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Lesson 2.03. Isotopes and Weighted Averages. Introduction. How can more than one form of atom make up an element? What can be different and what is always the same among the different atoms of a given element?. Objectives. After completing this lesson, you will be able to : - PowerPoint PPT Presentation

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Lesson 2.03Isotopes and Weighted AveragesIntroductionHow can more than one form of atom make up an element?

What can be different and what is always the same among the different atoms of a given element?

ObjectivesAfter completing this lesson, you will be able to:

Identify the structural differences between the isotopes of elements.

Determine the weighted average mass of an element when given the appropriate data.What do you know?The number of protons in an atoms nucleus identifies the atom as a specific element. Protons never change for a particular element

EXAMPLE: All oxygen atoms have eight protons. We know this because Oxygens atomic number = 8

What is an Isotope?Isotopes are atoms of the same element that have different masses due to their varying numbers of neutrons

Ways to represent an IsotopeMethod One: Add the mass number of the isotope, with a hyphen, to the end of the elements name.

EXAMPLES: an isotope of a carbon (C) atom that has six protons, six neutrons, and six electrons you can represent the isotope as Carbon-12. The 12 is the mass number of the isotope (six protons + six neutrons).Uranium-235would be an isotope of uranium with a mass of 235 (92 protons + 143 neutrons)Ways to represent an IsotopeMethod Two:The mass number is given as a superscript to the upper left side of the elements symbol, with the atomic number as a subscript to the lower left side of the symbol.

EXAMPLE: If you have an isotope of a carbon (C) atom that has six protons, eight neutrons, and six electrons you can represent the isotope by writing C with 14 a superscript on the upper left side and six as a subscript in the lower left side of the symbol. The 14 is the mass number and the six is the atomic number.C146Mass NumberAtomic NumberMethod Two ExamplesEXAMPLE:

An isotope of hydrogen (H) that has a mass of two would be written as H with two as a superscript on the upper left side and one as a subscript in the lower left side of the symbol. The two is the mass number and the one is the atomic number.H12PracticeRefer to the periodic table and each elements atomic number to answer the following questions:1- How many protons, neutrons, and electrons are in an atom of chlorine-37? Chlorine has an atomic number of 17.

2-How many protons, neutrons, and electrons are in an atom of Br with an atomic number of 35 and a mass number of 80?

More Practice3-How many protons, neutrons, and electrons in an atom of beryllium-9? Beryllium has an atomic number of 4.

4- How many protons, neutrons, and electrons are in an atom of Ar with an atomic number of 18 and a mass number of 38?

Check your Answers1- How many protons, neutrons, and electrons are in an atom of chlorine-37? Chlorine has an atomic number of 17. Protons = 17 (atomic number)Electron = 17 (atomic number) Neutrons = 20 (neutrons = mass number (37)atomic number (17)

2- How many protons, neutrons, and electrons are in an atom of Br with an atomic number of 35 and a mass number of 80? Protons = 35 (atomic number)Electron = 35 (atomic number) Neutrons = 45 (neutrons = mass number (80)atomic number (35)Check Your Answers3- How many protons, neutrons, and electrons in an atom of beryllium-9? Beryllium has an atomic number of 4. Protons = 4 (atomic number)Electron = 4 (atomic number) Neutrons = 5 (neutrons = mass number (9)atomic number (4)

4- How many protons, neutrons, and electrons are in an atom of Ar with an atomic number of 18 and a mass number of 38? Protons = 18 (atomic number)Electron = 18 (atomic number) Neutrons = 20 (neutrons = mass number (38)atomic number (18)Atomic MassHave you noticed that the atomic masses on the periodic table have decimal values even though the mass numbers are whole numbers?

Average Atomic MassThe average mass of an element calculated by averaging the mass values for all of that elements isotopes.

Because pure elements are made up of mixtures of isotopes, it is important to take into account the percentage and mass of each isotope, especially when some isotopes make up a greater percentage of the mixture than other isotopes.

the average atomic mass of each element is calculated using a weighted average.What is a weighted average?A weighted average accounts for the percent abundance and mass of each isotope in an element.Percent abundance is a measure of how common or rare that isotope is.The average atomic mass of an element can be calculated by multiplying the mass of each isotope by its relative abundance (the percentage represented in decimal form) and adding the resultsIsotope A [percent abundance mass] + Isotope B [percent abundance mass] + continue for each isotope = average atomic massExampleIsotopeMassPercent Abundance Oxygen- 1616.0 u99.72 %Oxygen- 1717.0 u0.038%Oxygen- 1818.0 u0.200%Step 1: Convert the percentages to decimals by dividing by 100IsotopeMassPercent Abundance (decimals)Oxygen- 1616.0 u99.72 /100= .99762Oxygen- 1717.0 u0.038/100= .00038Oxygen- 1818.0 u0.200/100= .00200Example continuedStep 2- multiply each isotopes mass by the decimal of the abundanceIsotopeMassDecimal of the AbundanceTotalOxygen- 1616.0 u.9976215.96192Oxygen- 1717.0 u.000380.00646Oxygen- 1818.0 u.002000.036Step 3- add the individual totals together(15.96192 + 0.00646 + 0.036) = 16.0uPractice 1Copper has two naturally occurring isotopes. Copper-63, with an atomic mass of 62.94 u, makes up 69.17 percent of the sample, and Copper-65, with an atomic mass of 64.93 u, makes up the other 30.83 percent. Calculate the weighted average atomic mass of copper based on the given information. Please give your answer in a total of four significant figures.IsotopeMassPercent Abundance Copper-6362.94 u69.17%Copper-6564.93 u30.83%Practice 2Three isotopes of argon (Ar) occur in nature. Argon-36 (35.968 u) makes up 0.337 percent of the sample, Argon-38 (37.963 u) makes up 0.063 percent of the sample, and Argon-40 (39.962 u) makes up the other 99.6 percent of the sample. Calculate the weighted average atomic mass of argon, giving your answer in a total of five significant figures.IsotopeMassPercent Abundance Argon-3635.968 u0.337 %Argon-3837.963 u0.063%Argon-4039.963 u99.6%Answers to PracticeQuestion 1-[.6917 62.94 u] + [.3083 64.93 u] =

63.55 u (average atomic mass)

Question 2-[0.00337 35.968 u] + [0.00063 37.963 u] + [0.99600 39.962 u] =

39.947 u (average atomic mass)ReviewTermDefinitionExampleIsotopeatoms of the same element that have different masses due to their varying numbers of neutrons

Carbon-12, Carbon-14Percent AbundanceA calculated percentage based on how common a specific isotope is in any given sample of an element Copper-63, with an atomic mass of 62.94 u, makes up 69.17% of the sample and Copper-65, with an atomic mass of 64.93 u, makes up the other 30.83%Average Atomic Massaverage mass of an element calculated by averaging the weighted mass values for all of that elements isotopes[.6917 62.94 u] + [.3083 64.93 u] = 63.55 u (average atomic mass)Congrats!You should now be able to:Identify the structural differences between the isotopes of elements.Determine the weighted average mass of an element when given the appropriate data.