## Monday, 6 June 2011

### Periodic table:Atomic mass

PERIODIC TABLE : ATOMIC MASS
123456789101112131415161718
1
He
1.007944.002602
2LiBe
 H
BCNOFNe
6.9419.01218210.81112.010714.0067415.999418.998403220.1797
3NaMgAlSiPSClAr
22.98977024.305026.98153828.085530.97376132.06635.452739.948
4KCaScTiVCrMnFeCoNiCuZnGaGeAsSeBrKr
39.098340.07844.95591047.86750.941551.996154.93804955.84558.93320058.693463.54665.3969.72372.6174.9216078.9679.90483.80
5RbSrYZrNbMoTcRuRhPdAgCdInSnSbTeIXe
85.467887.6288.9058591.22492.9063895.94(98)101.07102.90550106.42107.8682112.411114.818118.710121.760127.60126.90447131.29
6CsBa*HfTaWReOsIrPtAuHgTlPbBiPoAtRn
132.90545137.327178.49180.9479183.84186.207190.23192.217195.078196.96655200.59204.3833207.2208.98038(209)(210)(222)
7FrRa**RfDbSgBhHsMtUunUuuUub
(223)(226)(261)(262)(263)(262)(265)(266)(269)(272)(277)
*LaCePrNdPmSmEuGdTbDyHoErTm

151.964157.25158.92534162.50164.93032167.26168.93421173.04174.967
**AcThPaUNpPuAmCm
 Yb Lu 138.9055 140.116 140.90765 144.24 (145) 150.36
Bk
CfEsFmMdNoLr
(227)232.0381231.03588238.0289(237)(244)(243)(247)(247)(251)(252)(257)(258)(259)(262)
Element Groups (Families)
Alkali EarthAlkaline EarthTransition Metals
Rare EarthOther MetalsMetalloids
Non-MetalsHalogensNoble Ges
as

## HistoryMain article: HistoryThe first scientists to determine atomic weights were John Dalton between 1803 and 1805 and Jöns Jakob Berzelius between 1808 and 1826. Atomic weight was originally defined relative to that of the lightest element hydrogen taken as 1.00, and in the 1820s Prout's hypothesisstated that atomic masses of all elements would prove via a whole number rule to be exact multiples of this hydrogen weight. Berzelius, however, soon proved that this hypothesis did not always hold even approximately, and in some elements, such as chlorine, atomic weight falls almost exactly between two multiples of the hydrogen weight. Still later, as noted, this was shown to be an isotope effect, and that the atomic masses of pure isotopes, or nuclides, are multiples of the hydrogen mass, to within about 1%.In the 1860s Stanislao Cannizzaro refined atomic weights by applying Avogadro's law (notably at the Karlsruhe Congress of 1860). He formulated a law to determine atomic weights of elements: the different quantities of the same element contained in different molecules are all whole multiples of the atomic weight and determined atomic weights and molecular weights by comparing the vapor density of a collection of gases with molecules containing one or more of the chemical element in question.[6]In the early twentieth century, up until the 1960s chemists and physicists used two different atomic mass scales. The chemists used a scale such that the natural mixture of oxygen isotopes had an atomic mass 16, while the physicists assigned the same number 16 to the atomic mass of the most common oxygen isotope (containing eight protons and eight neutrons). However, because oxygen-17 and oxygen-18 are also present in natural oxygen this led to 2 different tables of atomic mass. The unified scale based on carbon-12, 12C, met the physicists' need to base the scale on a pure isotope, while being numerically close to the chemists' scale. of chemistry

### 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.

# 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.