{"id":159,"date":"2018-03-19T15:50:20","date_gmt":"2018-03-19T15:50:20","guid":{"rendered":"https:\/\/courses.lumenlearning.com\/suny-orgbiochemistry\/chapter\/atomic-masses\/"},"modified":"2018-08-06T16:10:46","modified_gmt":"2018-08-06T16:10:46","slug":"atomic-masses","status":"publish","type":"chapter","link":"https:\/\/courses.lumenlearning.com\/suny-monroecc-orgbiochemistry\/chapter\/atomic-masses\/","title":{"raw":"2.5 Atomic Masses","rendered":"2.5 Atomic Masses"},"content":{"raw":"
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Learning Objective<\/h3>\r\n
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  1. Define atomic mass and atomic mass unit.<\/li>\r\n<\/ol>\r\n<\/div>\r\n<\/div>\r\n<\/div>\r\n<\/div>\r\n<\/div>\r\n<\/div>\r\n
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    \r\n\r\nEven though atoms are very tiny pieces of matter, they have mass. Their masses are so small, however, that chemists often use a unit other than grams to express them\u2014the atomic mass unit.\r\n

    The atomic mass unit<\/strong>:<\/span><\/span>\u00a0(abbreviated u, although amu is also used) is defined as [latex]\\frac{1}{12}[\/latex] of the mass of a C-12 atom:<\/p>\r\n

    1\u00a0u=[latex]\\frac{1}{12}[\/latex] the\u00a0mass\u00a0of\u00a0C-12 atom <\/span><\/p>\r\n

    It is equal to 1.661 \u00d7 10\u221224<\/sup> g.<\/p>\r\n

    Masses of other atoms are expressed with respect to the atomic mass unit. For example, the mass of an atom of H-1 is 1.008 u, the mass of an atom of O-16 is 15.995 u, and the mass of an atom of S-32 is 31.97 u. Note, however, that these masses are for particular isotopes of each element. Because most elements exist in nature as a mixture of isotopes, any sample of an element will actually be a mixture of atoms having slightly different masses (because neutrons have a significant effect on an atom\u2019s mass). How, then, do we describe the mass of a given element? By calculating an average of an element\u2019s atomic masses, weighted by the natural abundance of each isotope, we obtain a weighted average mass called the atomic mass<\/span><\/span><\/strong>\u00a0(also commonly referred to as the atomic weight<\/em>) of an element.<\/p>\r\n

    For example, boron exists as a mixture that is 19.9% B-10 and 80.1% B-11. The atomic mass of boron would be calculated as (0.199 \u00d7 10.0 u) + (0.801 \u00d7 11.0 u) = 10.8 u. Similar average atomic masses can be calculated for other elements. Carbon exists on Earth as about 99% C-12 and about 1% C-13, so the weighted average mass of carbon atoms is 12.01 u.<\/p>\r\n

    The table in Chapter 21 \"Appendix: Periodic Table of the Elements\"<\/a> also lists the atomic masses of the elements.<\/p>\r\n\r\n

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    Example 6<\/h3>\r\n

    What is the average mass of a carbon atom in grams?<\/p>\r\n

    [reveal-answer q=\"9999\"]Show Answer[\/reveal-answer]\r\n[hidden-answer a=\"9999\"]This is a simple one-step conversion, similar to conversions we did in Chapter 1 \"Chemistry, Matter, and Measurement\". We use the fact that 1 u = 1.661 \u00d7 10\u221224 g:<\/p>\r\n

    [latex]12.01\\cancel{u}\u2009\\times\u2009\\frac{1.661\\times 10^{\u221224} g}{1 \\cancel{u}}\u2009=1.995\\times 10^{\u221223} g [\/latex] This is an extremely small mass, which illustrates just how small individual atoms are.[\/hidden-answer]<\/p>\r\n\r\n<\/div>\r\n

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    Skill-Building Exercise<\/h3>\r\n
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      What is the average mass of a tin atom in grams? The atomic mass of tin is 118.71 u.<\/p>\r\n\r\n<\/div><\/li>\r\n<\/ol>\r\n<\/div>\r\n<\/div>\r\n

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      Concept Review Exercises<\/h3>\r\n
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        Define atomic mass. Why is it considered a weighted average?<\/p>\r\n\r\n<\/div><\/li>\r\n \t

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        What is an atomic mass unit?<\/p>\r\n\r\n<\/div><\/li>\r\n<\/ol>\r\n<\/div>\r\n

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          The atomic mass is an average of an element\u2019s atomic masses, weighted by the natural abundance of each isotope of that element. It is a weighted average because different isotopes have different masses.<\/p>\r\n\r\n<\/div><\/li>\r\n \t

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          An atomic mass unit is 1\/12th of the mass of a C-12 atom.<\/p>\r\n\r\n<\/div><\/li>\r\n<\/ol>\r\n<\/div>\r\n<\/div>\r\n

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          Key Takeaway<\/h3>\r\n<\/div>\r\n
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