2.7.3 Applications of Mass Spectrometry

This topic explains the wide range of applications of mass spectrometry, including determining isotopic abundance, identifying unknown compounds, and calculating average atomic mass.

Definition

Applications of Mass Spectrometry: The practical uses of mass spectrometry in chemistry, biology, physics, and other scientific fields.

Main Idea

Mass spectrometry has a wide range of applications, including determining the relative abundance of isotopes of an element, identifying unknown compounds, quantifying the amounts of specific substances, elucidating molecular structures, and studying biomolecules like proteins and nucleic acids. It is also used to calculate the average atomic mass of an element from its isotopic composition.

Core Concept

• **Mass information:** The mz ratio tells us the mass of the ions.
• Abundance information: The peak heights tell us the relative abundance of each ion.
• Structural information: The fragmentation pattern tells us about the structure of the molecule.

Detailed Explanation

Applications of Mass Spectrometry

Mass spectrometry has a wide range of applications, including:

• Different isotopes of an element have different masses.
• The mass spectrum shows peaks for each isotope.
• The relative heights of these peaks give the relative abundance of each isotope.
• Example: Chlorine has two isotopes: Cl-35 (75.76%) and Cl-37 (24.24%).

• The molecular ion peak gives the molecular mass of the compound.
• The fragmentation pattern gives information about the structure of the compound.
• By comparing the mass spectrum with a library of known spectra, the compound can be identified.

• The height of a peak is proportional to the amount of that substance.
• By measuring the peak height, the amount of the substance can be determined.

• The fragmentation pattern provides information about the bonds in the molecule.
• Different functional groups fragment in different ways.
• This helps determine the structure of the molecule.

• Mass spectrometry is used to determine the mass and sequence of proteins.
• It is also used to study nucleic acids (DNA, RNA) and other biomolecules.

Calculating Average Atomic Mass

The average atomic mass of an element can be calculated using the following formula, taking into account the relative abundance of all isotopes:

Average atomic mass = (Mass No. × relative abundance) + (Mass No. × relative abundance) + …100

• Usually, the atomic masses written in the periodic table are not integral numbers.
• These non-integer atomic masses are the result of the inclusion of the relative abundance of different isotopes of an element.
• Relative abundance is usually taken in percentage units.
• They are multiplied by the mass numbers.
• The summation is divided by 100.

Example: Chlorine Isotopes

• Chlorine-35 with an abundance of 75.76%
• Chlorine-37 with an abundance of 24.24%

To calculate the relative atomic mass of chlorine:

Average atomic mass = (35 × 75.76) + (37 × 24.24)100

= 2651.6 + 896.88100

= 3548.48100

= 35.48 amu

Therefore, the relative atomic mass of chlorine is 35.48 amu.

Example: Magnesium Isotopes

Calculate the average atomic mass of Mg using the following mass spectrogram data:

Average atomic mass = (24 × 79) + (25 × 10) + (26 × 11)100

= 1896 + 250 + 286100

= 2432100

= 24.32 amu

Intermediate Level

Applications

Calculating Average Atomic Mass

Average atomic mass = ∑ (Mass No. × Relative Abundance)100

Example: Chlorine (35 × 75.76) + (37 × 24.24)100 = 35.48 amu

Example: Magnesium (24 × 79) + (25 × 10) + (26 × 11)100 = 24.32 amu

Important Points

• 1. Mass spectrometry determines the relative abundance of isotopes.
• 2. It identifies unknown compounds from their mass spectrum.
• 3. It quantifies the amounts of specific substances.
• 4. It elucidates molecular structures from fragmentation patterns.
• 5. It studies biomolecules like proteins and nucleic acids.
• 6. Average atomic mass = weighted average of isotope masses.
• 7. Example: Chlorine → 35.48 amu; Magnesium → 24.32 amu.

Remember

• - Applications: Isotope abundance, unknown compounds, quantification, molecular structure, biomolecules
• - Average atomic mass: ∑ (Mass No. × Relative Abundance)100
• - Examples: Chlorine (35.48 amu), Magnesium (24.32 amu)

Common Mistakes

• 1. Forgetting to divide by 100: Relative abundance is in percentage.
• 2. Mixing up mz ratio with mass: mz is mass-to-charge ratio.
• 3. Ignoring the charge: Ions with different charges have different mz ratios.
• 4. Confusing mass spectrum peaks with isotopes: Each peak corresponds to a specific ion.

Quick Revision

• Mass spectrometry determines the relative abundance of isotopes.
• It identifies unknown compounds from their mass spectrum.
• It quantifies the amounts of specific substances.
• It elucidates molecular structures from fragmentation patterns.
• It studies biomolecules like proteins and nucleic acids.
• Average atomic mass = weighted average of isotope masses.
• Example: Chlorine → 35.48 amu; Magnesium → 24.32 amu.

60-Second Revision

• Applications: Isotope abundance, unknown compounds, quantification, molecular structure, biomolecules
• Average atomic mass: ∑ (Mass No. × Relative Abundance)100
• Examples: Chlorine (35.48 amu), Magnesium (24.32 amu)

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