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Atoms, Ions and Molecules

Atoms, Ions and Molecules

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Atoms, Ions and Molecules

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  1. Atoms, Ions and Molecules Chapter 2

  2. Dalton’s Atomic Theory (1808) • All matter is composed of extremely small particles called atoms • Atoms of a given element are identical in size, mass, and other properties; atoms of different elements differ in size, mass, and other properties John Dalton • 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

  3. Atomic Theory of Matter • All matter is made of tiny indivisible particles called atoms. • This proposal has been verified experimentally. Single atoms of a variety of elements have been photographed with a scanning transmission electron microscope.

  4. Atomic Theory of Matter 2. Atoms of the same element are identical and atoms of a different element have different masses and chemical properties. • However you will learn that atoms of the same element can have different masses.

  5. Atomic Theory of Matter • Atoms of different elements combine in whole number ratios to form compounds • One molecule of water always consists of two atoms of hydrogen and one atom of oxygen.

  6. Atomic Theory of Matter 4. Chemical reactions involve the rearrangement of atoms. No new atoms are created or destroyed. • Modern research has altered this proposal. Atoms are not indestructible and may lose their identity when split during nuclear reactions. However Dalton’s proposal remains true, for chemical reactions.

  7. Law of Multiple Proportions • Atoms of two or more elements may combine in different ratios to produce more than one compound.

  8. Law of Constant Composition • A compound always contains two or more elements combined in a definite proportion by mass.

  9. Law of Conservation of Mass • The total mass of materials present after a chemical reaction is the same as the total mass before the reaction • This is the basis for which postulate?

  10. Parts of an Atom • J. J. Thomson - English physicist. 1897 • Made a piece of equipment called a cathode ray tube. • Determined the charge to mass ratio • It is a vacuum tube - all the air has been pumped out.

  11. Voltage source Thomson’s Experiment + - Vacuum tube Metal Disks

  12. Voltage source Thomson Experiment + - • Passing an electric current makes a beam appear to move from the negative to the positive end

  13. Voltage source Thomson’s Experiment + - • By adding an electric field

  14. Voltage source Thomson’s Experiment + - • By adding an electric field he found that the moving pieces were negative

  15. What Did Thomson Demonstrate? • Cathode rays: • Travel in straight lines • Are negatively charged • Are deflected by electric and magnetic fields

  16. Thomson’s Model • Found the electron • Said the atom was like plum pudding • A bunch of positive stuff, with the electrons able to be removed

  17. Robert Millikan • American Scientist • Goal: Determine the charge on the electron to determine its mass

  18. The Discovery of Atomic Structure • Cathode Rays and Electrons • Consider the following experiment: • Oil drops are sprayed above a positively charged plate containing a small hole. • As the oil drops fall through the hole, they are given a negative charge. • Gravity forces the drops downward. The applied electric field forces the drops upward. • When a drop is perfectly balanced, the weight of the drop is equal to the electrostatic force of attraction between the drop and the positive plate. Chapter 2

  19. Atomizer Oil droplets + - Oil Telescope Millikan’s Experiment

  20. Millikan’s Experiment X-rays X-rays give some drops a charge.

  21. Millikan’s Experiment Some drops would hover From the mass of the drop and the charge on the plates, he calculated the mass of an electron

  22. Radioactivity • Discovered by accident • French scientist Henri Bequerel • Studying pitchblende (oxides of uranium) • Discovered that it spontaneously emits high energy radiation • Three types • alpha- helium nucleus (+2 charge, large mass) • beta- high speed electron • gamma- high energy light

  23. (Uranium compound) 2.2

  24. Rutherford’s Experiment • Ernest Rutherford English physicist. (1910) • Believed in the plum pudding model of the atom. • Used uranium to produce alpha particles.

  25. Rutherford’s Experiment • Aimed alpha particles at gold foil by drilling hole in lead block. • Since the mass is evenly distributed in gold atoms alpha particles should go straight through. • Used gold foil because it could be made atoms thin.

  26. Rutherford’s Experiment • When the alpha particles hit a florescent screen, it glows.

  27. Rutherford’s Experiment Florescent Screen Uranium Lead block Gold Foil

  28. What he expected

  29. Why ?? • The alpha particles would pass through without changing direction very much • The negative charges were spread out evenly. Alone they were not enough to stop the alpha particles

  30. What he got

  31. How He Explained It • Atom is mostly empty • Small dense, positive piece at center • Alpha particles are deflected by it if they get close enough

  32. + How He Explained It

  33. Density and the Atom • Since most of the particles went through, it was mostly empty. • Because the pieces turned so much, the positive pieces were heavy. • Small volume, big mass, big density • This small dense positive area is the nucleus

  34. + Discovery of the Neutron + James Chadwick bombarded beryllium-9 with alpha particles, carbon-12 atoms were formed, and neutrons were emitted. *Walter Boethe Dorin, Demmin, Gabel, Chemistry The Study of Matter 3rd Edition, page 764

  35. Modern View • The atom is mostly empty space • Two regions • Nucleus- protons and neutrons • Electron cloud- region where you might find an electron

  36. Structure of Atom • There are two regions: • The nucleus: with protons and neutrons • Almost all the mass • Electron cloud- Most of the volume of an atom • The region where the electron can be found

  37. Size of Atom • Atoms are small. • Measured in picometers, 10-12 meters • Hydrogen atom, 32 pm radius

  38. Size of Atom • Nucleus tiny compared to atom • IF the atom was the size of a stadium, the nucleus would be the size of a marble. • Radius of the nucleus near 10-15m. • Density near 1014 g/cm3

  39. QUARKS equal in a neutral atom Most of the atom’s mass. Subatomic Particles ATOM NUCLEUS ELECTRONS NEUTRONS PROTONS Negative Charge Positive Charge Neutral Charge

  40. Subatomic particles Actual mass (g) Relative mass Name Symbol Charge Electron e- -1 1/1840 9.11 x 10-28 Proton p+ +1 1 1.67 x 10-24 Neutron no 0 1 1.67 x 10-24

  41. A X Mass Number Element Symbol Z Atomic Number 1 3 2 H (D) H (T) H 1 1 1 235 238 U U 92 92 Atomic number (Z) = number of protons in nucleus Mass number (A) = number of protons + number of neutrons = atomic number (Z) + number of neutrons Isotopes are atoms of the same element (X) with different numbers of neutrons in their nuclei

  42. Symbols • Find the • number of protons • number of neutrons • number of electrons • Atomic number • Mass Number 19 F 9

  43. Symbols • Find the • number of protons • number of neutrons • number of electrons • Atomic number • Mass Number 80 Br 35

  44. Symbols • if an element has an atomic number of 34 and a mass number of 78 what is the • number of protons • number of neutrons • number of electrons • Complete symbol

  45. Atomic Mass • How heavy is an atom of oxygen? • There are different kinds of oxygen atoms. • More concerned with average atomic mass. • Based on abundance of each element in nature. • Don’t use grams because the numbers would be too small

  46. Measuring Atomic Mass • Unit is the Atomic Mass Unit (amu) • One twelfth the mass of a carbon-12 atom. • Each isotope has its own atomic mass we need the average from percent abundance.

  47. Calculating averages • You have five rocks, four with a mass of 50 g, and one with a mass of 60 g. What is the average mass of the rocks? • Total mass = 4 x 50 + 1 x 60 = 260 g • Average mass = 4 x 50 + 1 x 60 = 260 g 5 5

  48. Calculating averages • Average mass = 4 x 50 + 1 x 60 = 260 g 5 5 5 • Average mass = .8 x 50 + .2 x 60 • 80% of the rocks were 50 grams • 20% of the rocks were 60 grams • Average = % as decimal x mass + % as decimal x mass + % as decimal x mass +

  49. Atomic Mass • Calculate the atomic mass of copper if copper has two isotopes. 69.1% has a mass of 62.93 amu and the rest has a mass of 64.93 amu.

  50. 17 Cl 35.4594 100 Mass spectrum of chlorine. Elemental chlorine (Cl2) contains only two isotopes: 34.97 amu (75.53%) and 36.97 (24.47%) 90 Cl-35 80 70 AAM = (34.97 amu)(0.7553) + (36.97 amu)(0.2447) 60 AAM = (26.412841 amu) + (9.046559 amu) AAM = 35.4594 amu 50 Abundance 40 30 Cl-37 20 10 0 36 37 35 34 Mass