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Isotopes and Average Atomic Mass

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)

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