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Chemistry Chapter 5

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Chemistry Chapter 5

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  1. ChemistryChapter 5

  2. Democritus – 4th century B.C. Greek philosopher who first proposed that matter is composed of tiny, indivisible particles called atoms. The first atomic model was not proposed until 2000 years later!

  3. John Dalton, an English schoolteacher, proposed the following atomic theory, in 1808. www.infoscience.fr/histoire/biograph/biograph.php3?Ref=96

  4. Dalton’s Atomic Theory: 1. All elements are composed of tiny, indivisible particles called atoms. 2. Atoms of the same element are identical and differ from those of other elements.

  5. 3. Atoms cannot be subdivided, created, nor destroyed. 4. Different elements combine in simple whole-number ratios to form chemical compounds.

  6. 5. Chemical reactions occur when atoms are separated, joined, or rearranged.

  7. Law of conservation of mass – mass is neither created nor destroyed. (See #3 and 5 above)

  8. Law of definite proportions – different samples of any pure compound contain the same elements in the same proportions by mass. For instance, both 23g of water and 72,000g of water will be 11.2% hydrogen and 88.9% oxygen by mass. (See #4 above)

  9. Law of multiple proportions – if two or more different compounds are composed of the same two elements, A and B, then the ratio of the masses of element B that combine with a certain mass of element A is always a ratio of small whole numbers. (See #4 above) Examples: NO, NO2, N2O4, N2O5

  10. The Modern Atomic Theory Not all aspects of Dalton’s theory have been proven correct: 1) atoms are divisible and 2) a given element can have atoms with different masses.

  11. The important aspects from Dalton’s theory are that: 1) all matter is composed of atoms and 2) atoms of any one element differ in properties from atoms of another element.

  12. An atom is defined as the smallest particle of an element that retains the properties of that element. We now know that atoms are not indivisible. They can be broken down into even smaller, more fundamental particles.

  13. 3 Subatomic Particles: electron, proton, and neutron.

  14. Discovery of the Electron Electrons are negatively charged subatomic particles.

  15. J. J. Thomson, an English physicist discovered them in 1897 by doing experiments involving passing an electric current through a tube containing gases at low pressure. http://www.sciencemuseum.org.uk/collections/treasures/thomp2.asp

  16. Thomson found that a beam, called a cathode ray tube, traveled from the cathode (-) to the anode (+), and that this beam was repelled by negative electrical charge. Since he knew opposite charges attract and like charges repel, he proposed that this cathode ray was a stream of tiny negatively charged particles moving at high speed. He named them electrons.

  17. Further studies by Robert Millikan, an American scientist, led to the determination of the properties of the electron in 1909 with his oil-drop experiment . Oil-drop experiment

  18. Oil-Drop Experiment: • His experiment involved measuring the charge on tiny oil drops. The charge of each drop was a multiple of 1.60x10-19 C. This was the charge of an individual electron. He used this charge, along with Thomson’s charge/mass ratio, to determine the mass of an electron. 1.6x10-19C x 1g = 9.11x10-28g 1.76x108C

  19. An electron has a charge of -1 and a mass of 1/1840, or 9.109x10-31 kg, the mass of a hydrogen atom.

  20. Discovery of the Atomic Nucleus Protons: Scientists knew that atoms are neutral (have no charge). So, if there are electrons with negative charge, there must also be particles with a positive charge. This led to the discovery of the proton.

  21. Protons have a +1 charge and are 1840 times more massive than the electron.

  22. Neutrons: Sir James Chadwick discovered them in 1932. Neutrons have no charge, but have essentially the same mass as a proton. http://nobelprize.org/physics/laureates/1935/

  23. Once subatomic particles were discovered, Dalton’s model of the atom had to be modified.

  24. The second model of the atom (Thomson’s “Plum Pudding Model”) proposed that electrons were evenly distributed throughout an atom filled uniformly with positively charged material.

  25. In 1911, Ernest Rutherford, a New Zealand native, tested this model with his gold foil experiment. http://www.vanderkroft.net/elements/images/portret/ernest_rutherford2.jpg

  26. Gold Foil Experiment: • He shot a beam of massive alpha particles (He+2) at a very thin sheet of gold foil. • He expected the alpha particles to pass easily through the foil, with little deflection. • He was shocked to see that even though most passed through without deflection, a small fraction was deflected at large angles and some even bounced straight backwards! • To explain this, he modified Thomson’s atomic model.

  27. Gold-Foil Experiment

  28. Rutherford’s model of the atom stated that the atom is mostly empty space with all the positive charge and almost all of the mass concentrated in a small region, which he called the nucleus.

  29. The tiny nucleus is composed of protons and neutrons. This is an image of silicon atoms arranged on a face of a crystal. It is impossible to "see" atoms this way using ordinary light. The image was made by a Scanning Tunneling Microscope, a device that "feels" the cloud of electrons that form the outer surface of atoms.

  30. How small is the nucleus? If the atom were the size of a football stadium, the nucleus would be about the size of a marble …10,000 times smaller!

  31. Properties of Subatomic Particles (Table 3-1, pg. 74)

  32. Going Further

  33. MORE SUBATOMIC PARTICLES???!!! • Leptons (elementary particles) • Electron • Mu-meson (muon) – more massive than electrons • Tau-meson (tau) – more massive than electrons • 3 types of neutrinos – almost massless

  34. Hadrons (made of quarks) • Mesons of many types • Composed of a quark and an antiquark • Baryons • Composed of 3 quarks of different colors • protons (2 up quarks + 1 down quark) • neutrons (1 up quark and 2 down quarks)

  35. Every particle has an antiparticle. The antiparticle of the electron is the positron (e+) Gluons – hold quarks together.

  36.  6 “flavors” of quarks 1. up 3. top (truth) 5. strange 2. down 4. bottom (beauty) 6. charm

  37. “Colors” of quarks (charge) +2/3 up, top, charm -1/3 down, bottom, strange

  38. Going Further

  39. Atomic number – the number of protons in the nucleus of an atom. The atomic number identifies an element.

  40. 1) What is the atomic number of aluminum? 13 2) Which element has 26 protons? Fe - iron

  41. In a neutral atom, the number of protons (p+) equals the number of electrons (e-). 3) How many p+ does a neon atom have? 10 4)How many e- does a neon atom have? 10 5)An element with atomic number = 6 has how many e-? 6

  42. Mass Number – The total number of protons and neutrons in the nucleus of an atom. mass number – atomic number = number of neutrons

  43. 6) Oxygen-18 has a mass number of 18. How many protons, neutrons and electrons does it have? Atomic # = 8 Therefore, p+ = 8 e- = 8 n0 = 18- 8 =10

  44. Isotopes – Atoms that have the same number of protons but different numbers of neutrons.

  45. When writing nuclear chemical symbols for isotopes, the mass number is written as a superscript and the atomic number is written as a subscript before the element symbol. Ex. 146C. This can also be written in hyphen notation as carbon-14.