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Development of the Atom

Development of the Atom. Essential Question. How were atoms and their subatomic particles discovered?. How small is the atom?. https://www.youtube.com/watch?v=yQP4UJhNn0I What makes up the nucleus? What surrounds the nucleus? What contributes most of the mass of the atom?

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Development of the Atom

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  1. Development of the Atom

  2. Essential Question • How were atoms and their subatomic particles discovered?

  3. How small is the atom? • https://www.youtube.com/watch?v=yQP4UJhNn0I • What makes up the nucleus? • What surrounds the nucleus? • What contributes most of the mass of the atom? • What is the atom mostly made up of (volume)?

  4. Democritus—400B.C. • Greek philosopher • World is made of two things • Empty space • Tiny particles, “atoma” • Believed shape determined properties • An idea only • Was NOT backed by experimental evidence

  5. Aristotle—350B.C. • Greek Philosopher • Proposed that matter is continuous • NOT made of smaller particles • Accepted until the 17th century

  6. John Dalton—Early 1800s • Dalton’s Atomic Theory • All matter is made of atoms, which are indivisible • All atoms of an element are identical • Compounds are formed from two or more elements

  7. Dalton’s Atom • Believed to be an indivisible sphere

  8. J.J. Thomson—1897 • Cathode Ray Tube Experiment • Discovered the existence of the electron

  9. Thomson’s Atom • Sphere with electrons embedded in it • Known as “Plum Pudding Model”

  10. Robert Millikan—1909 • Oil Drop Experiment • Charge of an electron = -1.6x10-19Coulombs • Standard Unit = -1

  11. Ernest Rutherford—1911 • Gold Foil Experiment • discovered the presence of a positive nucleus

  12. Rutherford Atom • Small dense nucleus, with electrons orbiting • Mostly empty space

  13. What are the parts of the atom? • Work with your partner to fill in what you know on the chart below:

  14. What are the subatomic particles?

  15. A really corny joke • A neutron walked into a restaurant, ordered a drink, and asked “How much?” • The waiter replied “For you, no charge.” • Take out your periodic tables!

  16. Atomic Number (Z) • Number of protons • Defines an element • Same # of protons = same element • Different # of protons = different elements • For neutral atoms = # e- • This number is always found on the periodic table!

  17. Essential Question • Why do the masses of individual atoms differ?

  18. Why do the masses of the same element differ? • Isotopes:atoms of the same element that differ in mass due to different numbers of neutrons • Same # protons (same element), different # neutrons (different isotope)

  19. Check Point • Element A has 6 protons, 6 neutrons, and 6 electrons • Element B has 5 protons, 6 neutrons, and 5 electrons • Element C has 6 protons, 8 neutrons, and 6 electrons • Make a statement about the relationship between Elements A, B, and C using the terms “element” and “isotope”

  20. Mass Number (A) • Defines a particular isotope • Total number of protons & neutrons • Not directly found on the periodic table • Standard notation: element-A • Uranium-235 • Uranium-238 To find # neutrons Mass # = p+ + no no = Mass # - p+ no = Mass # - atomic #

  21. QUICK CHECK! • Determine the number of protons, neutrons and electrons in Nickel-60 • Protons= • Neutrons= • Electrons= • What is the term used to describe the 60?

  22. Isotopes—Shorthand Notation • Z → identity of element • A→ particular isotope • #no = A - Z Mass # (p+ + no) Atomic # (p+=e-)

  23. U-235 and U-238 • Uranium-235 • Uranium-238

  24. You Try… • Silicon (Si) has three naturally occurring isotopes: Silicon-28, Silicon-29, and Silicon-30. Write the shorthand notation of each and determine the number of p+, n0, and e- in each. The atomic number of silicon is 14. Therefore: 28Si has 14p+, 14e-and 14n0(28-14) 29Si has 14p+, 14e-and 15n0(29-14) 30Si has 14p+, 14e-and 16n0 (30-14)

  25. What if the number of protons and electrons aren’t equal? Ion-an atom with a charge Forms by changing number of electrons Still the same element, only charged Lose electron + ion Cation Gain electron - ion Anion Finding # of electrons

  26. *Common mistake… • When you lose electrons, the ion becomes positive because there are more protons than electrons • Do not associate a positive charge with adding electrons! Example: Oxygen (Atomic number 8)

  27. Ion Notation Charge is represented as a superscript after the atomic symbol Examples: Na+ or O2-

  28. Another really corny joke • An atom walked into a police station and said he wanted to report a lost electron. The officer asked “Are you sure you lost it?” • He replied “I'm positive.”

  29. Practice 2. How many electrons are in the following ions? A. Ca2+ B. Cl-

  30. Essential Question How is the average atomic mass of an isotope calculated?

  31. Units for Masses on the P.T. • Units • amu’s (atomic mass units, referring to mass of 1 atom) • Grams (referring to mass of 1 mole of atoms) • Example (Carbon): • 12.01 amu = mass of 1 “average” carbon atom • 12.01 g = mass of 1 mole of “average” carbon atoms

  32. Decimal #s on the PT • What about the decimal numbers on the periodic table? What do they mean?

  33. Masses on the P.T. • Average Atomic Mass • the weighted average mass of an element’s isotopes • Weighted Average – depends on % abundance of each isotope • Percent Abundance—%of an element that is a particular isotope • More abundant isotopes = greater contribution

  34. Calculating a Weighted Grade Calculating average atomic mass is like finding a weighted grade. Joey’s grades are as follows: 85% on tests, 92% on homework, 87% on quizzes, 90% on labs His class is weighted as follows: 40% on tests, 30 % on labs, 20% on quizzes, 10% on labs What affects his overall grade the most? What is Joey’s overall grade in this class?

  35. Joey’s Grade Grade=(85)*(0.40)+92*(0.30)+87*(0.20)+90*(0.10) Grade=88%

  36. Calculating Avg. Atomic Mass • Convert all/any percents to decimal form • Multiply each mass by its percent abundance • Add together for avg. atomic mass

  37. Atomic Mass • The natural abundance for boron isotopes is: 19.9% 10B (10.013 amu) and 80.1% 11B (11.009amu). Calculate the atomic weight of boron.  Average Atomic Mass = 0.199 (10.013amu) + 0.801 (11.009amu) How does this number compare to the atomic mass on the periodic table?

  38. You Try… • Naturally occurring copper consists of 69.15% copper-63 (62.929601 amu) and 30.85% copper-65 (64.927794 amu). Find the average atomic mass of copper.

  39. Corny Joke  • What did the ion say to the other ion? • “I’ve got my ion you!”

  40. Finding Relative Abundances Work backwards to determine the relative abundances of isotopes. You must use a system of equations to do this.

  41. Relative Abundances Gallium consists of two naturally occuring isotopes with masses of 68.926 and 70.925 amu. The average atomic mass of Ga is 69.72 amu. Calculate the abundance of each isotope. 69.72amu=68.926x + 70.925y 1.0 = x + y

  42. Relative Abundances 69.72amu=68.926(1-y) + 70.925y 69.72amu = 68.926 – 68.926y + 70.925y 0.794 = 1.999y y = 0.3972 Isotope 70.925 = 39.72% Isotope 68.926 = 60.28%

  43. You Try…. Phosphorus has two isotopes, 32P (mass = 32.00amu) and 30P(mass = 30.00amu). The experimentally determined mass of phosphorus is 30.973amu. What are the percentage abundances of the two isotopes? 30.973= 32.00x + 30.00(1-x) 30.973 = 32.00x + 30.00 – 30.00x 0.973 = 2.00x x= 0.4865 32P= 48.65% 30P=51.35%

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