Massa atom relatif (Ar)

Summary

This video explains the concept of relative atomic mass (Ar), demonstrating how to determine the mass of an atom using a standard reference. Initially, hydrogen was used as the standard for atomic mass, but later, Carbon-12 (C-12) became the global standard due to its stability. The video also discusses how to calculate relative atomic mass using examples, including a detailed explanation of isotopes and their abundance. The tutorial further illustrates how mass spectrometry helps in the analysis of elements and their isotopes, providing a practical guide to understanding atomic masses in chemistry.

Takeaways

• 😀 The topic of the video is about determining relative atomic mass (Ar), also known as atomic weight (KR).
• 😀 Mass of a substance can be measured using a balance, and the unit for mass is grams (g).
• 😀 The challenge arises when determining the mass of an atom, which is too small to be seen with the naked eye.
• 😀 Atomic mass can be determined by comparing it to a standard atomic mass of another element, initially hydrogen.
• 😀 Hydrogen was once used as the standard for atomic mass, but later, oxygen with a mass of 16 SMA was used.
• 😀 With the invention of the mass spectrometer, it became possible to detect the mass of individual elements in compounds.
• 😀 Elements can have multiple isotopes, which are atoms of the same element with different numbers of neutrons.
• 😀 Carbon-12 (C-12) was established as the standard for atomic mass due to its stability and was set as the reference in 1961.
• 😀 The formula for calculating relative atomic mass (Ar) is: Ar = (mass of element / (1/12 of C-12 mass)).
• 😀 An example problem shows how to calculate the relative atomic mass using known values for average atomic mass and C-12 mass.
• 😀 The relative atomic mass can also be calculated using isotopic mass and abundance, as shown in the example of boron with isotopes B10 and B11.

Q & A

What is relative atomic mass (AR)?

Relative atomic mass (AR) is the mass of an atom of an element compared to the mass of a carbon-12 atom, which is defined as having a mass of exactly 12 atomic mass units (AMU).

Why was carbon-12 (C12) chosen as the standard for atomic mass?

Carbon-12 (C12) was chosen as the standard because it is the most stable isotope of carbon, and it was later established in 1961 as the standard for defining atomic mass.

How was the atomic mass of elements determined before the use of C12?

Initially, hydrogen was used as the standard for determining atomic mass, but later oxygen was also considered before carbon-12 became the definitive standard.

What is a mass spectrometer and what role does it play in atomic mass determination?

A mass spectrometer is an instrument that can detect the mass of atoms and isotopes in elements or compounds. It helped discover that atoms can have multiple isotopes, each with different masses.

What is an isotope?

An isotope refers to atoms of the same element that have the same number of protons but different numbers of neutrons, resulting in different atomic masses.

What is the formula for calculating relative atomic mass (AR)?

The formula for calculating relative atomic mass is: AR = (mass of one atom of element X) / (1/12 * mass of one atom of C12).

How do you calculate the relative atomic mass when dealing with isotopes?

The relative atomic mass can be calculated by multiplying the mass of each isotope by its percent abundance and summing these values. The formula is: AR = Σ (mass of isotope × percent abundance).

How do you calculate the relative atomic mass of boron using isotopes?

To calculate the relative atomic mass of boron, you would multiply the mass of each isotope (B10 and B11) by its respective abundance, then sum the results. For example, AR = (10.013 × 19.10%) + (11.01 × 80.90%).

What does the term 'percent abundance' refer to in the context of isotopes?

Percent abundance refers to the proportion of each isotope in a naturally occurring sample of an element, usually expressed as a percentage.

Why is the relative atomic mass of elements like boron not a whole number?

The relative atomic mass of elements like boron is not a whole number because it is a weighted average of the masses of different isotopes, each contributing according to its abundance in nature.