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The Atom Atomic Number and Mass Number Isotopes

The Atom Atomic Number and Mass Number Isotopes. Democritus (460 - 370 B.C.).

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The Atom Atomic Number and Mass Number Isotopes

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  1. The Atom Atomic Number and Mass Number Isotopes

  2. Democritus(460 - 370 B.C.) The goal of the Greek philosophers was to explain the natural world. In an effort to organize many of the phenomena they observed, philosophers believed that a single "primary matter" existed. It was of this primary matter, modified in various ways, that all other things were created. Democritus expanded the idea to state that matter was composed of small particles called "atoms" that could be divided no further.

  3. John Dalton(1766 - 1844) During the 19th century, a vast amount of data on how substances react with each other was collected. From this data, some simple laws of chemical reactivity had been devised. Among these were the law of conservation of matter and law of multiple proportions. While others had proposed very similar theories, John Dalton is usually credited with developing the atomic theory.

  4. Dalton's theory can be summarized as follows: 1. Matter is composed of small particles called atoms. 2. All atoms of an element are identical, but are different from those of any other element. 3. During chemical reactions, atoms are neither created nor destroyed, but are simply rearranged. 4. Atoms always combine in whole number multiples of each other. For example, 1:1, 1:2, 2:3 or 1:3.

  5. J.J. Thomson(1856 - 1940 At approximately the same time as radioactivity was being investigated, J.J. Thomson and others were performing experiments with cathode ray tubes. A cathode ray tube is an evacuated tube that contains a small amount of gas between two metallic plates. When a potential is placed between the cathode (the negatively charged plate) and the anode (the positively charged plate) a "ray" of electric current passes from one plate to the other. Thomson discovered that this ray was actually composed of particles.

  6. JJ Thomson Plum-Pudding Model Cathode Ray Tube

  7. When a second set of plates is placed around the tube, the ray is bent toward the positive plate indicating that the ray is composed of negatively charged particles. By varying the potential on the plates, Thomson was able to determine the mass to charge ratio of these particles.

  8. He concluded that the negatively charged particles were subatomic particles that were part of every atom. He further surmised that, since atoms were electrically neutral, the atom must also contain some positive charge. Based on these conclusions Thomson proposed that an atom was composed of a spherical ball of positive charge with negative charge imbedded in it. The negative charge particles would later become known as electrons. This was later known as “The Plum Pudding Model”

  9. In 1909, Rutherford set a fellow scientist, Hans Geiger, and a student, Ernest Marsden, to work on this problem. They devised a system that allowed alpha particles (the nuclei of helium atoms) to be shot at a very thin piece of gold foil and the trajectory of the particles monitored. They observed that while most of the particles passed through the foil with little or no deflection, some were deflected to a great degree.

  10. Rutherford

  11. Since the gold film was so thin, Rutherford proposed that all of the deflections observed were from single encounters of alpha particles with the atom. In order to deflect the relatively large and swiftly moving alpha particles to such a large extent, a large force was required. This force, he contended, could only be caused by a large concentration of positive charge within the atom. This large concentration of charge was located at the center of the atom and became known as the nucleus.

  12. Bohr

  13. Bohr expanded upon this theory by proposing that electrons travel only in certain successively larger orbits. He suggested that the outer orbits could hold more electrons than the inner ones, and that these outer orbits determine the atom's chemical properties.

  14. Bohr’s Model – Quantum Model Nucleus Electron Circular Orbit Energy Levels

  15. Atomic Theory • Atoms are building blocks of elements • Similar atoms in each element • Different from atoms of other elements • Two or more different atoms bond in simple ratios to form compounds

  16. Subatomic Particles Particle Symbol Charge Relative Mass Electron e- 1- 0 Proton p+ + 1 Neutron n 0 1 LecturePLUS Timberlake

  17. Location of Subatomic Particles

  18. Atomic Number Counts the number of protons in an atom LecturePLUS Timberlake

  19. Periodic Table • Represents physical and chemical behavior of elements • Arranges elements by increasing atomic number • Repeats similar properties in columns known as chemical families or groups LecturePLUS Timberlake

  20. Periodic Table 1 2 3 4 5 6 7 8 11 Na LecturePLUS Timberlake

  21. Atomic Number on the Periodic Table 11 Na Atomic Number Symbol LecturePLUS Timberlake

  22. All atoms of an element have the same number of protons 11 Na 11 protons Sodium LecturePLUS Timberlake

  23. Learning Check State the number of protons for atoms of each of the following: A. Nitrogen 1) 5 protons 2) 7 protons 3) 14 protons B. Sulfur 1) 32 protons 2) 16 protons 3) 6 protons C. Barium 1) 137 protons 2) 81 protons 3) 56 protons LecturePLUS Timberlake

  24. Number of Electrons • An atom is neutral • The net charge is zero • Number of protons = Number of electrons • Unless it is an ion. LecturePLUS Timberlake

  25. Mass Number Counts the number of protons and neutrons in an atom LecturePLUS Timberlake

  26. Atomic Symbols • Show the mass number and atomic number • Give the symbol of the element mass number 23 Nasodium-23 atomic number11 LecturePLUS Timberlake

  27. How many neutrons does each of the following have? 16 31 65 O P Zn 8 15 30

  28. Practice An atom has 14 protons and 20 neutrons. A. It has atomic number 1) 14 B. It has a mass number of 3) 34 C. The element is 1) Si D. Another isotope of this element would be 3) 36X 14 LecturePLUS Timberlake

  29. Models of the Atom Dalton’s model (1803) Thomson’s plum-pudding model (1897) Rutherford’s model (1909) Bohr’s model (1913) Charge-cloud model (present) 1897 J.J. Thomson, a British scientist, discovers the electron, leading to his "plum-pudding" model. He pictures electrons embedded in a sphere of positive electric charge. 1911 New Zealander Ernest Rutherford states that an atom has a dense, positively charged nucleus. Electrons move randomly in the space around the nucleus. 1926 Erwin Schrödinger develops mathematical equations to describe the motion of electrons in atoms. His work leads to the electron cloud model. 1803 John Dalton pictures atoms as tiny, indestructible particles, with no internal structure. 1913 In Niels Bohr's model, the electrons move in spherical orbits at fixed distances from the nucleus. 1800 1805 ..................... 1895 1900 1905 1910 1915 1920 1925 1930 1935 1940 1945 1932 James Chadwick, a British physicist, confirms the existence of neutrons, which have no charge. Atomic nuclei contain neutrons and positively charged protons. 1924 Frenchman Louis de Broglie proposes that moving particles like electrons have some properties of waves. Within a few years evidence is collected to support his idea. 1904 Hantaro Nagaoka, a Japanese physicist, suggests that an atom has a central nucleus. Electrons move in orbits like the rings around Saturn. Greek model (400 B.C.) Dorin, Demmin, Gabel, Chemistry The Study of Matter , 3rd Edition, 1990, page 125

  30. Isotopes and AAM

  31. Isotopes • Atoms with the same number of protons, but different numbers of neutrons. • Atoms of the same element (same atomic number) with different mass numbers Isotopes of chlorine 35Cl 37Cl 1717 chlorine - 35 chlorine - 37

  32. Atomic Mass Na 22.99 • Listed on the periodic table • Gives the mass of “average” atom of each element compared to 12C • Average atom based on all the isotopes and their abundance % • Atomic mass is not a whole number LecturePLUS Timberlake

  33. Calculating Average Atomic Mass • Percent(%) abundance of isotopes • Mass of each isotope of that element • Weighted average = AAM = mass isotope1(% abundance) + mass isotope2(% abundance) + …

  34. Atomic Mass of Magnesium Isotopes Mass of Isotope Abundance 24Mg = 24.0 amu 78.70% 25Mg = 25.0 amu 10.13% 26Mg = 26.0 amu 11.17% Atomic Atomic Mass= LecturePLUS Timberlake

  35. Learning Check Gallium is a metallic element found in small lasers used in compact disc players. In a sample of gallium, there is 60.2% of gallium-69 (68.9 amu) atoms and 39.8% of gallium-71 (70.9 amu) atoms. What is the average atomic mass of gallium? LecturePLUS Timberlake

  36. Solution Ga-69 68.9 amu x 60.2 = 41.5 amu for 69Ga 100 Ga-71 (%/100) 70.9 amu x 39.8 = 28.2 amu for 71Ga 100 Atomic mass Ga = 69.7 amu LecturePLUS Timberlake

  37. Finding An Isotopic Mass A sample of boron consists of 10B (mass 10.0 amu) and 11B (mass 11.0 amu). If the average atomic mass of B is 10.8 amu, what is the % abundance of each boron isotope?

  38. Learning Check Copper has two isotopes 63Cu (62.9 amu) and 65Cu (64.9 amu). What is the % abundance of each isotope? (Hint: Check periodic table for atomic mass) 1) 30% 2) 70% 3) 100% LecturePLUS Timberlake

  39. Solution 2) 70% Solution 62.9X + 6490 = 64.9X = 6350 -2.0 X = -140 X = 70% LecturePLUS Timberlake

  40. LecturePLUS Timberlake

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