Atoms average atomic masses

Use the following terms to create a concept map atoms, average atomic mass, molecules, mole, percentage composition, and molar masses. 

The atomic mass is the average mass of an element in atomic mass units ("amu"). 

Though individual atoms always have an integer number of amus, the atomic mass on the periodic table is stated as a decimal number because it is an average of the various isotopes of an element. Isotopes can have a weight either more or less than the average. The average number of neutrons for an element can be found by subtracting the number of protons (atomic number) from the atomic mass. 

A common mistake for beginners in mass spectrometry is to confuse average atomic mass and isotopic mass. For example, the average atomic mass for chlorine is close to 35.45, but this average is of the numbers and masses of Cl and Cl isotopes. This average must be used for instruments that cannot differentiate isotopes (for example, gravimetric balances). Mass spectrometers do differentiate isotopes by mass, so it is important in mass spectrometry that isotopic masses be used... 

Atomic mass (Section 1.1) The weighted average mass of an element s naturally occurring isotopes. 

Relative masses of atoms of different elements are expressed in terms of their atomic masses (often referred to as atomic weights). The atomic mass of an element indicates how heavy, on the average, one atom of that element is compared with an atom of another element... 

Notice that hydrogen has an atomic mass of 1.008 amu helium has an atomic mass of 4.003 amu. This means that, on the average, a helium atom has a mass that is about one-third that of a C-12 atom ... 

The average atomic mass shown in the periodic table is not equal to the mass number. 

Click Coached Problems for a self-study module on average atomic mass from isotopic abundance. 

Bromine is a red-orange liquid with an average atomic mass of 79.90 amu. Its name is derived from the Greek word bromos (fipofios), which means stench. It has two naturally occurring isotopes Br-79 (78.92 amu) and Br-81 (80.92 amu). What is the abundance of the heavier isotope ... 

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. 

The greater the mass of an individual atom, the greater the molar mass of the substance. However, most elements exist in nature as a mixture of isotopes. We saw in Section B, for instance, that neon exists as three isotopes, each with a different mass. In chemistry, we almost always deal with natural samples of elements, which have the natural abundance of isotopes. So, we need the average molar mass, the molar mass calculated by taking into account the masses of the isotopes and their relative abundances in typical samples ... 

STRATEGY First calculate the average atomic mass of the isotopes by adding together the individual masses, each multiplied by the fraction that represents its abundance. Then obtain the molar mass, the mass per mole of atoms, by multiplying the average atomic mass by Avogadro s constant. 

E.29 The king of Zirconia is naturally fond of the element zirconium and has established an independent definition of the mole based on zirconium. The mass of one zirconium-90 atom is 1.4929 X 10 22 g. If zirconium were the standard used for molar mass (instead of carbon-12), 1 mol would be defined as the amount of substance that contains the same number of entities as there are atoms in exactly 90 g of zirconium-90. In that case, what would be (a) the molar mass of carbon-12 (b) the (average) molar mass of gold ... 

Two forms of the same element are called isotopes. The isotopes of an element have the same atomic number but have different atomic masses. Iron has several isotopes that, when weighted by their naturally occurring abundance, gives an average mass of 55.845 amu. A simple example would be an element with only two isotopes, one with a mass of 10 amu, the other of 12 amu. If the isotopes were equally common, then the average atomic mass for that element would be 11. If 90% of the element occurred naturally as the isotope with a mass of 10 amu, then the average atomic mass would be 10.2, as calculated below ... 

The atomic mass of an element is the weighted average mass of the isotopes of that element. Based on this definition, which of these does NOT show the correct atomic mass for an element ... 

D The average atomic mass of one atom of each element in the molecule... 

Table 6.6 Some frequently encountered atoms with their monoisotopic and average atomic masses...

Most mass spectrometers will resolve ions with unit resolution up to at least 2000 Da, and so monoisotopic atomic masses are used in these cases. Above 2000 Da, the resolution should be checked and, if it is insufficient to resolve adjacent isotopes, then average atomic masses can be used in calculations. 

Ans. (a) The nucleus is a distinct part of the atom. Neutrons are subatomic particles which, along with protons, are located in the nucleus. (b) Mass number refers to individual isotopes. It is the sum of the numbers of protons and neutrons. Atomic weight refers to the naturally occurring mixture of isotopes, and is the relative mass of the average atom compared to l2C. (/) Atomic mass is the same as atomic weight [see (b)]. Atomic mass unit is the unit of atomic weight. 

Naturally occurring boron consists of approximately 20% of 10B and 80% of UB, leading to an average atomic mass of 10.8 amu. Because 10B has a relatively large cross-section for absorption of slow (thermal) neutrons, it is used in control rods in nuclear reactors and in protective shields. In order to obtain a material that can be fabricated into appropriate shapes, boron carbide is combined with aluminum. 

The atomic mass of chlorine is reported as 35.5 to three significant figures. No single atom of chlorine has that mass because the atomic mass of any element is the weighted average of all the isotopes, not the mass of any one atom. Chlorine is 76% 35C1 and 24% 37C1. 

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

A The average atomic mass of boron is 10.811, which is closer to 11.009305 than to 10.012937. Thus, boron-11 is the isotope present in greater abundance. 

The weighted-average atomic mass of the element iridium is just slightly more than 192 u. The mass of the first isotope is a bit less than 191 u. Hence, the mass of the second isotope must more than 192 u that isotope must be 193 Ir. 

Each of the isotopic masses is multiplied by its fractional abundance. The resulting products are summed to obtain the average atomic mass. 

To determine the average atomic mass, we use the following expression average atomic mass = (isotopic mass x fractional natural abundance)... 

We use the expression for determining the weighted-average atomic mass. 

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