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How Atoms Differ Chapter 4 Section 4.3.

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Presentation on theme: "How Atoms Differ Chapter 4 Section 4.3."— Presentation transcript:

1 How Atoms Differ Chapter 4 Section 4.3

2 Objectives Calculate the number of protons, neutrons, and electrons in an atom given its mass number and atomic number. Define an isotope and explain why atomic masses are not whole numbers. To calculate how we get atomic masses of elements with isotopes.

3 Determining the Number of Protons in an Element:
Find the element on the periodic table. Find the Atomic Number of the element. Atomic Number = # of protons in the element. *Remember the # of protons of an atom determines what element it is representing.

4 Examples How many protons do the following elements have? Oxygen Zinc
8 Zinc 30 Bismuth 83

5 Determining the number of Neutrons in an atom:
Atomic Mass = protons + neutrons If we know the # of protons from the Atomic Number, then … Atomic Mass - # of protons = # of neutrons *Remember you should take the atomic mass on the periodic table and round it first

6 *There are 8 neutrons in a neutral oxygen atom.
Example Oxygen Atomic Mass from the Periodic Table: > 16 Subtract the # of protons (Atomic Number) from the Atomic Mass 16 – 8 = 8 *There are 8 neutrons in a neutral oxygen atom.

7 Electrons In a neutral atom, there must be the same number of positive and negative charges. Therefore… # of protons = # of electron The # of electrons = Atomic Number.

8 What are Isotopes? Atoms that have the same number of protons but different numbers of neutrons. They will have the same Atomic Number but different Atomic Masses.

9 Isotopes In nature most elements are found as a mixture of isotopes.
No matter where a sample of an element is obtained, the relative abundance of each isotope is the constant. Example: In a banana there is 93.25% potassium with 20 neutrons, % have 22 neutrons, and % have 21 neutrons.

10 Ways to write Isotopes 39 Potassuim-39 or K-39 or 19K protons-19
electrons-19 neutrons-20

11 Masses of subatomic Particles
Mass of Atoms *Both protons and neutrons are very close in mass to 1 amu. *Atomic Mass Unit = 1/12 the mass of carbon-12 Masses of subatomic Particles Particle Mass (amu Electron Proton Neutron

12 How do isotopes affect the mass?
We take weighted averages of those elements based on the abundance of the isotopes. Weighted Average = mass of isotope x percent abundance (decimal form) Then add the values that each isotope gives.

13 Example Element X has 2 natural isotopes. The isotopes with mass amu has a relative abundance of 19.91%. The isotope with mass amu has a relative abundance of 80.09%. What is this element? Calculate the average atomic mass. Compare this mass to the Periodic Table to find your element.

14 Isotope A: mass = 10. 012 amu @ 19. 91% Isotope B: mass = 11
Isotope A: mass = % Isotope B: mass = % A: x = B: x = Now Add them: = 10.81amu Average Atomic Mass for Element X = amu

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