Periodic table:Atomic mass

PERIODIC TABLE : ATOMIC MASS

1.00794

4.002602

6.941

9.012182

10.811

12.0107

14.00674

15.9994

18.9984032

20.1797

22.989770

24.3050

26.981538

28.0855

30.973761

32.066

35.4527

39.948

39.0983

40.078

44.955910

47.867

50.9415

51.9961

54.938049

55.845

58.933200

58.6934

63.546

65.39

69.723

72.61

74.92160

78.96

79.904

83.80

85.4678

87.62

88.90585

91.224

92.90638

95.94

(98)

101.07

102.90550

106.42

107.8682

112.411

114.818

118.710

121.760

127.60

126.90447

131.29

132.90545

137.327

178.49

180.9479

183.84

186.207

190.23

192.217

195.078

196.96655

200.59

204.3833

207.2

208.98038

(209)

(210)

(222)

Uun

Uuu

Uub

(223)

(226)

(261)

(262)

(263)

(262)

(265)

(266)

(269)

(272)

(277)

151.964

157.25

158.92534

162.50

164.93032

167.26

168.93421

173.04

174.967

Yb	Lu
	138.9055	140.116	140.90765	144.24	(145)	150.36

(227)

232.0381

231.03588

238.0289

(237)

(244)

(243)

(247)

(251)

(252)

(257)

(258)

(259)

(262)

Element Groups (Families)
Alkali Earth	Alkaline Earth	Transition Metals
Rare Earth	Other Metals	Metalloids
Non-Metals	Halogens	Noble Ges

History
Main article: History
The 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

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

M u

is the Molar mass constant and

N A

is the Avogadro constant.

Atomic mass

Atomic massThe atomic mass (m_a) 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.

Monday, 6 June 2011

Periodic table:Atomic mass

History

History

Conversion factor between atomic mass units and grams

Conversion factor between atomic mass units and grams

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]

How Atomic Mass is Measured

Introduction Atomic mass

Atomic massThe atomic mass (m_a) 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]