Difference Between Emission and Absorption Spectrum in Physics
Emission spectrum is a core topic in modern Physics that helps us understand the behavior of atoms and molecules when they emit light. When atoms absorb energy, their electrons get excited to higher energy levels. As these electrons return to their lower energy states, they emit the extra energy in the form of photons.
- The resulting light, when passed through a prism or a spectroscope, splits into distinct lines of specific wavelengths, forming the emission spectrum.
- Each element’s emission spectrum is unique, acting like its fingerprint.
- There are three main types of spectra: line spectrum, continuous spectrum, and band spectrum.
- A line spectrum consists of sharp, distinct lines and is observed when gaseous atoms are excited.
- A continuous spectrum appears as an unbroken band of colors seen from heated solids or dense materials.
- The band spectrum is formed by groups of closely spaced lines, typically seen with molecules rather than single atoms.
Definition, Types, and Explanation
An emission spectrum is the set of wavelengths emitted by atoms or molecules when electrons fall from higher to lower energy levels. When white light is passed through a dispersing medium like a prism, it splits into its constituent colors, forming a spectrum. However, when atoms are excited and their emitted light is passed through a prism, we obtain an emission spectrum of discrete wavelengths.
For example, when copper is heated in a flame, the resulting green light seen is part of its emission spectrum. Hydrogen, when excited, emits light in multiple series such as Lyman, Balmer, Paschen, Brackett, and Pfund.
Key Formulas and Their Applications
The main formula used to determine the wavelengths of emission lines in hydrogen is:
Where:
- λ = wavelength of emitted light
- RH = Rydberg constant (1.097 × 107 m-1)
- n1 = lower energy level
- n2 = higher energy level (n2 > n1)
This formula allows you to calculate the exact position of spectral lines for hydrogen and is widely used for theoretical as well as practical studies of atomic spectra.
Step-by-Step Problem Solving Approach
To solve problems based on the emission spectrum:
- Identify the initial (n2) and final (n1) energy levels involved in the electron transition.
- Use the formula 1/λ = RH(1/n12 – 1/n22).
- Substitute values of n1, n2, and RH.
- Calculate the difference and find the wavelength λ of the emitted photon.
This method can be used for any emission spectrum question, especially those involving hydrogen or hydrogen-like atoms.
Example: Wavelength Calculation in Hydrogen Spectrum
Suppose an electron in a hydrogen atom drops from n = 3 to n = 2 (Balmer series):
- Use n1 = 2, n2 = 3.
- 1/λ = 1.097 × 107 × (1/4 – 1/9)
(¼ – 1/9 = (9 – 4)/36 = 5/36) - 1/λ = 1.097 × 107 × 5/36 = 1.523 × 106 m-1
- λ = 1 / 1.523 × 106 = 6.57 × 10-7 m = 657 nm
This line lies in the visible red region of hydrogen’s spectrum.
| Types of Spectrum | Description |
|---|---|
| Line Spectrum | Discrete sharp lines from excited atoms (e.g., hydrogen, sodium) |
| Continuous Spectrum | Unbroken range of colors, produced by hot solids/liquids |
| Band Spectrum | Closely spaced lines forming bands, characteristic of molecules |
| Hydrogen Spectral Series | n1 | n2 (n2 > n1) | Region |
|---|---|---|---|
| Lyman Series | 1 | 2, 3, ... | Ultraviolet |
| Balmer Series | 2 | 3, 4, ... | Visible |
| Paschen Series | 3 | 4, 5, ... | Infrared |
| Brackett Series | 4 | 5, 6, ... | Infrared |
| Pfund Series | 5 | 6, 7, ... | Infrared |
| Formula | What It Represents |
|---|---|
| 1/λ = RH(1/n12 – 1/n22) | Wavelength of hydrogen spectral lines |
| ΔE = hν = hc/λ | Energy released or absorbed during electron transition |
| f = c/λ | Relationship between light frequency and wavelength |
Applications and Next Steps
The emission spectrum allows scientists to identify elements both in laboratories and in distant stars. Spectral lines also confirm the quantum nature of electrons in atoms. For further study, explore other related Physics topics such as Atomic Spectra, Line Spectrum, the Rydberg Constant, and the Bohr Model of Hydrogen Atom.
Practice more questions on these concepts and solve previous year numerical for deeper understanding. To continue building strong fundamentals, refer to practice sets, examples, and topic-specific lessons available across Vedantu’s Physics resources.
Summary
Understanding the emission spectrum strengthens your grasp of atomic structure, quantum transitions, and related Physics problems. Mastering this topic helps you excel in Physics exams while building your analytical skills in modern Physics. Regular problem-solving and revision using Vedantu’s structured resources are recommended for efficient preparation.
FAQs on Emission Spectrum Explained: Physics Concepts, Formulas & Uses
1. What is an emission spectrum in simple terms?
An emission spectrum is the set of specific wavelengths (colors) of light that an atom or molecule emits when its electrons drop from higher to lower energy levels.
• It appears as bright lines or bands against a dark background.
• Each element has a unique emission spectrum, acting like a 'fingerprint' for identification.
• The emission spectrum arises due to quantum transitions within the atom.
2. How is the emission spectrum different from the absorption spectrum?
The emission and absorption spectrum are opposites in how they appear and are formed:
• The emission spectrum shows bright, colored lines on a dark background as atoms release energy.
• The absorption spectrum shows dark lines on a continuous bright background when atoms absorb specific wavelengths of light.
• Emission involves electrons moving to lower energy levels; absorption involves electrons absorbing energy to move to higher levels.
3. What causes the formation of an emission spectrum?
An emission spectrum is formed when electrons in an atom or molecule return from a higher energy level to a lower one, releasing energy as photons.
• The specific wavelength of each spectral line depends on the energy difference between the levels.
• Each transition produces a photon of a particular color (wavelength), creating the pattern of the emission spectrum.
4. What are the types of spectra seen in Physics?
The three main types of spectra are:
• Continuous spectrum: All wavelengths are present without interruption (e.g., white light from the sun).
• Line spectrum (atomic emission spectrum): Consists of distinct lines, each corresponding to a specific wavelength and produced by gaseous atoms.
• Band spectrum: Appears as groups of closely spaced lines, typically observed in molecular spectra.
5. What is the main formula for the hydrogen emission spectrum?
The main formula for the hydrogen spectral lines is:
1/λ = RH (1/n12 – 1/n22)
• RH = Rydberg constant = 1.097 × 107 m-1
• n2 > n1; n1 and n2 are principal quantum numbers
• This formula helps calculate the wavelength (λ) for different spectral series (Lyman, Balmer, etc.)
7. What is the significance or use of emission spectra?
Emission spectra are important for several reasons:
• Identifying elements (every element has a unique spectrum)
• Studying atomic structure and quantum mechanics
• Applications in astronomy to determine the composition of stars
• Used in forensic science, lighting (e.g., neon lamps), and lasers
• Helps in analyzing unknown substances in laboratories
8. What is meant by spectral lines and why are they unique for each element?
Spectral lines are individual bright or dark lines in a spectrum, each corresponding to a specific wavelength emitted or absorbed by an atom.
• These lines are unique because every element has a distinct set of energy levels.
• Transitions between these levels produce a signature pattern, similar to a fingerprint, that can be used to identify the element.
9. How do you calculate the energy of emitted light in an emission spectrum?
The energy of emitted light (photon) is given by:
• ΔE = hν = hc/λ
where:
• ΔE = energy difference between levels
• h = Planck's constant (6.626 × 10-34 Js)
• ν = frequency
• c = speed of light (3 × 108 m/s)
• λ = wavelength of emitted light
10. What is a line spectrum and which elements produce it?
A line spectrum is a type of emission or absorption spectrum consisting of discrete lines at specific wavelengths.
• Produced mostly by gaseous elements (e.g., hydrogen, helium, mercury) and atoms in low-pressure discharge tubes
• Each line corresponds to a quantum electronic transition within the atom
11. Why does hydrogen have different spectral series (Lyman, Balmer, etc.)?
Hydrogen's different spectral series occur because electrons drop to different lower energy levels (n1), emitting photons of different energies.
• Lyman: Electron drops to ground level (n=1), emitting ultraviolet light
• Balmer: Electron drops to n=2, emitting visible light
• Paschen, Brackett, Pfund: Electron drops to n=3, 4, or 5, emitting infrared light
12. How do emission spectra help in astronomy?
Emission spectra enable astronomers to identify the elements present in stars and galaxies.
• The unique set of spectral lines seen from a star reveals its chemical composition.
• Helps estimate temperature, movement (via redshift), and physical conditions of astronomical objects.
