Element relative atomic mass

Number of atoms of each element Relative atomic mass Total mass due to element %... 

Element Relative atomic mass IsHJkJmol b.HJkJg- ... 

By 1850. values of atomic weights (now called relative atomic masses) had been ascertained for many elements, and a knowledge of these enabled Newlands in 1864 to postulate a law of octaves. When the elements were arranged in order ol increasing atomic weight, each... 

Each element that has neither a stable isotope nor a characteristic natural isotopic composition is represented in this table by one of that element s commonly known radioisotopes identified by mass number and relative atomic mass. 

Titanium, Ti, atomic number 22, relative atomic mass 47.90, is the ninth most common element (ca 0.6% by weight) and is widely distributed in the earth s cmst. It is found particularly in the ores mtile, Ti02, and ilmenite, FeTiO. ... 

Accurate atomic weight values do not automatically follow from precise measurements of relative atomic masses, however, since the relative abundance of the various isotopes must also be determined. That this can be a limiting factor is readily seen from Table 1.3 the value for praseodymium (which has only 1 stable naturally occurring isotope) has two more significant figures than the value for the neighbouring element cerium which has 4 such isotopes. In the twelve years since the first edition of this book was published the atomic weight values of no fewer than 55 elements have been improved, sometimes spectacularly, e.g. Ni from 58.69( 1) to 58.6934(2). 

Atomic weights are known most accurately for elements which have only 1 stable isotope the relative atomic mass of this isotope can be determined to at least 1 ppm and there is no possibility of variability in nature. There are 20 such elements Be, F, Na, Al, P, Sc, Mn, Co, As, Y, Nb, Rh, I, Cs, Pr, Tb, Ho, Tm, Au and Bi. (Note that all of these elements except beryllium have odd atomic numbers — why )... 

The relative error is the absolute error divided by the true value it is usually expressed in terms of percentage or in parts per thousand. The true or absolute value of a quantity cannot be established experimentally, so that the observed result must be compared with the most probable value. With pure substances the quantity will ultimately depend upon the relative atomic mass of the constituent elements. Determinations of the relative atomic mass have been made with the utmost care, and the accuracy obtained usually far exceeds that attained in ordinary quantitative analysis the analyst must accordingly accept their reliability. With natural or industrial products, we must accept provisionally the results obtained by analysts of repute using carefully tested methods. If several analysts determine the same constituent in the same sample by different methods, the most probable value, which is usually the average, can be deduced from their results. In both cases, the establishment of the most probable value involves the application of statistical methods and the concept of precision. 

As the term mole refers to an amount of substance with reference to the specified mass of carbon-12, it is possible to express the relative molecular mass (the basis for the mole) for any substance as the additive sum of the relative atomic masses (R.A.M.s) of its component elements, for example ... 

One problem with Mendeleev s table was that some elements seemed to be out of place. For example, when argon was isolated, it did not seem to have the correct mass for its location. Its relative atomic mass of 40 is the same as that of calcium, but argon is an inert gas and calcium a reactive metal. Such anomalies led scientists to question the use of relative atomic mass as the basis for organizing the elements. When Henry Moseley examined x-ray spectra of the elements in the early twentieth century, he realized that he could infer the atomic number itself. It was soon discovered that elements fall into the uniformly repeating pattern of the periodic table if they are organized according to atomic number, rather than atomic mass. 

The relative atomic mass is the weighted average of the mass numbers of all the isotopes of a particular element. 

From the mass spectrum obtained, the relative atomic mass of an element may be calculated. 

One mole of atoms of an element is its relative atomic mass expressed in grams. 

To calculate the relative molecular mass of one mole of molecules of a substance, add up the relative atomic masses of the constituent elements. 

Table B.l summarizes the ground-state electron configuration and formal APH indices (turn number t, angular number l-n) for each known element, together with atomic number (Z) and relative atomic mass). As shown by the asterisks in the Anal column, 20 elements exhibit anomalous electron configurations (including two that are doubly anomalous - Pd and Th), compared with idealized t/l-n APH descriptors. These are particularly concentrated in the first d-block series, as well as among the early actinides. Such anomalies are indicative of configurational near-degeneracies that may require sophisticated multi-reference approximation methods for accurate description.

Table B.l. The currently known chemical elements, showing atomic number (Z), chemical symbol, name, relative atomic mass, ground-state electron configuration, and APH indices (t = turn number l-n = angular number) asterisks (, ) symbolize anomalous (APH non-conforming) ground-state electronic configurations, which are indicative of configurational near-degeneracy...

Mass spectrometry is based upon the separation of charged ionic species by their mass-to-charge ratio, m/z. Within the general chemical context however, we are not used to taking into concern the isotopes of the elemental species involved in a reaction. The molecular mass of tribromomethane, CHBrs, would therefore be calculated to 252.73 g mol using the relative atomic masses of the elements as listed in most periodic tables. In mass spectrometry we have to leave this custom behind. Because the mass spectrometer does not separate by elements but by isotopic mass, there is no signal at m/z 252.73 in the mass spectmm of tribromomethane. Instead, major peaks are present at m/z 250, 252, 254 and 256 accompanied by some minor others. 

The relative atomic mass or the atomic weight as it is also often imprecisely termed is calculated as the weighted average of the naturally occurring isotopes of an element. [3]The weighted average is calculated from... 

The relative molecular mass, M, or molecular weight is calculated from the relative atomic masses of the elements contributing to the empirical formula. [3]... 

Note The calculation of relative molecular mass, Mr, of organic molecules exceeding 2000 u is significantly influenced by the basis it is performed on. Both the atomic weights of the constituent elements and the natural variations in isotopic abundance contribute to the differences between monoisotopic- and relative atomic mass-based values. In addition, they tend to characteristically differ between major classes of biomolecules. This is primarily because of molar carbon content, e.g., the difference between polypeptides and nucleic acids is about 4 u at Mr = 25,000 u. Considering terrestrial sources alone, variations in the isotopic abundance of carbon lead to differences of about 10-25 ppm in Mr which is significant with respect to mass measurement accuracy in the region up to several 10 u. [41]... 

Table A. 1 comprises the stable elements from hydrogen to bismuth with the radioactive elements technetium and promethium omitted. Natural variations in isotopic composition of some elements such as carbon or lead do not allow for more accurate values, a fact also reflected in the accuracy of their relative atomic mass. However, exact masses of the isotopes are not affected by varying abundances. The isotopic masses listed may differ up to some 10 u in other publications.

Table A.l. Isotopic mass, isotopic composition, and relative atomic mass [u] of non-radioactive elements. lUPAC 2001.

Table8.6. Relative atomic mass Mr of naturally occurring elements (basis = 12.00)...

Precise measurements of the amounts of different isotopes can be important. You need to know the exact measurements if you re asked to figure out an element s atomic mass. To calculate an atomic mass, you need to know the masses of the isotopes and the percentage of the element that occurs as each isotope (this is called the relative abundance ). To calculate an average atomic mass, make a list of each isotope along with its mass and its percent relative abundance. Multiply the mass of each isotope by its relative abundance. Add the products. The resulting sum is the atomic mass. 

The relative atomic mass of an element is the weighted mean of the isotopic masses. The weighted mean is calculated from the masses of all the possible isotopes of the element, taking into account the natural relative abundance of each isotope. 

Mendeleev s chemical insight led him to leave gaps for elements that would be needed to complete the pattern but were unknown at the time. When they were discovered later, he turned out to be strikingly correct. For example, his pattern required an element he named eka-silicon below silicon and between gallium and arsenic. He predicted that the element would have a relative atomic mass of 72 (taking the mass of hydrogen as 1) and properties similar to... 

One problem with Mendeleev s table was that some elements seemed to be out of place. For example, when argon was isolated, it did not seem to have the correct mass for its location. Its relative atomic mass of 40 is the same as calcium s, but argon is an inert gas and calcium a reactive metal. Such anom-... 

The atomic mass unit (amu) is defined as 1/12 the mass of a carbon-12 isotope. The relative atomic mass of an element is the weighted average of the isotopes relative to I/12(of jhe carbon-12 isotope. For example, the atomic mass of neon is 20.17 amu and is calculated from the following data neon-19 (amu of 19.99245, natural abundance of 90.92%), neon-20 (amu of 20.99396, natural abu dan c of 0,260%) and ncon-21 (amu of 21.99139, natural abundance ofc 82%) ... 

The average mass of a large number of atoms of an element is called its relative atomic mass (symbol AT). This quantity takes into account the percentage abundance of all the isotopes of an element which exist. 

This small quantity is not easy to work with so, as you saw in Chapter 3, a scale called the relative atomic mass scale is used. In this scale an atom of carbon is given a relative atomic mass, An of 12.00. All other atoms of the other elements are given a relative atomic mass compared to that of carbon. 

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