Showing posts with label Atomic Mass. Show all posts
Showing posts with label Atomic Mass. Show all posts

Thursday, November 8

Standard for Atomic Mass

Introduction to standard for atomic mass:

Atomic mass is the mass of a single atom and it is usually expressed in atomic mass units (or amu). Most of the atomic mass is concentrated in the nucleus, in the protons and neutrons contained. Since each proton or neutron weighs about 1 amu, the atomic mass is always very close to the mass number (or the total number of protons and neutrons in the nucleus). The atomic masse is usually determined by mass spectrography and the atoms of the isotopes of an element have the same atomic mass. It has been determined with great relative accuracy, but its absolute value is less certain.Looking out for more help on Carbon Atomic Mass in algebra by visiting listed websites.

Standard Atomic Mass

Atomic mass is the mass of an atom or molecule on a scale where the mass of a carbon-12 (12C) atom is exactly 12. Atomic mass or the mass of any atom is approximately equal to the total number of its protons and neutrons multiplied by the atomic mass unit (u) which is 1.660539 × 10-24 gram since electrons are much lighter. The atoms do not differ much from this simple formula, but only by less than 1%.

Atomic mass unit or amu

Atomic mass unit or amu is the unit defined as exactly 1/12th of the mass of a carbon-12 atom. The carbon-12 is an isotope of carbon with six protons and six neutrons in its nucleus. One amu is equivalent to 1.66 × 10−24 grams or 1.66 × 10--27 kg, approximately. The masses of individual atoms are expressed in terms of Atomic Mass Unit or amu. The standard is the unit of mass equals to one-twelfth the mass of the carbon atom, having as nucleus the isotope with mass number 12. Atomic mass unit has the abbreviation as amu and also known as dalton. I am planning to write more post on Rotational Energy, The first Law of Thermodynamics. Keep checking my blog.

Atomic Mass of Isotopes:

The discovery of isotopes complicated the situation. For example, pure oxygen is composed of a mixture of isotopes in nature and some of the oxygen atoms are more massive than others. This was not problem for the calculations since the relative abundance of the isotopes remained as the same [or constant]. In such cases, the average mass of its atoms expressed in amu is taken as the relative atomic mass of a chemical element.