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Atomic Mass Definition: Definition of Weight and Mass, Units of Weight and Mass, Terms in Use, History, Difference Between Atomic and Molar Masses, Practice Problems, FAQS

Atomic Mass Definition: Definition of Weight and Mass, Units of Weight and Mass, Terms in Use, History, Difference Between Atomic and Molar Masses, Practice Problems, FAQS

There is a famous Egyptian story of King Osiris. Osiris is said to be the ultimate ruler of the kingdom of afterlife. How did one gain entry into the afterlife? There was a simple test that Osiris personally looked over. According to him, a good person has a pure heart, which would be lighter than a feather. If you passed that test, you went into the field of reeds or the Egyptian understanding of heaven. However, if you failed this test, you would be confined to Duat or the endless desert.

We know enough to understand that the above story is one of pure myth but has adaptations across most religions.

In practice, the human heart weighs anywhere between 250 and 350 grams! That is safely much heavier than a feather which typically weighs about 0.01 grams.

Speaking of light and heavy, the lightest known material in the universe is an up quark which is 1/477th of the mass of a hydrogen atom. Atomic mass as an idea itself was under scrutiny for about a quarter of a century and now we have much clarity about these numbers. However, such small masses deserve a discussion of their own. Let us try and dive deeper into what exactly atomic mass is.

TABLE OF CONTENTS:

Definition of Weight and Mass

Mass is a term indicating the total content in the system, while weight indicates the force experienced by the system due to gravity. Mass is fixed for a substance under all conditions of measurement while weight can change. Mass of a human is the same whether on earth or in a spacecraft. But,the weight of the human is not the same.

Units of Atomic weight and Atomic Mass

Atomic mass is, the mass contained in an atom. Nearly all of an atom's mass is made up of its protons, electrons and neutrons, with only a small amount of the mass being lost by way of nuclear binding energy.

Evidently this is a very small mass and hence, we use a different unit to talk about atomic masses. It is typically denoted with Da, which stands for Dalton which is equal to 1/12th of the mass of a Carbon-12 atom in ground state. It is also interchangeably called “amu” (atomic mass unit) or simply “u”. as an unified mass. So we can say the mass of a hydrogen atom is 1 amu or 1 u or 1 Da. As a result, the atomic mass expressed in Da has a value that is equivalent to the mass number, which is a whole number. As most of the elements exist with isotopes the weighted average of their atomic masses is referred to as atomic weight. It should be remembered atomic mass and atomic weight are not same if the atom exists in isotopes. For carbon, atomic mass is 12, while its atomic weight is 12.0107. But, the actual weight, for example, of a hydrogen atom in SI units is 1.6735575 × 10-27 kg.

Weight and Mass Terms In Use

  • Relative isotopic mass: By dividing the atomic mass of an isotope, ma, by the atomic mass constant, mu, a dimensionless value known as relative isotopic mass can be derived. Consequently, a carbon-12 atom's relative isotopic mass is 12 whereas its atomic mass is 12 Da by definition. The relative molecular mass is the total of all the atoms' respective isotopic masses. The relative isotopic mass and the atomic mass of an isotope refer to a particular isotope of an element. The elemental atomic mass, which is the average atomic mass of all isotope of an element, in terms of relative concentration.
  • Relative atomic mass: The ratio of the average mass of atoms of a chemical element to the atomic mass constant is known as the atomic weight, or relative atomic mass, which is a dimensionless physical quantity. The average atomic mass is the average weighted of the isotopes taken in proportion to their fractional existence compared to the atomic mass of carbon-12. The value in this comparison is dimensionless because it is the product of the two weights (having no unit). The name "relative" is also explained by this quotient: the sample mass value is compared to that of carbon-12.

The term "relative atomic mass" is more inclusive and can be used to refer to samples obtained from extraterrestrial environments as well as highly specific terrestrial environments. .

  • Standard atomic weight: A chemical element's standard atomic weight, denoted by the symbol Ar°(E), is the weighted arithmetic mean of all of its isotopes' relative isotopic weights multiplied by their relative abundances on Earth. For instance, isotope 63Cu (Ar = 62.929) constitutes 69% of thecopper on Earth, the rest being 65Cu (Ar = 64.927), soAr ° ( Cu ) = 0.69 × 62.929 + 0.31 × 64.927 = 63.55. This weighted mean has no dimensions since relative isotopic masses are dimensionless quantities. It can be multiplied by the dalton, also referred to as the atomic mass constant, to create a measure of mass (with dimension M). The standard atomic weight (Ar°) is the most prevalent and useful of the several variations of the concept of atomic weight (Ar, also known as relative atomic mass) used by physicists. The atomic weight of carbon in a specific bone from a certain archaeological site is an example of an element whose non-standardized atomic weight is unique to sources and samples. Such values are averaged to the range of atomic weights that a scientist may anticipate obtaining from numerous randomly selected samples from the Earth. The interval notation used for several standard atomic weight values is justified by this range. Eighty of the 118 known chemical elements have stable isotopes, and 84 of them have this value based on the Earth's environment.

History of Atomic Mass

John Dalton, Thomas Thomson, and Jöns Jakob Berzelius were the first researchers to calculate the relative atomic masses of atoms. According to Prout's idea, which was put forth in the 1820s, all atomic masses would turn out to be precise multiples of hydrogen. Relative atomic mass (or atomic weight) was first based on its mass relative to hydrogen, which was assumed to be equal to 1.00. However, Berzelius quickly demonstrated that this wasn't even close to being accurate; in fact, for some elements, like chlorine, the relative atomic mass, at 35.5, is almost exactly halfway between two integral multiples of hydrogen. This was later found out to be because of isotopes.

By using Avogadro's law, Stanislao Cannizzaro improved relative atomic masses in the 1860s. By comparing the vapour density of a group of gases with molecules containing one or more of the chemical element in question, he developed a law to determine the relative atomic masses of elements which states that: the amounts of the same element present in different compounds will be an integer multiples of its atomic weight.

There were two distinct atomic-mass scales used by chemists and physicists in the 20th century, up to the 1960s. While oxygen was assigned the 16 as the atomic mass, as the natural oxygen exists with isotopes of atomic mass 17 and 18,necessitating two distinct tables of atomic mass, having oxygen as standard was dropped.

The carbon-12, or 12C-based unified scale satisfied the physicists' requirement that the scale be based on a pure isotope and was numerically comparable to the scale used by chemists. As the "unified atomic mass unit," this was chosen.

The Dalton and symbol "Da" are the current International System of Units (SI) major recommendations for this unit's designation. The accepted names and symbols for the same unit are "unified atomic mass unit" and "u."

A secondary synonym for atomic weight called "relative atomic mass" was introduced in 1979 as a compromise. Twenty years later, the term "relative atomic mass" has replaced these synonyms as the preferred word.

Difference Between Atomic and Molar Masses

Difference between atomic and molar masses
MOLAR MASS ATOMIC MASS
Refers to mass of one mole of a substance Refers to the mass one atom of an element
SI unit is g/mol to use in higher calculations Measured in u or Da
It is defined as the mass of Avogadro number of atoms/molecules or compounds Defined as the mass of a single atom of any pure element.
Measurement given to compounds, atoms or molecules. Determined only in atoms.
Less accurate. Accurate to use in higher calculations.
Example: Mass of 1 mole of oxygen is 15.9994 grams. Therefore, the molar mass = 15.9994 g/mol Example: Atomic mass of oxygen-16 is 16 u.

Practice Problems

Q1. Some meteorites have an abundance ratio of O-18 :O-16 that is higher than the value determined by the average atomic mass of oxygen on earth?
Ans.
Is the average oxygen atomic mass in these meteorites higher, lower, or the same as that of an oxygen atom on Earth? Answer: Earth has a predominant isotopic presence of O-16, generally. This naturally leads to the atomic mass of oxygen being closer to 16. However, if O-18 is more abundant in a certain sample, we can say that the average atomic mass will be higher, or closer to 18 (as compared to a regular sample on earth).

Q2. What is the relative atomic mass of Lithium?
Ans:
Since nothing has been mentioned in the question, we take our sample to be from earth. The composition in this case will be approximately 7% Li-6 and 93% Li-7. Hence,

relative atomic mass will be: = 6.93

Q3. Earth has about 76% Cl-35 and 24% Cl-37. What is the standard atomic weight of Chlorine?
Ans:
To calculate the standard atomic weight, we need to look at the relative abundances of the isotopes. As we can see, relative abundance of Cl-35 is 0.76 and that of Cl-37 is 0.24.

Hence, we can write: Ar ° ( Cl ) = 0.76 × 35 + 0.24 × 37 = 35.48

Q4. What is the atomic mass of potassium? What is its relative isotopic mass?
Ans:
The atomic mass of Potassium is 39 u. Since the abundance of K-39 on earth is about 93%, we can assume that to be the most abundant isotope. Hence, since mass number is 39, we can say relative isotopic mass of potassium is 39 .

Frequently Asked Questions

Q1. What is mass number?
Ans:
Mass number is the sum of the number of protons and neutrons in a certain nuclide. Alternatively, it can be understood as the number of nucleons. A quick trick is that mass number is generally equal to the numerical value of isotopic mass.

Q2. What is the best definition of atomic weight?
Ans:
The term "atomic weight" refers to an atom's overall weight. With the addition of a small amount from the electrons, it is roughly equal to the sum of the protons and neutrons. It is a dimensionless quantity as it is a relative quantity.

Q3. What is the heaviest known element?
Ans:
Uranium is the heaviest naturally occurring nuclide and its atomic number is 92. However, we know that 118 elements have been discovered so far. Scientists have been able to synthesise Oganesson, which is currently the heaviest element we know of.

Q4. Which is the densest substance in the universe?
Ans:
It is well known that Osmium is the densest element found on Earth with a density of 22 g/cc. However, the densest material in the universe is actually a neutron star which is known to be capable of warping space time due to immense gravitational pull.

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