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The Atom

The Atom. Nuclear Model of the Atom. An atom is an electrically neutral particle C omposed of protons, neutrons, and electrons. Atoms are spherical in shape Have a tiny, dense nucleus of positive charge surrounded by one or more negatively charged electrons

pandora-gilmore
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The Atom

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  1. The Atom

  2. Nuclear Model of the Atom • An atom is an electrically neutral particle • Composed of protons, neutrons, and electrons. • Atoms are spherical in shape • Have a tiny, dense nucleus of positive charge surrounded by one or more negatively charged electrons • Nucleus contains 99.7 % of the mass of an atom.

  3. Basic Structure of the Atom

  4. Subatomic Particles

  5. Subatomic Particles • Most of the atom consists of fast-moving electrons traveling through the space around the nucleus. • Atoms are neutral in charge • The number of protons = the number of electrons • Atomic Number: • The number of Protons in the nucleus • determines the identity of an atom • No two types of atoms have the same number of protons • An Element consists of atoms with the same number of protons.

  6. Atomic Number(Z): the number of protons in the nucleus of each atom ofthe element • Atomic Number = Protons • Elements are arranged in the periodic table from left to right in order of increasing atomic number. • Atomic number identifies the element

  7. Isotopes and Mass Number • Atoms can have different mass. • Naturally occurring elements are a mixture of atoms that have different numbers of neutrons. • Atoms with the same number of protons, but different number of neutrons are called Isotopes.

  8. Hydrogen Has Three Isotopes: • Protium 1 proton 1 electron 0 neutrons • Deuterium 1 proton 1 electron 1 neutron • Tritium 1 proton 1 electron 2 neutrons

  9. Isotopes and Mass Number • Isotopes: atoms of the same element that have different masses. • Same number of protons • Different number of neutrons • Different Mass • Atoms of different isotopes have different masses so identity is given by name and mass. • Mass Number = protons and neutrons

  10. Isotope (Nuclear) Symbols: Consists of three parts • the symbol of the element • the atomic number of the element • the mass number of the specific isotope.

  11. Nuclear Symbol • Read as Helium-4

  12. Isotope Names Element Name-Mass Number Helium-4 Potassium-39 Hydrogen-3

  13. Neutrons = Mass Number – Atomic Number Atomic Number: _________ Mass Number: __________ # of Protons: _________ # of Electrons: _______ # of Neutrons: ________ Name of Isotope: ____________

  14. Example Atomic Number: _________ Mass Number: __________ # of Protons: _________ # of Electrons: _______ # of Neutrons: ________ Name of Isotope: ____________

  15. Example Atomic Number: _________ Mass Number: __________ # of Protons: _________ # of Electrons: _______ # of Neutrons: ________ Isotope Name: _________

  16. Draw the isotope symbol for Calcium with 21 neutrons.

  17. Comparing Potassium Isotopes Potassium-41 • protons = ________ • electrons = ______ • neutrons = _______ Potassium-40 • protons = ________ • electrons = ______ • neutrons = _______ Potassium-39 • Protons = __________ • Electrons = _________ • Neutrons = __________ 

  18. Relative Atomic Masses • Atomic masses measured in grams are very small. • Example: An atom of Oxygen-16 has a mass of 2.657 x 19-23grams. • It is most convenient to measure mass in relative atomic masses.

  19. To set up a relative scale of atomic masses: • One atom is chosen and assigned a relative mass value, • All the other masses are expressed in relation to this defined standard. • Carbon-12 is the chosen standard. • A single atom of C-12 is assigned a mass of exactly 12 atomic mass units(u).

  20. An AMU • One amu is exactly 1/12th of the mass of a carbon-12 atom or 1.660 5402 x 10-24grams. • The atomic mass of a carbon-12 atom is exactly 12 u. • Atomic Mass: The mass of an atom expressed in atomic mass units.

  21. Average Atomic Mass • Atomic Masses given for the elements on the periodic table are weighted averages for the naturally occurring mixtures of isotopes, called average atomic mass. • Average Atomic Mass depends on mass and relative abundance of the isotopes.

  22. To calculate average atomic mass: • Atomic mass x relative abundance • Add results.

  23. Example Problem: Naturally occurring copper consists of 69.17% Cu-63, mass of 62.939 598u and 30.83% Cu-65, mass of 64.927 793u • 0.6917 x 62.939 598u = 43.535u • 0.3083 x 64.927 793u = 20.017

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