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Published byMitchell Melton Modified over 6 years ago

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**Isotopes Atoms of the same element that different mass numbers**

(which means they have different numbers of neutrons) They can be written two ways: element name - mass number Symbol mass # atomic # carbon - 12 6 p+ and 6 n0 C 12 6 C 13 6 carbon - 13 6 p+ and 7 n0 C 14 6 carbon - 14 6 p+ and 8 n0 Remember: mass number = p+ + n0 X 35 17 (a) X 35 18 (b) X 37 17 (c) X 38 19 (d) Which of the following are isotopes of the same element?

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**Atomic Mass x 75.77% = x 24.23% = 896.51 + = = 35 amu 2651.95 37 amu**

A weighted average of the masses of all the isotopes of an element. Step 1: Multiply the mass by its percent for each isotope Step 2: Add them together Step 3: Divide by 100 Use "amu" for its unit of measure Example: Chlorine has two naturally occurring isotopes, chlorine-35 and chlorine-37. Their relative abundances are 75.77% and 24.23%, respectively. Calculate the atomic mass for chlorine. 35 amu x 75.77% = 37 amu x 24.23% = 896.51 + = = 35.5 amu 100

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**Atomic Mass x 19.91% = x 80.09% = + = = 10.012 amu 199.34 11.009 amu**

Example: Element X has two natural isotopes. The isotope with a
mass of amu has a relative abundance of 19.91%. The
isotope with a mass of amu has a relative abundance of
80.09%. Calculate the atomic mass of this element. amu x 19.91% = 199.34 amu x 80.09% = 881.71 + = = 10.8 amu 100

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