Atomic Mass Definition
o The mass of atom can be defined as the 1/12 of the mass of an atom of carbon-12 or 1.660538921×10-24 grams.
o The atomic mass unit is not same as the mass of the proton or neutron.
o The SI unit of mass is represented in Kilogram (kg), but atomic mass is usually measured in Dalton (symbol D, or u).
Atom
o It is defined as the smallest unit of a matter that may or may not exist independently and shows all the properties of that given element.
o Most of the space in an atom is empty. The small area which is occupied by positively charged sphere is named as called nucleus.
o The nucleus comprises two particles named as Proton and Neutron. Nucleus is surrounded by negatively charged particle named as Electron.
o The whole mass of the atom is present in nucleus. Electron are attracted towards nucleus by electric force.
o Electrons also revolve around nucleus in fixed path called shell or orbit.
Atomic Number
o The number of protons present in nucleus of the atom is known as atomic number.
o It is denoted by symbol Z. A neutral atom contains equal number of protons and electrons.
Mass Number
o The sum of total number of protons and neutrons is called mass number.
Isotopes
The different form of element having same atomic number but different mass number are named as isotopes. Isotopes shows the following properties as mentioned below:
1. They have same number of protons.
2. Difference in mass is due to different number of neutrons.
3. Isotopes also displays different physical properties but their chemical properties are usually always same.
o For Instance: The hydrogen atom has three isotopes which are Protium, Deuterium, Tritium. Protium has 1 proton and no neutron, deuterium has 1proton and 1 neutron and tritium have 1 proton and 2 neutrons.
o Isotopes are shown as: 1H1, 1H2, 1H3
Average Atomic Mass
o Average atomic mass of an element is defined as the sum of the masses of the isotopes of the element multiplied by its natural abundance on earth.
o Natural abundance is defined as availability of a particular isotopes that are found in nature.
Atomic Mass vs Average Atomic Mass
o The atomic mass is the mass of a single atom, which can be only one isotope at a time, and is not an average of all the masses of isotopes present in nature.
o An element has only one naturally occurring dominant isotopes, the difference between the atomic mass and average atomic mass can either be small or negligible.
o Average atomic mass is frequently close to whole number, but never exactly a whole number, because of the following reason given below:
1. Different masses of protons and neutrons and different isotopes have different ratio of their protons and neutrons.
2. Atomic masses are reduced to some extents by their binding energy.
3. The relative atomic mass is thought to be precise for the planet Earth and is a standardized number.
o An average atomic mass is true for a given sample, as the value can change over geologically and certain process change the ratio of presence of isotopes.
Average Atomic Mass Calculation
o Average Atomic Mass =f1M1 + f2M2 +f3M3 +…………… + fnMn
f = is defined as the fraction of the natural abundance of the isotopes
M = is the mass number (weight) of the isotopes.
o In the periodic table atomic number is given at the base of each element symbol. If data is available for the natural abundance of various isotopes, it is very easy to calculate the average atomic mass.
o In helium, only one isotope of Helium-3 is present for every million isotopes of Helium-4; so, the average atomic mass is very close to 4 amu (4.002602).
Average Atomic Mass Calculation Steps
1. Initially, convert the percentage into fraction by dividing the number by 100.
2. Now calculate the mass number by adding number of neutrons to number of protons.
3. Multiply the fraction of all the given isotopes with their respective mass number.
4. Lastly, add up all to get average atomic mass of the given element.
Average Atomic Mass Example
Calculate the average atomic mass of chlorine
o Chlorine comprises of two isotopes one with 18 neutrons (75.77 percent of natural chlorine atoms) and other with 20 neutrons (24.23 percent of natural chlorine atoms).
o The atomic number of chlorine is 17 as it has 17 protons in its nucleus.
Percentage of first isotope 17 Cl35 is equal to 75.77 or 75/100
Percentage of second isotope 17 Cl37 is equal to 24.23 or 25/100
Mass number of first isotope is equal to 35
Mass number of second isotope is equal to 37
Average atomic mass = 35×75/100 + 37×25/100
= 142/4
= 35.5u
o The method mentioned above is also used to calculate average atomic mass of other elements. For all the calculation in chemistry average atomic mass is very necessary.
Atomic Mass Citations
- Proton-neutron interactions and the new atomic masses. Phys Rev Lett . 2005 Mar 11;94(9):092501.
- Direct and indirect methods for molar-mass analysis of fragments of the capsular polysaccharide of Haemophilus influenzae type b. Anal Biochem . 1997 Aug 1;250(2):228-36.
- Molecular mass and location of the most abundant peak of the molecular ion isotopomeric cluster. J Mol Model . 2005 Sep;11(4-5):271-7.
- Effective atomic numbers, water and tissue equivalence properties of human tissues, tissue equivalents and dosimetric materials for total electron interaction in the energy region 10 keV-1 GeV. Appl Radiat Isot . 2014 Dec;94:1-7.