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Exploring the Evolution of Atomic Theory

Discover the historical icons and fundamental laws of early atomic theory, from Democritus to Rutherford. Learn about Dalton's Atomic Theory, subatomic particles, and the structure of the atom. Unravel the mysteries of the nucleus and electron cloud in this educational journey.

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Exploring the Evolution of Atomic Theory

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  1. Honors Chemistry chapter 3 Atoms: The Building Blocks of Matter

  2. PART 1 Icons ofEarly Atomic Theory

  3. Icons in Early Atomic Theory • Democritus [400 B.C] • Greek philosopher • Hypothesized: Nature has a basic indivisible particle of which everything is made of • Called this particle an atom • Greek “atomos” = indivisible

  4. 1790s – Discovery of Basic Laws • Law of Conservation of Mass • Mass is neither created nor destroyed during ordinary chemical reactions or physical changes • Law of Definite Proportions • A chemical compound contains the same elements in exactly the same proportions by mass regardless of size of sample or source of compound • i.e. Every sample of table salt is made of 39.34% Na and 60.66% Cl • i.e. H2O always has 2 atoms of H and 1 atom of O

  5. Basic Laws Continued • Law of Multiple Proportions • If two or more different compounds are composed of the same two elements then the ratio of the masses of the second element combined with a certain mass of the first element is always a ratio of small whole numbers • i.e. CO and CO2 • CO = 1.00g of C and 1.33 g of O • CO2 = 1.00 g of C and 2.66 g of O • The ratio of the second element is 2.66 to 1.33 or 2 to 1

  6. Icons ofEarly Atomic Theory Continued • John Dalton [1808] • English schoolteacher – liked nature and weather • Developed: Dalton’s Atomic Theory

  7. Dalton’s Atomic Theory • All matter is composed of extremely small particles called atoms • Atoms of a given element are identical in size, mass and other properties and are different from atoms of other elements • Atoms cannot be subdivided, created, or destroyed • Atoms of different elements combine in simple whole number ratios to form chemical compounds • In chemical reactions, atoms are combined, separated or rearranged

  8. Issues withDalton’s Atomic Theory • Atoms can be split into even smaller particles (nuclear chemistry) and aren’t indivisible • i.e. nucleus, protons, electrons • A given element can have different masses • i.e. isotopes

  9. Structure of the Atom • Today’s definition of the atom • Atom = Smallest particle of an element that retains the chemical properties of that element • Two regions • Nucleus • Very dense, small center of the atoms • Protons and neutrons • Electron Cloud • Region occupied by electrons • Subatomic particles • Protons, neutrons, electrons

  10. Icons of Early Atomic Theory Continued • J.J. Thomson [1897] • Discovered: The 1st subatomic particle: the negatively charged electron • Used a Cathode Ray Experiment • Cathode Ray Tube – Electric current passed through a metal disk to another metal disk in a gas at low pressure (vacuum sealed tube) • i. e. neon signs and ‘old-fashioned’ television sets

  11. Cathode Ray Experiment • When a current passed through the cathode ray tube, the surface of the tube opposite the cathode glowed • Glow was hypothesized to be stream of particles called a cathode ray • Ray affected by magnetic fields • Attracted to positive charge • Deflected from negative charge • http://www.youtube.com/watch?v=7YHwMWcxeX8&NR=1

  12. Discovery of the1st subatomic particle • Thomson measured the ratio of the charge of the particles to their mass • Same ratio no matter what metal or gas was used • Named this particle an electron • http://www.youtube.com/watch?v=IdTxGJjA4Jw&feature=related

  13. Thomson’sPlum Pudding Model • Atoms are electrically neutral • Must have positive charges to balance the negatively charged electrons • Electrons have a lot less mass than atoms • Other particles must account for their mass • Plum Pudding Model • positively charged sphere with electrons dispersed through it

  14. Icons ofEarly Atomic Theory Continued • Robert Millikan [1909] • Discovered: The measurement of an electron charge • Oil Drop Experiment • Measured the difference in velocity of oil droplets • Charged droplets (ionizing radiation) vs. uncharged • http://www.youtube.com/watch?v=XMfYHag7Liw&feature=related

  15. Icons ofEarly Atomic Theory Continued • Ernest Rutherford (with Hans Geiger and Ernest Marsden) [1911] • Discovered: A new atomic model • Gold Foil Experiment • Bombarded thin piece of gold foil with alpha particles • Expected alpha particles to pass through with minimal deflection • Surprised when 1 in 8000 deflected back to source • It was “as if you had fired a 15 inch artillery shell at a piece of tissue paper and it came back and hit you” • http://www.youtube.com/watch?v=wzALbzTdnc8&feature=related

  16. Rutherford’sNew Model of the Atom • Discovered the nucleus is a small densely packed volume of positive charge • Size comparison • Nucleus = marble • Whole Atom = football field • At this point in history, we were not sure where the electrons were – stay tuned for more in Chapter 4

  17. PART 2 Inside the Atom

  18. Inside the Nucleus • 2 types of particles • Protons • positively charged = +1 • made up of quarks • Neutrons • neutral = 0 charge • Made up of quarks • Mass in the nucleus • Protons = 1.673 x 10-27 • Neutrons = 1.675 x 10-27 • To simplify, both have mass of 1 amu (atomic mass unit)

  19. How does the nucleus stay together? • Strong Nuclear Forces • Two protons extremely close = strong attraction • Two neutrons extremely close = strong attraction • Neutrons and Protons extremely close = strong attraction • Strong nuclear forces overcome the repulsion of like positive charges to keep the nucleus together!!!

  20. Where are the Electrons? • In the Electron Cloud • A cloud of negative charge outside of the nucleus • More on this later........ • Electrons = Negatively charged particles with almost no mass (9.109 x 10-31)

  21. Characteristics of Atoms • Atomic Number • Equal to the number of protons and specific to each type of element • Identifies the element • # of protons is what give that element its characteristic properties • Elements with different protons are NOT THE SAME ELEMENT!!!

  22. Neutral Atoms • Neutral atoms • total positive charge equals the total negative charge • # protons (+1 each) = # electrons (-1 each)

  23. Isotopes • Atoms of the same element (i.e. same # of protons) that have differing number of neutrons • Isotopes of the same element • have different masses • do not differ significantly in chemical behavior

  24. Mass Number • Mass number = #protons + # neutrons • Average Atomic Mass • Every element has isotopes • The periodic table takes into account all naturally occurring isotopes of an element and averages them

  25. Properties of Subatomic Particles

  26. Ions • Atoms with a charge • Negative – more electrons than protons • Positive – more protons than electrons • Charge = #protons - # electrons • Magnesium atom with 12 protons and 10 electrons has a charge of +2

  27. Average Atomic Mass • Average Atomic Mass listed on the periodic table • UNIT is amu = atomic mass unit • 1 amu is a standard • Equal to 1/12 the mass of a C-12 atom • Takes into account all an elements isotopes and the frequency of each isotopes occurrence in nature • How to Calculate Average Atomic Mass

  28. Example of Calculating the Average Atomic Mass – Hydrogen • There are two naturally occurring isotopes of hydrogen • Hydrogen with 1 proton and zero neutrons • Hydrogen with 1 proton and one neutron • Differentiating between the two isotopes (symbol – mass number) • Calculation:

  29. PART 3: The Mole

  30. Amadeo Avogadro • Amadeo Avogadro [1776] • Lawyer turned professor of mathematical physics • Theorized: equal volumes of all gases at the same temperature and pressure contain the same number of particles. • After Avogadro’s death Avogadro’s number was determined • Avogadro’s number is simply a unit of measure • 1 mole = 6.023 x 1023 of any substance • Typically used to talk about particles (atoms, compounds, etc.)

  31. The Mole Unitput into perspective!!! • One mole of rice grains is more grains than the total number of grains grown since the beginning of time. • A mole of rice would occupy a cube about 120 miles on each edge. • A mole of marshmallows would cover the US to a depth of 600 miles • A mole of hockey pucks would be equal in volume to the moon • A mole of basketballs would just about fit perfectly into a ball bag the size of the earth.

  32. Mole Video • http://www.youtube.com/watch?v=Hj83oRHdezc

  33. Mole Calculations • 1 mole = 6.02 x 1023 of anything (atoms, molecules, formula units, particles, etc.) • Use dimensional analysis when solving: • Conversion factor: 1 mole = 6.02 x 1023 atoms, particles, formula units, etc. • Practice: • If I have 3.5 moles of carbon atoms, how many molecules do I have? • If I have 5.43 x 1031 molecules of carbon dioxide, how many moles do I have?

  34. Molar Mass • Molar Mass • The mass of one mole of a substance • The molar mass of an element can be found on the periodic table • Same as the average atomic mass • 1 amu = 1 gram/mole • E.g. Average atomic mass of C = 12.011 amus • Molar mass of C = 12.011 grams/1 mole

  35. Calculations Using the Molar Mass – Use Dimensional Analysis • Calculate the number of grams of carbon in 3.25 moles of carbon. • Calculate the number of moles of hydrogen in 6.05 grams of hydrogen. • Calculate the number of atoms of carbon in 15.00 grams of carbon.

  36. Mole Sing A Long October 24, 2005

  37. It’s a Unit After All”http://www.youtube.com/watch?v=ZNnbZgnSohk • A mole of laughter, a mole of tears • A mole of atoms, a mole of cheer • The name of that measure • Is a real chemist’s treasure • It’s a unit after all • Chorus • It’s a unit, after all • It’s a unit, after all • It’s a unit after all • It’s a unit after all

  38. It’s a Unit After All” • http://www.youtube.com/watch?v=at_9A_gfln0 • A chemist’s friend, tried and ture, • An Avogadro would stand by you. • And any chemist anywhere, • Would stand up and swear, • It’s a unit after all • Chorus • It’s a unit after all, etc.

  39. “It’s a Unit After All” • Six point oh two times ten to twenty-three • A number to live by in chemistry • So this is October 24th • Don’t be absurd, for • It’s a unit after all • Chorus • It’s a unit after all • It’s a unit after all • It’s a unit after all • I’ts a unit after all

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