What's the Difference between Mass Number and Atomic Mass?

Summary

This video explains the difference between mass number and atomic mass. Mass number refers to the total number of protons and neutrons in an atom's nucleus, providing an estimate of the atom's weight in atomic mass units (amu). In contrast, atomic mass is the weighted average mass of an element’s isotopes, considering the relative abundance of each isotope. Using Boron as an example, the video explains how to calculate atomic mass, emphasizing the influence of isotope abundance on the final average. The key distinction is that mass number applies to individual isotopes, while atomic mass considers all isotopes of an element.

Takeaways

• 😀 Mass number is the sum of protons and neutrons in an atom's nucleus.
• 😀 Atomic mass is also known as average atomic mass, relative atomic mass, or atomic weight.
• 😀 Atomic mass is not the same as mass number; mass number is a whole number while atomic mass is an average.
• 😀 Mass number helps estimate the weight of an atom in atomic mass units (amu).
• 😀 Atomic mass takes into account the abundance of different isotopes of an element.
• 😀 The number of protons in an atom determines the element it is, and thus the element's position on the periodic table.
• 😀 Isotopes are versions of an element that have the same number of protons but different numbers of neutrons.
• 😀 The script uses Boron as an example, showing two isotopes: Boron-10 and Boron-11.
• 😀 Boron-10 and Boron-11 make up about 20% and 80% of Boron atoms, respectively.
• 😀 To calculate the atomic mass, the relative abundance of isotopes is multiplied by their respective mass numbers, then summed.
• 😀 The atomic mass of Boron (10.8 amu) is closer to Boron-11 because it is more abundant, reflecting the weighted average of isotopes.

Q & A

What is the primary difference between mass number and atomic mass?

Mass number refers to the sum of protons and neutrons in the nucleus of an atom, while atomic mass is the weighted average of the masses of all isotopes of an element, considering their relative abundances.

What does the mass number tell us about an atom?

The mass number tells us the total number of protons and neutrons in the nucleus of an atom, which gives an estimate of how much the atom weighs in atomic mass units (amu).

Why are electrons ignored when calculating an atom's mass?

Electrons are ignored in mass calculations because they are very light and their mass is negligible compared to that of protons and neutrons.

What is the mass number of an atom with 5 protons and 6 neutrons?

The mass number is 11, calculated by adding the number of protons (5) and neutrons (6).

How does atomic mass differ from mass number when considering isotopes?

Atomic mass takes into account the relative abundance of each isotope of an element and averages their masses, while mass number is specific to each isotope based on its protons and neutrons.

Why can't we simply add the mass numbers of isotopes and divide by two to find atomic mass?

Because atomic mass involves considering the different abundances of isotopes, simply averaging the mass numbers would be inaccurate. The actual atomic mass is a weighted average based on the percentage of each isotope.

How is the atomic mass of Boron calculated in the script?

The atomic mass of Boron is calculated by taking into account that 20% of Boron atoms are Boron-10 (10 amu) and 80% are Boron-11 (11 amu). The weighted average is: (0.20 * 10) + (0.80 * 11) = 10.8 amu.

What is the mass number of Boron-10 and Boron-11?

Boron-10 has a mass number of 10 (5 protons + 5 neutrons), and Boron-11 has a mass number of 11 (5 protons + 6 neutrons).

Why does the atomic mass of Boron end up being closer to 11 than 10?

Because 80% of Boron atoms are Boron-11, which has a mass of 11 amu, the atomic mass is weighted more heavily toward 11 amu, making the average closer to 11 than 10.

What is the significance of the atomic mass value listed on the periodic table?

The atomic mass value on the periodic table reflects the average mass of all naturally occurring isotopes of an element, taking into account their relative abundances. It is typically not a whole number because it represents a weighted average.