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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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