Chapter 4 - The Atom
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Chapter 4 - The Atom

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  • 1. ATOMIC STRUCTURE AND THE PERIODIC TABLE Or How I Learned To Love Dead White Guys
  • 2. The Atom Chemistry is all about the investigation of matter and what it is. Eventually, chemists ask the question: what is matter made of? To answer that question, you need to go back 2000 yrs to a man named Democritus. Democritus believed that matter was made of a particle so small that it could not be broken any further than itself. He called this particle the atom . Aristotle, however, disagreed and downplayed this theory and so Democritus was pushed aside (and so was his theory). Democritus’ theory also didn’t explain chemical behavior so was disregarded.
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  • 5. It wasn’t until 1803 that the concept of the atom was widely introduced again. John Dalton, an English teacher, was able to relate atoms to matter. He proposed an atomic theory: 1. All matter is composed of atoms. 2. Atoms of an element are identical in size, mass, and other properties; different elements atoms’ differ in size, mass, and other properties. 3. Atoms cannot be created, destroyed or subdivided. 4. Atoms of different elements combine in simple ratios. 5. In chemical reactions, atoms are combined, separated, or rearranged.
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  • 7. Dalton’s theory was widely accepted for the rest of the 1800’s because it explained so many new observations of matter. The start of the 20th century saw a break down of Dalton’s theory, however. In 1903, JJ Thomson found that matter is made of negatively charged particles. The name of these negatively charged particles, originally called corpusles by Thomson, is now called electrons.
  • 8. This device produced what was known as “cathode rays”.
  • 9. Because cathode rays always have the same properties, regardless of the element used to produce them, it was concluded that electrons are present in atoms of all elements.
  • 10. After the discovery of the proton, Thomson came up with a model of the atom: This model had too many problems, however, and it would be 20 years before another one would clarify things. In 1909, Ernest Rutherford bombarded a thin sheet of gold foil w/ positively charged particles. He found that these (+) particles had no effect on the gold atoms. That can only mean the (+) charges aren’t equally surrounding the atom. This is called the raisin-filled fruitcake model. Also called “plum pudding” model.
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    • Coolio Animation
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  • 14. They must be in a small dense portion of the atom. And this small area must be (+) charged. The name of the positively charged area of the atom is called the nucleus . Rutherford then came up w/ a new model of the atom (which he stole from a Japanese scientist):
  • 15. In 1930 the final part of the atom was found: the neutron. New Model: Neutrons have no charge but roughly the same mass as a proton. e - = electrons p + = protons n 0 = neutron
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  • 17. Atomic number equals the number of protons Atoms are composed of electrons, protons, and neutrons. Protons and neutrons make up the small, dense nucleus. Electrons surround the nucleus and occupy most of the volume of the atom. How, then, are atoms of hydrogen different than atoms of oxygen? Elements are different because they contain different numbers of protons. The atomic number, Z, tells you how many protons are in an atom.
  • 18. Remember that atoms are always neutral because they have an equal number of protons and electrons. Therefore, the atomic number also equals the number of electrons the atom has . Each element has a different atomic number . For example, the simplest atom, hydrogen, has just one proton and one electron, so for hydrogen, Z = 1. The largest naturally occurring atom, uranium, has 92 protons and 92 electrons, so Z = 92 for uranium. The atomic number for a given element never changes.
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  • 20. Mass number equals the total number of subatomic particles in the nucleus The mass number, A , of an atom equals the number of protons plus the number of neutrons. A fluorine atom has 9 protons and 10 neutrons, so A = 19 for fluorine. Oxygen has 8 protons and 8 neutrons, so A = 16 for oxygen. This mass number includes only the number of protons and neutrons (and not electrons) because protons and neutrons provide most of the atom’s mass. Although atoms of an element always have the same atomic number, they can have different mass numbers.
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  • 22. Calculating the number of neutrons in an atom Atomic numbers and mass numbers may be included along with the symbol of an element to represent different isotopes. The two isotopes of chlorine are represented below. If you know the atomic number and mass number of an atom, you can calculate the number of neutrons it has.
  • 23. Mass number ( A ) - Atomic number ( Z ) = # of neutrons
  • 24. Isotopes of an element have different numbers of neutrons Neutrons can be added to an atom without affecting the number of protons and electrons the atom is made of. Many elements have only one stable form, while others have different “versions” of their atoms. Each version has the same number of protons and electrons as all other version but a different number of neutrons. These different versions, or isotopes , vary in mass but are all atoms of the same element because they each have the same number of protons. Three isotopes of hydrogen exist and have the same chemical properties because each is made of one proton and one electron.
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  • 26. Atomic Mass The mass of a proton or neutron is very small; an electron even smaller still. The mass of the largest atom is still incredibly small, because of its components. In determining the masses of atoms, and not having to work with the incredibly small masses of protons, etc., scientists used a reference isotope as a standard. The isotope of carbon was chosen, carbon-12. The isotope of carbon was assigned a mass of exactly 12 atomic mass units.
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  • 28. An atomic mass unit (amu) is defined as one-twelfth the mass of a carbon-12 atom . (For comparison, a helium-4 atom, with a mass of 4.0026 amu, has about one-third the mass of a carbon-12 atom.) A carbon-12 atom has six protons and six neutrons in its nucleus, and its mass is set as 12 amu. Therefore the mass of a single proton or single neutron is about one-twelfth of 12 amu, or 1 amu. You might predict, then, that the atomic mass of an element should be a whole number. However, that is not usually the case. For example, the atomic mass of chlorine (Cl) is 35.453 amu. How can this be decimal mass be explained? The atomic mass of an element is a weighted average mass of the atoms in a naturally occurring sample of the element.
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    • A weighted average mass reflects both the mass and the relative abundance of the isotopes as they occur in nature.
    • Now that you know that the atomic mass of an element is a weighted average of the masses of its isotopes, you can calculate atomic mass based on relative abundance.
    • To do this, you must know three values:
    • the number of stable isotopes of the element
    • the mass of each isotope, and
    • the natural percent abundance of each isotope .
    • Once you have these values for an element, multiply the atomic mass of each isotope by its abundance, expressed as a decimal, then add the results.
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  • 31. Example: Copper consists of 69.17% copper-63 (amu of 62.939 u); and, copper-65 which is 30.83% (64.927 u). The average atomic mass is: (.6917 x 62.939 u) + (.3083 x 64.927 u) = 63.55 u Example: An element has two naturally occurring isotopes. They have the following abundances: 50.69% (78.92 amu), and 49.31% (80.92 amu). What is the average atomic mass? (.5069 x 78.92) + (.4931 x 80.92) = 79.90 u