Average Atomic Mass

An element’s atomic mass (also known as atomic weight) is the weighted average mass of the atoms. Average masses are commonly stated in unified atomic mass units (u), where 1 u equals one-twelfth the mass of a neutral carbon-12 atom.

Obtaining the Average Atomic Mass

Consider the following about carbon: Carbon-12 and carbon-13 have natural abundances of 98.90% and 1.10%, respectively. Carbon-12 has an atomic mass of 12.00000 amu (atomic mass unit), while carbon-13 has an atomic value of 13.00335 amu.

The percentage abundance must first be converted to decimals before the average atomic mass can be calculated. This means that when the abundances are combined, they must equal one. Divide the percent abundance by 100 to get the decimal equivalent.

Carbon-12 has a fractional abundance of 0.9890, while carbon-13 has a fractional abundance of 0.0110.

The fractional abundance is then multiplied by the mass and added.

Therefore, carbon’s average atomic mass is 12.01 amu. This is the value shown on the periodic table. It makes logical sense that the average atomic mass is closer to 12 because carbon-12 has a significantly higher fractional abundance than carbon-13.

Average Atomic Mass Formula

How to calculate the Average Atomic Mass?

To understand the average amount of an average atomic mass of an element can be done by multiplying the sum of isotopes and its natural abundance (the decimal associated with the percent of atoms of that element that are of a given isotope).

• Average atomic mass = f1M1 + f2M2 +… + fnMn, where f is the natural abundance percentage and M is the mass number of the isotope (weight).

• The average atomic mass of an element is generally located under the elemental symbol on the periodic table. 

• Calculating the average atomic mass is simple when data on the natural abundance of different isotopes of an element is available.

• Regarding helium, one isotope of Helium-3 exists for every million isotopes of Helium-4; hence, the average atomic mass is extremely near to 4 amu (4.002602 amu).

• There are two primary isotopes of chlorine: including one 18 neutrons (75.77 percent of natural chlorine atoms) and one with 20 neutrons (75.77 percent of natural chlorine atoms) (24.23 percent of natural chlorine atoms). 

• Chlorine has an atomic number of 17. (it has 17 protons in its nucleus).

• Transform the proportions into fractions before calculating the average mass (divide them by 100). Determine the mass numbers after that. 

• The chlorine isotope with 18 neutrons has a mass number of 35 amu and an abundance of 0.7577. Multiply the proportion by the mass number for each isotope to find the average atomic mass and then mix them together.

• Chlorine’s average atomic mass is (0.7577 35 amu) + (0.2423 37 amu) = 35.48 amu.

• One instance is calculating the atomic mass of boron (B), which contains two isotopes: B-10, which has a natural abundance of 19.9%, and B-11, which has an abundance of 80.1 percent. Therefore,

• Boron’s average atomic mass is (0.199 10 amu) + (0.801 11 amu) = 10.80 amu.

• We usually utilise average atomic masses when doing mass calculations using elements or compounds (combinations of elements).

Important Considerations about average atomic mass

• The nucleus of an element may include various numbers of neutrons, but it always contains the same amount of protons.
• Isotopes are variations of an element that have differing neutron masses.
• To calculate the Average atomic mass of an element, you can add isotopes and multiply it by the abundant element of earth availability. 
• Always use the periodic table’s average atomic weight when making mass calculations with elements or compounds.
• The mass number of an atomic nucleus is the total number of protons and neutrons in the nucleus.
• The quantity of a certain isotope found in nature on Earth is called natural abundance.
• The atomic number of an element establishes its identity and denotes the number of protons in the nucleus of one atom.
• For example, the lightest element, hydrogen, will always have one proton in its nucleus. Helium always has two protons in its nucleus.

Isotopes

Atoms of an element, on either hand, may have varying numbers of neutrons in their nuclei. For example, stable helium atoms, including one or two neutrons, exist, but both atoms have two protons. Isotopes are different types of helium atoms with different masses (3 or 4 atomic mass units). Every isotope’s mass number is the sum of the number of protons and neutrons in the nucleus. This is because each proton and neutron weighs one atomic mass unit (amu). The mass of an atom may be calculated by multiplying protons and neutrons by one atomic mass unit—each physical quality as a collection of isotopes.

The name ‘isotope’ is derived from the Greek terms ‘isos’ (meaning ‘same’) and ‘topes’ (meaning ‘place’) because elements with different subatomic structures can maintain the same position on the periodic table.

Conclusion

It is critical to remember that while the average atomic mass of magnesium is 24.31 amu and the average atomic mass of carbon is 12.01 amu, these values do not represent the mass of any individual atom. For example, if you could pick up a single carbon atom (imagine it’s feasible! ), it would be either 12.00 amu or 13.000335 amu. It would not have been 12.01 a.m.

Similarly, if you picked up any solitary magnesium atom, it would be either 23.98504 amu, 24.98583 amu, or 25.98259 amu. The average atomic mass is just a method for scientists to account for the stable isotopes of an element found on Earth. It does not denote the mass of a single atom unless the element has one stable isotope.