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Isotopes and Average Atomic Mass

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jeneil
15 diapositives · 455 Views · Oct 07, 2014

Isotopes and Average Atomic Mass. Objectives. Explain what an isotope is. Compare and contrast two different isotopes Calculate the average atomic mass of an element. Review: How to read symbols. ion. When you change the number of electrons , you get an __________________________

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Isotopes and Average Atomic Mass
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  1. Isotopes andAverage Atomic Mass

  2. Objectives • Explain what an isotope is. • Compare and contrast two different isotopes • Calculate the average atomic mass of an element.

  3. Review: How to read symbols ion • When you change the number of electrons, you get an __________________________ • When you change the number of protons, you get an _____________________________ • When you change the number of neutrons, you get ____________________________ • Symbols contain the mass number and the atomic number. completely new element Isotopes of the same element U Mass Number → Can change! 238 Atomic Number → NEVER Changes 92

  4. Isotopes • Atoms of the same element can have different numbers of neutrons. • Atoms with the same number of protons, but different mass numbers are called isotopes.

  5. Naming & Writing Isotopes • There are two ways we can write isotopes. Isotopes of Carbon include: 14C and 12C • We can also put the mass number after the name of the element: • carbon-12 • carbon-14 • uranium-235

  6. Isotopesare atoms of the same element having different masses, due to varying numbers of neutrons. 0 1 2

  7. Elements occur in nature as mixtures of isotopes.

  8. Check for Understanding • Are isotopes? No. Isotopes must be the same element. • Are all isotopes man-made? No. Isotopes occur in nature. Right now, every living thing has in them. • Are all isotopes radioactive? No. Both Carbon -12 and Carbon -14 are isotopes. Only Carbon-14 is unstable. We will learn how to predict when an isotope is radioactive or not, later.

  9. Why Average Atomic Mass? • The majority of the masses listed on the periodic table are decimals. Why? • Because natural samples of elements are a mixture of naturally occurring isotopes. • Ex: How heavy is an atom of cesium? • It depends, because there are different kinds of cesium atoms. Most have a mass of 133, but some have a mass of 132 and 134. • To account for the mixture of isotopes, we report the masses of elements as the average atomic mass. • This is based on the abundance (percentage) of each isotope of that element found in nature.

  10. How do we measure Atomic Mass? • We use grams to measure the mass of most things in chemistry, but not for atomic mass. Why? • Because the masses would be too small if measured in grams. • Instead of grams, the unit we use is the Atomic Mass Unit (amu) • It is defined as one-twelfth the mass of a carbon-12 atom. • Don’t worry about why we use it, just memorize this as a fact! It is like 1 gallon = 4 quarts or why a dozen = 12. It just is. • Carbon-12 chosen because of its isotope purity.

  11. Atomic Masses Atomic mass is the average of all the naturally occurring isotopes of that element. 6 protons 6 neutrons 6 protons 7 neutrons 6 protons 8 neutrons Carbon = 12.011 ** This rounds to the major isotope

  12. To calculate the Average Atomic Mass: • Convert the percentages into decimals (divide by 100) • Multiply the percentage (in decimal form) by the mass of the isotope • Add the masses from step 2 Example:A sample of cesium is 75% Cesium-133, 20% Cesium-132 and 5% Cesium-134. What is its average atomic mass?

  13. Example 1: A sample of cesium is 75% 133Cs, 20% 132Cs and 5% 134Cs. What is its average atomic mass? • What are the three isotopes in this problem? 1. Convert percents to decimals (divide by 100) 2. Multiply the percent (in decimal form) by the mass (0.75) x (133) = 99.75 (0.20) x (132) = 26.4 (0.05) x (134) = 6.7 Total 132.85 amu 3. Add the masses together to get the avg atomic mass.

  14. Example 2: Boron has two naturally occuring isotopes, 19.8% Boron-10 and 80.2% Boron-11. What is its average atomic mass? • What are the two isotopes in this problem? 1. Convert percents to decimals (divide by 100) 2. Multiply the percent (in decimal form) by the mass (0.198) x (10) = 1.98 (0.802) x (11) = 8.82 Total 10.8 amu 3. Add the masses together to get the avg atomic mass.

  15. Atomic Mass vs. Mass # • Mass Number = Total number of particles in the nucleus (always a whole number!) • Atomic Mass – weighted average of all the isotopes of an element (a decimal number)