy low and stable. However, the atomic emission spectrum is produced by analyzing the radiations emitted by excited atoms.

The atoms can be excited in the following ways:

• (a) By heating the substance intensely 
• (b) By passing an electric discharge through a gas at low pressure 
•  (c) By passing an electric current through a thin metal filament

When any of the above methods are applied, the electron in the lowest energy state of an atom absorbs energy, causing it to jump from the ground state to a higher energy state.

The ground state electron cannot absorb just any amount of energy; the energy of the absorbed photon must be equal to or greater than the energy difference between the two energy states in the atom.

∆E = E₂-E₁

Where, ∆E = The amount of energy emitted or absorbed

E₂ = Energy of higher stationary level of an atom

E₁ = Energy of lower stationary level of an atom

It is evident that the potential energy of an excited electron is greater than that of its ground state, making it unstable. Therefore, the excited electron returns to its initial lower energy state by emitting the energy of the absorbed photon.

The released photon's energy must equal the energy gap between the two stationary shells involved in the electron transition, following the law of conservation of energy.

When the radiations emitted by excited gaseous substances are analyzed using a spectrograph, a discontinuous spectrum of sharp bright lines separated by dark bands is observed. This photographic recording of emitted light radiations is known as an emission spectrum.

Each bright line corresponds to the wavelength of a photon emitted during a specific electron transition, while the dark background indicate the photon energies absorbed during that electron movement.

Emitted photons are colored when their wavelengths lie in the visible region, but those beyond the visible spectrum are not perceptible to the human eye without special spectroscopic equipment.

You might wonder why atoms in their normal state do not exhibit an atomic emission spectrum. The reason is that electrons can only undergo excitation after absorbing energy. Once in the excited state, the unstable electron emits light as it returns to the lower, stable shell. This emitted light is analyzed by the spectroscope to generate the emission spectrum.

To obtain an atomic emission spectrum, we supply either heat or electric energy. Additionally, according to Bohr’s atomic theory, an atom in its ground state does not emit energy. Therefore, atoms at normal temperatures do not exhibit an emission spectrum but rather an absorption spectrum.

Typically, noble gases appear colorless due to their complete octet configuration. However, when an electric discharge is passed through them at low pressure, each noble gas element emits a distinct color.

Helium emits a pale yellow glow, neon emits a reddish-orange hue, argon emits blue light, krypton emits a whitish-purple glow, xenon emits blue or lavender light, and radon emits a yellow-green color in the discharge tube. Therefore, noble gases are commonly used in fluorescent lighting and discharge lamps.

Table-1: A concise table detailing the visible region atomic emission lines of noble gases.
Name of the noble gas element	The color displayed by the emission line.
Helium	Pale yellow
Neon	Reddish-orange
Argon	Blue
Krypton	Whitish-purple
Xenon	Blue or lavender light
Radon	Yellow-green

In addition to noble gases, other elements exhibit specific colors when subjected to similar conditions. Nitrogen emits a pink glow, oxygen emits violet to lavender light, sodium vapor emits bright orange to yellow hues, and mercury vapor emits light blue.

Both hydrogen and water vapor emit a lavender glow at low electric current and shift to pink or magenta at currents over 10 mA. Consequently, in the signage industry, neon signs utilize long luminous gas-discharge tubes containing rarefied neon or other gases for illumination.

Each species possesses a varying number of electrons capable of various excitations under diverse temperature and pressure conditions. Consequently, each gas emits a unique glow aiding in its identification. Thus, the atomic spectrum serves as the fingerprint of atoms.

In addition to electric current, flame is also utilized in analyzing atomic emissions. Flame emission spectroscopy involves heating a sample in a flame and observing the emitted light to analyze its composition.

Alternatively, a flame test involves heating a sample with metal ions in a flame, causing the excited metal ions to emit light of distinct colors.

However, it's noteworthy that not all metal ions emit light in the visible spectrum when subjected to flame, but some of the metals that do are listed below:

Table-2: Colours emitted by metals in the flame test
Name of the metal ion	Colour imparted to the flame
Lithium	Red
Sodium	Golden yellow
Potassium	Light pink
Rubidium	Red-violet
Caesium	Violet/blue
Magnesium	Bright white
Calcium	Orange-red
Strontium	Red

The diverse array of flame colors indicates the unique electronic configurations within the atoms' orbits, which give rise to the distinct hues observed. This understanding is utilized in the creation of fireworks, where specific metal compounds are employed to generate radiant colors upon detonation.

The emission spectrum of a molecule is termed a band spectrum or molecular spectrum. Typically exhibited by compounds or gases such as N₂ and O₂ at low pressure and temperature, the band spectrum comprises numerous bright bands interspersed with dark regions.

Each band is sharply defined at one end, gradually fading towards the other—a characteristic appearance often likened to flutes, hence its alternate name, the fluten spectrum.

Line emission spectra are highly valuable in the investigation of electronic structure, as each element possesses a distinct line emission spectrum. It's impossible for any two elements to have identical line spectra. Comparing the lines of an unknown sample's emission spectrum with those of a known element swiftly confirms the identity of the former.

Spectroscopic analysis revealed the existence of elements like rubidium (Rb), caesium (Cs), thallium (Tl), indium (In), gallium (Ga), and scandium (Sc) when examining their minerals. Helium (He) was discovered in the sun through spectroscopic methods.

Atomic absorption spectrum

It's the photographic representation of the wavelengths absorbed by an atom in its ground state, characterized by dark lines interrupting an otherwise continuous spectrum.

When electromagnetic radiation passes through the vapors of an element or a solution of a substance, and the transmitted light is analyzed using a spectrograph, the resulting spectrum is the absorption spectrum of that element. The dark regions in the absorption spectrum correspond to the wavelengths of light absorbed by the element.

When light traverses through an organic compound sample, certain wavelengths are absorbed, while others remain unaltered. The wavelengths absorbed by a molecule are a result of changes in its electronic, vibrational, or rotational energy levels, permissible for its constituent atoms.

A molecule absorbs light of specific frequencies if there's an energy transition within it, indicated by ΔE, equal to hν. The absorbed energy may either heat up the sample or be re-emitted.

A spectrometer is employed to gauge the wavelengths of absorbed photons. When changes in absorption are plotted against wavelengths, distinct absorption bands emerge, which are indicative of a compound's unique characteristics. This technique serves as an invaluable tool in determining the molecular structure of unfamiliar compounds.

The absorption of photons at specific wavelengths reveals the presence of certain functional groups within the compound.

For instance, visible and ultraviolet radiation, spanning the wavelength range of 200-800 mμ in the electromagnetic spectrum, leads to the excitation of π-electrons in conjugated or unconjugated diene systems.

In conjugated systems, the energy difference between the ground state and the excited energy level is lower, resulting in absorption at longer wavelengths.

Moreover, the carbonyl group present in an aldehyde or a ketone absorbs light at specific wavelengths within the visible and UV regions of the electromagnetic spectrum. Hence, an ultraviolet or visible spectrum is quite useful for detecting conjugation and carbonyl groups, but it may not offer information about the rest of the molecule.

Table-3: Summary of spectroscopic techniques;
Radiation absorbed	Effect on the molecule of a substance
Ultraviolet (190-400 nm)	Changes in electronic energy levels within the molecule
Visible (400-800 nm)	Conjugated unsaturation, conjugation with non-bonding electrons, extent of pi-electron system
Infra-red (667-400 cm-1)	Changes in the vibrational and rotational movements of the molecule. Detection of all functional groups which have specific vibrational frequencies such as C=O, O-H,-NH2,C=C- etc.
Radio frequency (60-300 MHz)	Nuclear magnetic resonance, induces changes in the magnetic proportion of certain atomic nuclei, notably that of hydrogen (hydrogen atoms in different environments can be detected, counted and analysed for structure determination)
Electron beam impact 70eV, 6000 kJ/mol	Ionisation and fragmentation of the molecule into spectrum of fragment ions (determination of molecular weight and deduction of molecular structure from the fragments obtained)

Excited state and the ground state:

In a photochemical process, the absorption of a quantum of energy elevates an electron from its ground state to a higher energy level. Because the energy levels of a molecule are quantized, the energy needed to raise an electron from one level to another is a fixed amount.

Therefore, only light with a frequency matching this energy difference can cause excitation. Consequently, light of any other frequency will not be absorbed and will pass through the sample without any loss of intensity.

If a compound absorbs light in the visible region, it appears colored and displays a color complementary to the absorbed wavelength. For instance, a compound that absorbs blue light will appear orange.

Table-4: Relationship of absorbed and observed colours
Wavelength absorbed expressed in A0	Absorbed colour	Observed complimentary colour
4000-4500	Violet	Yellow
4500-5000	Blue	Orange
5000-5500	Green	Red
5500-6000	Yellow	Violet
6000-6500	Orange	Blue
6500-7000	Red	Green

Formation of absorption bands;

We expect the atomic absorption spectrum to display sharp peaks, with each peak corresponding to the promotion of an electron from one energy level to another.

However, instead of sharp peaks, broad absorption bands are typically observed. This broadening occurs because the excitation of electrons is also accompanied by constant vibrational and rotational motions of the molecules, which are quantized. Thus, a molecule in a particular electronic state also exists in various vibrational and rotational states.

The difference between two adjacent electronic levels is greater compared to adjacent rotational levels, with the difference between adjacent vibrational levels being intermediate.

Electronic excitation is superimposed on rotational and vibrational levels. Thus, during promotion, an electron transitions from a specific vibrational and rotational level within one electronic state to another vibrational and rotational level within the next electronic state.

This results in numerous possible transitions, close together, involving changes in electronic, rotational, and vibrational levels. Consequently, a large number of closely spaced wavelengths are absorbed, forming bands rather than single peaks.

In complex molecules with many atoms, the multitude of vibrational sublevels and their proximity cause the discrete bands to merge, resulting in broad bands.

Difference between emission and absorption spectrum

Table-5: Difference between Emission and absorption spectrum
Emission spectrum	Absorption spectrum
It is produced by analysing the radiant energy emitted by the excited substance	It is produced when white light is passed through a substance and transmitted light is analysed by a spectrograph
It consists of some bright lines separated by dark spaces	It consists of dark lines in otherwise continuous spectrum

Atomic Spectra of Hydrogen and Sodium Atoms

Hydrogen atomic spectrum

(a) Hydrogen emission spectrum:

Consider a sample of hydrogen gas in the glass discharge tube. The electric current is passed through the hydrogen gas present in the discharge tube under low pressure.

When the hydrogen atoms absorb energy from the electric discharge, they get excited to higher energy states. And the unsettled electron in the excited state then returns to its initial position with the emission of photons of suitable wavelengths.

Then the hydrogen gas in the discharge tube glows red due to the electron transition between the energy levels.

And the emitted light radiation is passed through the slit and made to fall on the glass prism that separates the light radiation into constituent wavelengths. Finally, the photographic plate over there records the line emission spectrum of hydrogen.

The hydrogen atom, being the simplest atom with only one electron, exhibits a well-defined emission spectrum known as the hydrogen spectrum. This spectrum is composed of several series of lines, each corresponding to electron transitions between different energy levels.

The line spectrum of hydrogen displays a series of spectral lines across the ultraviolet, visible, and infrared regions. The emissions in the visible region are detectable by the human eye. Consequently, the entire hydrogen spectrum spans a frequency range from 10¹⁶ Hz to 10¹³ Hz. It is divided into five spectral series, namely:

1. Lyman series 
2. Balmer series 
3. Paschen series
4.  Bracket series
5.  Pfund series

In the Lyman series, spectral lines result from electron transitions from the first energy level (n=1) to higher states. Since these transitions require highly energetic electromagnetic radiation, this series lies in the UV region.

The Balmer series occurs due to electron transitions from the second energy level (n=2) to higher states. These transitions require less energy than those in the Lyman series, hence they appear in the visible region and produce colored spectral lines. The electron transitions to the 3rd, 4th, 5th, and 6th energy levels result in red, green, blue, and violet spectral lines, with wavelengths of 656.5 nm, 486.3 nm, 432.4 nm, and 410.3 nm, respectively.

Paschen Series: Transitions from higher energy levels (n > 3) to the third energy level (n = 3), occurring in the infrared region.

Brackett and Pfund Series: Transitions to the fourth and fifth energy levels (n = 4 and n = 5) respectively, also in the infrared region.

The wavenumber of a spectral line of hydrogen atom may be calculated using the Rydberg's equation. According to this equation:

⊽ = ¹⁄_λ = R [¹⁄_n1² - ¹⁄_n2²]

Where, R = Rydberg constant and is equal to 109678 cm^-1

In the above, n₁ and n₂ represent the principal quantum numbers of an atom's energy levels, both being integers.n₂ indicates the excited state of the atom, while n₁ represents the ground state. Consequently, n₂ is always greater than n₁. Additionally, for a specific spectral series, n₁ remains constant while n₂ varies for each line within the series.

To better understand the above formulas, let's look at an example that demonstrates how to calculate the wavelength of a photon emitted during an electron transition.

Problem-6: Calculate the wavelengths of the first line and the series limit for the Lyman series of Hydrogen.

Answer:

For first line in Lyman series n₁=1 and n₂=2

⊽ = ¹⁄_λ = 109678 cm^-1 [¹⁄_n1² - ¹⁄_n2²]

 =109678 cm^-1 [¹⁄₁ - ¹⁄₄]

 =109678 x 0.75 cm^-1 = 82259 cm^-1

λ = 1.2157 x 10^-5 cm = 1215.7 A⁰

For series limit in Lyman series n₁=1 and n₂=∞

⊽ = ¹⁄_λ =  109678 cm^-1 [¹⁄₁ - ¹⁄_∞²]

 =109678 cm^-1

λ = 9.1176 x 10^-6 cm = 911.76 A⁰

Careful observation of the hydrogen spectrum reveals that within each spectral series, the gap between spectral lines narrows at higher frequencies. This phenomenon is likely due to the smaller energy differences between the involved stationary levels.

It is known that the energy levels of a hydrogen atom are not evenly spaced; the energy difference between successive levels decreases as the value of n increases. The following formula is used to determine the energy of the different stationary orbits of hydrogen:

E_n =  - [¹³¹²⁄_n²] kJ mol^-1

Where, E_n denotes the energy of the stationary orbit n

n is an integer

Table-6:Calculates the energy gap between two different stationary levels of hydrogen atom
n value	En value in kJ mol-1	Energy difference between two successive energy states in kJ mol-1
1	-1312	-1312
2	-328	984
3	-145.78	182
4	-82	64
5	-52.48	30
6	-36.44	16
7	-26.78	10

The table above illustrates that as the energy gap between successive energy states decreases, the spacing between them also diminishes accordingly. Consequently, spectral lines appear closer together when electronic transitions occur between energy orbits with n values greater than or equal to three.

This is because only a specific range of frequencies can induce significant electronic transitions, resulting in a band-like appearance. For example, electromagnetic radiation in the IR region can cause hydrogen electron transitions from the third, fourth, and fifth main energy levels to higher levels. Therefore, the Paschen, Brackett, and Pfund series are considered compact due to the closeness of their spectral lines.

Conversely, the Lyman series is the longest series because of the significant energy gap between the first main energy level and higher energy states, resulting in spectral lines that are far apart.

From the table, we can conclude that when the energy gap between two different energy states is small, lower frequency photons are sufficient to induce electron excitations between them. Consequently, as we move from higher to lower wavelengths, the lines will be farther apart on the higher wavelength side and gradually closer on the lower wavelength side.

The point at which the series ends is called the series limit. The electronic transition involving n = ∞ as the higher energy state to the ground state represents the series limit of every series in the hydrogen emission spectrum.

Hydrogen absorption spectrum:

When white light passes through the vapors or solution of hydrogen gas, the transmitted light then passes through a spectroscope, resulting in a spectrum of dark lines at specific wavelengths. These dark lines correspond to the wavelengths of light absorbed by the hydrogen atoms. Thus, in the hydrogen absorption spectrum, we observe dark lines against a continuous white background.

If you carefully observe the hydrogen absorption spectrum, you will notice all five spectral series, as seen in the hydrogen emission spectrum, appear as dark absorption lines at the same wavelength ranges. This occurs because the hydrogen electron absorbs the same energy photons as those emitted during specific electron transitions due to the constant energy difference between the two states.

Since the energy of stationary orbits is fixed and electron transitions are quantized, the electron must absorb or emit radiation of the same wavelength to transition between specific energy states. Therefore, the wavelengths of the dark lines in the absorption spectrum of a substance are identical to the wavelengths of the bright lines in the emission spectrum of the same substance.

Sodium atomic spectrum

Sodium emission spectrum:

Vaporizing a sodium crystal in a Bunsen flame emits bright light. Passing the emitted sodium light through the prism and allowing it to fall on the photographic plate generates a sodium emission spectrum. It shows two yellow spectral lines at 5890 A⁰ and 5896 A⁰ separated by a dark space.

For an electron in an n_s orbital (l=0) , j can only be ¹⁄₂. When the electron is promoted to an n_p orbital, j may be ³⁄₂ or ¹⁄₂, and the energies corresponding to the different j values are not quite equal.

In the emission spectrum of sodium, for example, transitions from the 3p_3/2 and 3p_1/2 levels to the 3s_1/2 level therefore correspond to slightly different amounts of energy, and this spin–orbit coupling is the origin of the doublet structure of the strong yellow line in the spectrum of atomic sodium.

D-Lines:

The most notable lines in the sodium spectrum are the D-lines, consisting of two close lines at wavelengths 589.0 nm and 589.6 nm, found in the visible region. These lines are due to transitions of electrons between the 3p and 3s orbitals.

Sodium absorption spectrum:

We can observe sodium absorption through either of the following methods:

(a) Pass white light through the vapors or a solution of sodium chloride and analyze the transmitted light with a spectrograph. (b) Pass sunlight through a prism after it has passed through a sodium chloride solution.

In both cases, the sodium absorption spectrum shows two dark lines at 5890 Å and 5896 Å, known as D₁ and D₂ lines. These dark lines indicate the absorption of yellow light at these wavelengths by the sodium atom. As a result, the wavelengths of the dark lines in the sodium chloride absorption spectrum are the same as those of the two yellow lines at the same positions in the sodium emission spectrum.

Uses of spectral studies:

1. It is helpful in radar, aeronautical, and marine voice communication systems.
2. We use it in weather forecasting, radio, television, and mobile/phone communication broadcasting.
3. Besides, spectral studies have much significance in the composition analysis of astronomical matter.
4. The study of the emission and absorption spectra of the substance provides information on its structure. In particular, astronomers use this spectroscopic data to determine the constituents of stars and interstellar matter.
5. With the help of spectral line data, astronomers can determine the temperature and density of the chemical substance in the stars.
6. And it even tells about the magnetic field strength of stars.
7. The width of spectral lines informs the velocity of celestial bodies in space.
8. And the presence of water traces or winds on the stars can be confirmed by investigating their spectral information.
9. The spectral lines shifting back and forth in the experimental observations reveal the orbiting trends of the astronomical object determining its mass and size.
10. The brightness or fade of the spectral lines informs the physical changes accompanying their strength.
11. The study of light radiation's path in the intergalactic medium intimates us about the stuff like gas or dust that fills the space between the stars and galaxies. It tells the materials around the stars, galaxies, and other interstellar objects.
12. It is vital in biomedical spectroscopy, such as tissue analysis and medical imaging.

What is spectroscopy in simple terms?

Spectroscopy is the branch of science that deals with the electromagnetic spectrum as a function of its wavelength or frequency. It enables us to investigate the correlation between matter and electromagnetic radiation.

In laboratories, light radiation is allowed to pass through the sample to be analyzed. The transmitted light radiation then goes through the diffraction grating equipment and is afterward captured by a photodiode. It generates an electromagnetic spectrum of the selected sample.

It is a fundamental tool in studying the electronic structure and composition of matter at atomic, molecular, and astronomical distances. Generally, spectrometers, spectrophotometers, spectrographs, or spectral analyzers are necessary for spectral measurements.

Isaac Newton studied visible light and was the pioneer of spectroscopy. Then it was further extended by James Clerk Maxwell’s prediction on the electromagnetic nature of light. It paved the way for the evolution of the electromagnetic spectrum.

The theory of spectroscopy is each chemical element of the periodic table has a unique atomic spectrum corresponding to the photon energies emitted or absorbed by it. Hence, the line spectrum is a fingerprint to identify the atoms in unknown chemical substances.

Besides, spectroscopy is a tool to study the atomic properties of universal matter. It plays a vital role in astronomy and remote sensing on earth.

The spectroscopy techniques of chemical analysis involve the study of line spectra of atoms. Instead, biomedical spectroscopy focuses on diffraction and scattering of light to study biological tissue.

The most common branches of spectroscopy used in chemistry are; atomic spectroscopy, molecular spectroscopy, vibrational spectroscopy, rotational spectroscopy, visible spectroscopy, ultraviolet spectroscopy, infrared spectroscopy, Raman spectroscopy, and nuclear magnetic resonance (NMR) spectroscopy.

1. What is the difference between spectrum and spectra?

Answer:

The terms "spectrum" and "spectra" are related but differ in number. Here's a brief explanation:

Spectrum: This is the singular form of the word and refers to a single range of wavelengths of light or electromagnetic radiation. For example, when discussing the range of colors seen in a rainbow, you would refer to "the spectrum of visible light."

Spectra: This is the plural form of the word and refers to multiple ranges of wavelengths of light or electromagnetic radiation. For instance, when comparing the light emissions from different elements, you would refer to "the spectra of hydrogen and sodium."

In summary, "spectrum" refers to one range, while "spectra" refer to multiple ranges.

2. What are spectrum and spectroscopy?

Answer:

A spectrum is the range of different colors produced when light is dispersed by a prism or diffraction grating, showing the different wavelengths or frequencies of light. It can also refer to a range of other types of electromagnetic radiation.

Spectroscopy is the scientific study of how light interacts with matter. It involves analyzing the spectrum of light emitted, absorbed, or scattered by materials to understand their properties, composition, and structure.

3. What is the difference between a spectrum and a continuum?

Answer:

A spectrum refers to the range of different wavelengths of electromagnetic radiation, typically displayed as distinct lines or bands that correspond to specific energy transitions in atoms or molecules. Each line or band represents a specific wavelength or frequency.

A continuum, on the other hand, is a continuous range of wavelengths or frequencies without distinct lines or gaps. It represents a smooth, uninterrupted spread of electromagnetic radiation across a wide range of wavelengths, often seen in the radiation emitted by solid objects or very hot gases.

In summary, a spectrum shows discrete lines or bands, while a continuum is a seamless distribution of wavelengths.

4. What are band spectrum and line spectrum?

Answer:

Band Spectrum:

A band spectrum consists of groups of closely spaced spectral lines, forming bands. Each band has a sharp edge on one side and fades out on the other. This type of spectrum is typically observed in molecular spectra where the transitions involve both electronic and vibrational energy levels. Band spectra are characteristic of gases at low pressure and are seen in compounds like nitrogen (N2) and oxygen (O2).

Line Spectrum:

A line spectrum, also known as an atomic spectrum, consists of distinct, isolated lines, each corresponding to a specific wavelength or frequency of light emitted or absorbed by an atom. These lines represent the quantized energy levels of electrons in atoms. Line spectra are characteristic of individual elements, with each element producing a unique set of spectral lines. Examples include the hydrogen emission spectrum and the sodium emission spectrum.

In summary, a band spectrum features closely spaced lines forming bands, typical of molecular transitions, while a line spectrum consists of isolated lines, characteristic of atomic transitions.

5. What are the types of spectrum of line?

Answer:

The types of line spectra are generally classified into two main categories:

Emission Spectrum:

Continuous Emission Spectrum: Produced when a solid, liquid, or densely packed gas is heated to a high temperature, resulting in a continuous range of wavelengths.

Line Emission Spectrum: Produced by excited atoms or molecules emitting light at specific wavelengths, resulting in discrete lines. For example, the hydrogen emission spectrum.

Absorption Spectrum:

Continuous Absorption Spectrum: Occurs when a continuous spectrum of light passes through a cooler, less dense gas, absorbing specific wavelengths. This results in dark lines (absorption lines) on the otherwise continuous spectrum.

Line Absorption Spectrum: Similar to the continuous absorption spectrum, but here the absorption lines are specific to the elements in the gas phase. For instance, the absorption lines observed in the sun's spectrum due to the presence of various elements.

In summary, the types of line spectra include the line emission spectrum and the line absorption spectrum, both of which feature discrete lines corresponding to specific wavelengths associated with transitions between energy levels in atoms or molecules.

Multiple choice questions and answers on the spectrum and its types:

1. A source of light emits photons of different energies. Which of the following statement may best suits the light source?

1. The source is monochromatic
2. The source is polychromatic
3. The source has the highest energy than the surrounding
4. The existence of such a source is hypothetical

Answer:
The source is polychromatic
Explanation:
Sunlight or white light are examples of polychromatic sources which can emit light radiations of different wavelengths. Each wavelength denotes the photon energy associated with it.

2. The absorbed light radiations by an atom represented as ____________ in its absorption spectrum.

1. Bright lines
2. Colored bands
3. Dark lines
4. White lines

Answer:
Dark lines
Explanation:
The absorption spectrum is a pattern of dark lines or bands corresponding to the light radiations absorbed by the substance.

3. The atomic spectrum results from emitted----------------------

1. Electrons
2. Protons
3. Photons
4. Particle matter

Answer:
Photons
Explanation:
The absorbed or emitted photon energies of an atom generate its spectrum. And it is an arrangement of photon energies in the increasing or decreasing order of frequencies or wavelengths.

4. The amount of light scattered in the air varies_________________

1. Directly with its wavelength
2. Inversely with its wavelength
3. On the density of air
4. Remains constant at all wavelengths

Answer:
Inversely with its wavelength
Explanation:
The quantity of light scattered in air molecules depends on the wavelength of light. Shorter wavelength radiations scatter more than a longer wavelength. As a result, blue light scatters more than red light.

5. The frequency of _________ colored visible light radiation is maximum

1. Blue
2. Red
3. Violet
4. Navy blue

Answer:
Violet
Explanation:
The visible light (or) sunlight consists of seven colored radiations. They are violet, indigo, blue, green, yellow, orange, and red. Among them, violet-colored light radiation associate with the maximum photon energy. Hence, it has more frequency than the other.

6. The pattern of light radiations obtained due to dispersion is called__________

1. Spectrum
2. Spectrograph
3. Rainbow
4. Spectroscope

Answer:
Spectrum
Explanation:
The polychromatic radiation that passes through the prism splits into its constituent wavelengths. This process is known as dispersion. When we pass the dispersed light through a spectrometer exhibits the spectrum on the photographic plate. Alternatively, a spectrum is an arrangement of dispersed radiations of polychromatic light.

7. The absorption and re-radiation of light by the molecules of the atmosphere is __________

1. Dispersion
2. Scattering
3. Interference
4. Compression

Answer:
Scattering
Explanation:
The scientist Rayleigh predicted this phenomenon. According to him, scattering is the phenomenon of absorption of light and re-radiating it by the molecules of the atmosphere in different directions.

8. Rainbow forms when sunlight passes through the raindrops. The formation of a rainbow depends on which of the following phenomenon?

1. Dispersion
2. Total internal reflection
3. Refraction
4. Interference

• Options 1 and 2 are correct
• Options 1, 2, and 3 are correct
• Options 2 and 3 are correct
• All options are correct

Answer:
Options 1,2,and 3 are correct
Explanation:
The formation of a rainbow does depend on the interference phenomenon.

9. Which colored radiation of visible light scatters the least during sunrise and sunset?

1. Blue
2. Violet
3. Red
4. Crimson red

Answer:
Red
Explanation:
Out of the seven colored radiations of sunlight, red light has a longer wavelength. And we already discussed that shorter wavelength radiations scatter more than a longer wavelength. Therefore red light scatters the least.

10. VIBGYOR is________________

1. Continuous solar spectrum
2. Discontinuous solar spectrum
3. Hydrogen color spectrum
4. Atomic line spectrum

Answer:
Continuous solar spectrum
Explanation:
When sunlight passes through a prism shows a series of seven-colored diffused bands called the solar spectrum. It is an example of a continuous emission spectrum. The seven colors observed are Violet, Indigo, Blue, Green, Yellow, Orange, and Red, called VIBGYOR.

11. Atomic spectrum is called ___________

1. Band spectrum
2. Line spectrum
3. Color spectrum
4. Electromagnetic spectrum

Answer:
Line spectrum
Explanation:
An atomic spectrum consists of a series of bright or dark lines corresponding to the photon energies emitted or absorbed by the atom. The discrete lines of the atomic spectrum represent the quantized energy levels of the atom. Hence, it is also known as the line spectrum.

12. The phenomenon by which a polychromatic light splits into its component radiations of wavelengths while passing through the prism is ____________

1. Dispersion
2. Interference
3. Diffraction
4. Total internal reflection

Answer:
Dispersion
Explanation:
When the polychromatic radiation passes through a prism splits into its constituent wavelengths. This process is known as dispersion.

13. In a discharge tube, the sparking position of a gas depends upon

1. Type of gas chosen
2. Pressure and volume of a taken gas sample
3. Temperature of a gas
4. Pressure of gas and separation of electrodes

Answer:
The pressure of gas and separation of electrodes
Explanation:
In the discharge tube, the glowing or sparking of gas indicates the emission of light radiation due to electron transitions. The position of sparking thus depends upon the pressure of gas and the separation of electrodes.

14. The spectral lines of the solar absorption spectrum are known as_________________

1. VIBGYOR lines
2. Fraunhofer lines
3. Bright colored lines
4. Rainbow lines

Answer:
Fraunhofer lines
Explanation:
The solar spectrum shows dark absorption lines due to the absorption of light radiations by atoms in cooler regions of the atmosphere around the sun. These spectral lines are known as Fraunhofer lines, named after the scientist who discovered them.

15. Yellow lines of the solar emission spectrum occur at__________

1. 456 nm
2. 434 nm
3. 589 nm
4. 656 nm

Answer:
589 nm
Explanation:
The emission spectrum of the sodium atom shows two yellow-colored lines at wavelengths 5890 angstroms and 5896 angstroms.

16. Who discovered the word ‘spectrum’?

1. Fraunhofer
2. James Clerk Maxwell
3. Isaac Newton
4. Neil Bohr

Answer:
Isaac Newton
Explanation:
In 1666, Isaac Newton predicted that white light is a mixture of seven colored radiations of sun rays. He proved it by passing white light through the prism. He obtained a continuous rainbow of sunlight that he named spectrum. Hence, Sir Isaac Newton coined the term spectrum.

17. How many types of the spectrum are there?

1. 2
2. 3
3. 4
4. 1

Answer:
3
Explanation:
There exist three kinds of the spectrum. Emission spectrum Absorption spectrum Continuous spectrum

18. The range of wavelengths emitted by hot celestial objects such as stars is known as _________ spectrum

1. Emission spectrum
2. Absorption spectrum
3. Continuous spectrum
4. Discontinuous spectrum

Answer:
Continuous spectrum
Explanation:
The source of continuous spectra is hot celestial matter and stars. The range of wavelengths that occur in the continuous spectrum depends upon the temperature of the emitted light source.

19. Who discovered the helium atom from the solar spectrum?

1. Isaac Newton
2. Alfred Fowler
3. Johann Rydberg
4. Pierre Janssen

Answer:
Pierre Janssen
Explanation:
During the solar eclipse of 1868, Janssen discovered the Helium atom in the solar spectrum.

20. An atom shows ____________

1. Only emission spectrum
2. Only absorption spectrum
3. Both emission and absorption spectrum
4. Continuous and discontinuous spectrum

Answer:
Both emission and absorption spectrum
Explanation:
Each atom has its characteristic emission and absorption spectrum. The wavelengths of bright lines in the emission spectrum are the same as that of the dark lines in the absorption spectrum.

21. Which of the following statement is true?

1. Spectral studies help in the analysis of the composition of astronomical matter.

2. Using spectral line data, astronomers can determine the temperature and density of chemical substances in stars.

3. It tells about the magnetic field strength of stars.
4. All are correct

Answer:
All are correct
Explanation:
Spectral studies have much significance in the composition analysis of astronomical matter. The study of the emission and absorption spectra of the substance provides information on its structure. In particular, astronomers use this spectroscopic data to determine the constituents of stars and interstellar matter.

22. Atomic spectrum is evidence of ______________

1. Quantized angular momentum
2. Quantized energy levels of an atom
3. Spin-orbit coupling in an atom
4. Electric field strength of atoms

Answer:
Quantized energy levels of atoms
Explanation:
We know that the electron transitions of atoms and molecules give spectral emissions in various regions of the electromagnetic spectrum. Hence, the atomic spectrum is evidence for the existence of quantized electron orbits in the atom.

23. Blue colored light of street lights is due to________________

1. Mercury emission lines
2. hydrogen emission lines
3. sodium emission lines
4. copper emission lines

Answer:
Mercury emission lines
Explanation:
Streetlights seem blue due to intense photon emissions of Mercury occurring at 450nm.

24. The strongest atomic emission line from the sun is__________

1. Hydrogen-alpha line
2. Lyman alpha line
3. Helium emission line
4. Oxygen spectral line

Answer:
Lyman alpha line
Explanation:
The hydrogen electron transition from n=1 to n=2 proceeds with the emission of the Lyman alpha line at 121.5 nm. Due to the plethora of hydrogen in the universe, it is the strongest atomic emission from the sun.

25. Why do the gases in the discharge tube conduct electric current at low-pressure conditions?

1. At low pressure, the volume of gas increases. It increases the number of gas molecules in the discharge tube. More gas molecules mean more conducting power.
2. It increases the mobility of electrons. Hence freely moving electrons conduct electric current rapidly.
3. The colliding electrons acquire higher kinetic energy due to the increased mean free path. It leads to the ionization of atoms.
4. Low pressure means high temperature. So the kinetic energy of gas molecules increases. So they became good conductors of electricity.

Answer:
The colliding electrons acquire higher kinetic energy due to the increased mean free path. It leads to the ionization of atoms.
Explanation:
Gases are poor conductors of electricity. But they conduct electric current at low-pressure conditions. The number of gaseous molecules is less at low-pressure conditions allowing the free movement of electrons from the cathode to the anode in the discharge tube. Additionally, it reduces the number of collisions between the electrons and the gas molecules. Thus, the colliding electrons acquire high kinetic energy. It leads to the ionization of atoms.

Additional reference:

I believe our blog article, “Spectrum and Its Types,” covers the concepts you're looking for, including the different types of spectra like solar and atomic. Here are some additional references to help you understand the concept better.

1. Electromagnetic radiation and its types

2. An infographic on the applications of spectral studies