Monday, 6 June 2011

Periodic table:Atomic mass


Element Groups (Families)
Alkali EarthAlkaline EarthTransition Metals
Rare EarthOther MetalsMetalloids
Non-MetalsHalogensNoble Ges



Conversion factor between atomic mass units and grams


Conversion factor between atomic mass units and grams

                        The standard scientific unit for dealing with atoms in macroscopic quantities is the mole (mol), which is defined arbitrarily as the amount of a substance with as many atoms or other units as there are in 12 grams of the carbon isotope C-12. The number of atoms in a mole is calledAvogadro's number, the value of which is approximately 6.022 × 1023 mol−1. One mole of a substance always contains almost exactly therelative atomic mass or molar mass of that substance (which is the concept of molar mass), expressed in grams; however, this is almost never true for the atomic mass. For example, the standard atomic weight of iron is 55.847 g/mol, and therefore one mole of iron as commonly found on earth has a mass of 55.847 grams. The atomic mass of an 56Fe isotope is 55.935 u and one mole of 56Fe will in theory weigh 55.935g, but such amounts of pure 56Fe have never been found on Earth.
The formulaic conversion between atomic mass units and SI mass in grams for a single atom is:
1\ {\rm{u}}={M_{\rm{u}} \over N_{\rm A}}\ = {{1\ \rm{g/mol}} \over N_{\rm A}}
where Mu is the Molar mass constant and NA is the Avogadro constant.         

Atomic mass

Atomic massThe atomic mass (ma) is the mass of a specific isotope, most often expressed in unified atomic mass units.[1] The atomic mass is the total mass of protons, neutrons and electrons in a single atom(when the atom is motionless).[2]

The atomic mass is sometimes incorrectly used as a synonym of relative atomic mass, average atomic mass and atomic weight; these differ subtly from the atomic mass. The atomic mass is defined as the mass of an atom, which can only be one isotope at a time and is not an abundance-weighted average as in the case of atomic weight. In the case of many elements that have one dominant isotope the actual numerical similarity/difference between the atomic mass of the most common isotope and the relative atomic mass or standard atomic weights can be very small such that it does not affect most bulk calculations—but such an error can be critical when considering individual atoms. For elements with more than one common isotope the difference even to the most common atomic mass can be half a mass unit or more (e.g. chlorine). The atomic mass of an uncommon isotope can differ from the relative atomic mass or standard atomic weight by several mass units.

How Atomic Mass is Measured

Masses of atoms are measured by a mass spectrometer. Mass Spectrometry is the process of sending ionized vapors throughan electromagnetic field that seperates the ions relative to their mass-to-charge ratio. The ions are then detected and compared to the mass spectrum. Elements  seen on the Periodic Table are not simply the mass of that certain element but a weighted average of their Isotopes which will be discuessed later in this page.The atomic mass of an element is a weighted average of its isotopic masses.

Introduction Atomic mass

Introduction Atomic  mass:/The first scientists to measure atomic mass were John Dalton (between 1803 and 1805) Jons Jacoband Berzelius (between 1808 and 1826). Early atomic mass theory was purposed by the English chemist William Prout in a series of published papers in 1815 and 1816. Known was Prout's Law, Prout propsed that the known elements all had atomic weights that were whole number multiples of the atomic mass of hydrogen. Berzelius discovered that this was not always true by showing that Chlorine(Cl) had a mass of 35.45 which was not a whole number multiple of hydrogen's mass. The SI unit for atomic mass is the atomic mass unit (u), or Dalton (Da). The unit Dalton was derived from the carbon-12 isotope, as 12 u is the exact atomic mass of that isotope. So 1 u is 1/12 of the mass of a carbon-12 isotope (see Isotopic Atomic Mass section below). Later, in the 1860s, Cannizzario applied Avogadro's number to atomic masses, allowing stoichiometric conversions to be carried out. Over time, our knowledge of atomic mass has expanded greatly.