Isotopes and Average Atomic Mass

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Isotopes andAverage Atomic Mass
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Objectives

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

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Review How to read symbols

• 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.

ion
completely new element
Isotopes of the same element
U
Mass Number ? Can change!
238
Atomic Number ? NEVER Changes
92
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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.

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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

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Isotopes are atoms of the same element having
different masses, due to varying numbers of
neutrons.
Isotope Protons Electrons Neutrons Nucleus
Hydrogen1 (protium) 1 1
Hydrogen-2 (deuterium) 1 1
Hydrogen-3 (tritium) 1 1
0
1
2
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Elements occur in nature as mixtures of
isotopes.
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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.

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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.

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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.
• Dont 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.

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Atomic Masses
Atomic mass is the average of all the naturally
occurring isotopes of that element.
Isotope Symbol Composition of the nucleus in nature
Carbon-12 12C 98.89
Carbon-13 13C 1.11
Carbon-14 14C lt0.01
6 protons 6 neutrons
6 protons 7 neutrons
6 protons 8 neutrons
Carbon 12.011
This rounds to the major isotope
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To calculate the Average Atomic Mass

1. Convert the percentages into decimals (divide by
   100)
2. Multiply the percentage (in decimal form) by the
   mass of the isotope
3. 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?
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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.

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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.

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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)