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Exploring Atoms: The Building Blocks of Matter

This chapter examines the development of atomic theory and provides an overview of the structure of atoms. Topics covered include early atomic theory, Dalton's atomic theory, modern atomic theory, and atomic structure.

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Exploring Atoms: The Building Blocks of Matter

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  1. Chapter 3 ATOMS: The Building Blocks of Matter West Valley High School General Chemistry Mr. Mata Chapter #3 • Atoms: The Building Blocks of Matter

  2. Standard 2a • To compare and contrast the atomic number and atomic mass.

  3. Essential Question • Describe how atomic theory developed to give us our modern view of the atom.

  4. 3-1 Early Atomic Theory • Atoms – smallest particle of matter. • Democritus (400 B.C) stated world was made of atoms. (atomos = “indivisible”). • Antoine Lavoisier (1800’s) discovered mass didn’t change after chemical rxn. • Proposed “matter can be changed, but it cannot be created or destroyed “ (Law of Conservation of Mass).

  5. Dalton’s Atomic Theory • All matter composed of atoms. • Atoms of same element identical in size, mass, properties; atoms of diff elements diff. in size, mass, properties. • Atoms can’t be subdivided, created, destroyed. • Atoms of different elements combine in whole-number ratios to form chemical compounds. • In chemical reactions, atoms are combined, separated, or rearranged.

  6. Modern Atomic Theory • All matter is composed of atoms. • Atoms of any one element differ in properties from atoms of another element. • Element’s average mass unique to element. • Atoms cannot be subdivided, created, or destroyed in ordinary chemical rxns. • Changes CAN occur in nuclear rxns!

  7. Atomic Structure • Atom- smallest particle of element that retains chemical properties of element. • Nucleus- positively charged, dense central portion of the atom; contains nearly all mass (~ 99.7%).

  8. Subatomic Particles

  9. The Atomic Scale • Most of the mass of the atom is in the nucleus (protons and neutrons) • Electrons are found outside of the nucleus (the electron cloud). • Most volume of atom=empty space.

  10. Famous Scientists

  11. Discovery of the Electron • J.J. Thomson (1897) used a cathode ray tube to deduce the presence of a negatively charged particle. • Discovered the – electron particle.

  12. (1906)Thomson awarded the Nobel Prize in chemistry for discovery of the electron. • The atom could be broken down into smaller particles.

  13. Thomson’s Atomic Model • Thomson believed electrons were like plums in a + charged “pudding”. • He called it the “plum pudding” model.

  14. Rutherford’s Gold Foil Experiment • Alpha particles are helium nuclei. • Particles fired at a thin sheet of gold foil. • Particle hits on screen (film) are detected.

  15. Rutherford’s Findings • Most particles passed right through screen. • Few particles deflected. • VERY FEW were greatly deflected. • Conclusions: • The nucleus is small. • The nucleus is dense. • The nucleus is + charged.

  16. Atomic number: number of protons in the nucleus of atom. # p(+) = # e(-) 6 C Carbon 12.011 Section 3-3

  17. Atomic Mass • number of protons & neutrons in the nucleus. • Atomic Mass=protons + neutrons • Atomic Mass C = 12.011 • Mass number = rounded atomic mass • Mass Number C = 12

  18. Nuclear Symbols 235 92 Mass number (p+ + n) Atomic Mass (p+ + n) Element symbol Atomic number (# of p+) U 235 92

  19. Hyphen Notation Sodium-23 (23 is the atomic mass) Sooo… 23- 11 (atomic #) = 12 for the # of neutrons. Atomic number of 11 is the # of protons (11) and electrons(11).

  20. Molar Relationships • 1 mole = 6.02 x 1023 atoms, molecules, or particles • Molar mass = mass of 1 mole of substance(Hint: Add up atomic masses) • Ex: Find the molar mass of: H2O = (H=1(2), 0=16), 18 grams CO2 = (C=12, O=16(2)), 44 grams H2SO4=(H=1(2),S=32,O=16(4)), 98 g

  21. Molar Masses Find molar masses: NaCl = 58 grams = 1 mole CaCl2 = 110 grams = 1 mole Fe(OH)2 = 90 grams = 1 mole H3PO4 = 98 grams = 1 mole C6H12O6 = 180 grams = 1 mole

  22. Calculations:Converting moles to grams • Given # of mole X ? g (look at periodic table)= g 1 mole • How many grams of lithium are in 3.50 moles of lithium? • 3.50 mole Li X 7 g = 24.5 g Li 1 mol

  23. Your turn… • How many grams of carbon are in 8.25 moles of carbon? • 8.25 mole C X 12 g = 99.0 g C 1 mol

  24. Your turn… • How many grams of uranium are in 21.5 moles of uranium? • 21.5 mole U X 238 g = 5117 g U 1 mol

  25. Calculations:Converting grams to moles • Given # of g X 1 mol = mol g (look at periodic table) • How many moles of lithium are in 18.2 grams of lithium? • 18.2 g Li X 1 mol Li = 2.6 mol Li 7 g

  26. Your turn… • How many moles of gold are in 150 grams of gold? • 150 g Au X 1 mol Au = 0.8 mol Au 197 g

  27. Your turn… • How many moles of H2O are in 120 grams of water? • 120 g H2O X 1 mol H2O= 6.7 mol H2O 18 g

  28. Avogadro’s Number • Number of particles in exactly one mole of a pure substance. • 6.02 x 1023 is called “Avogadro’s Number”. • Named in honor of the Italian chemist Amadeo Avogadro (1776-1855). I didn’t discover it. Its just named after me!

  29. Important Skill • To find electrons, protons, and neutrons for an element: • 40 Ca atomic mass (#p + #n) 20  atomic number (#p = #e) • # electrons = 20 (atomic number) • # protons = 20 (atomic number) • # neutrons = 20 (atomic mass–atomic #) • 40 – 20 = 20 neutrons

  30. Chapter 3 SUTW Prompt • Describe the contributions of Thomson, Millikan, & Rutherford to our understanding of the atom. • Complete an 11-13 sentence paragraph using the SUTW paragraph format. Hilite using green, yellow, and pink. • Due Date: Monday, September 11, 2017 (beginning of class).

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