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UNIT 2

UNIT 2. Atoms , Molecules, and Ions. The Power of 10. http://micro.magnet.fsu.edu/primer/java/scienceopticsu/powersof10/. What to study for Unit 2. Look at syllabus (Objectives) List of ions to learn Homework recommendations Online Quiz, due Monday, 9/15 ( 5 pm ) Exam Tuesday, 9/16

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UNIT 2

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  1. UNIT 2 Atoms, Molecules, and Ions

  2. The Power of 10 http://micro.magnet.fsu.edu/primer/java/scienceopticsu/powersof10/

  3. What to study for Unit 2 • Look at syllabus (Objectives) • List of ions to learn • Homework recommendations • Online Quiz, due Monday, 9/15 (5 pm) • Exam Tuesday, 9/16 • Don’t forget Unit 1!!!

  4. A HISTORY OF THE STRUCTURE OF THE ATOM

  5. History • Greek Philosopher Democritus (460-370 B.C.): all matter composed of small atoms atomos = indivisible • 1803, John Dalton (brit.): atoms are the fundamental building blocks of matter

  6. Dalton's Postulates Each element is composed of extremely small particles called atoms.

  7. Dalton's Postulates All atoms of a given element are identical to one another in mass and other properties, but the atoms of one element are different from the atoms of all other elements.

  8. Dalton's Postulates Atoms of an element are not changed into atoms of a different element by chemical reactions; atoms are neither created nor destroyed in chemical reactions.

  9. Dalton’s Postulates Compounds are formed when atoms of more than one element combine; a given compound always has the same relative number and kind of atoms.

  10. John Dalton’s Atomic Theory(ca 1803) Each element is composed of extremely small particles called atoms. All atoms of a given element are identical. The atoms of different elements are different and have different properties (including different masses). Atoms of an element are not changed into different types of atoms by chemical reactions. Atoms are neither created nor destroyed in chemical reactions. This is theLaw of Conservation of Mass. Compounds are formed when atoms of more than one element combine. A given compound always has the same relative number and kind of atoms. This is theLaw of Constant Composition.

  11. John Dalton’s Atomic Theory Led him to deduce the Law of Multiple Proportions: When two or more elements combine to form more than one compound, the relative masses of the elements which combine will be in in the ratio of small whole numbers. In carbon monoxide, CO, 12 g carbon combine with 16 g oxygen. C:O ratio is 12:16 or 3:4. In carbon dioxide, CO2, 12 g carbon combine with 32 g oxygen. C:O ratio is 12:32 or 3:8.

  12. John Dalton’s Atomic Theory Almost right. A good start. very small Structure of the atom after Dalton (ca. 1810)

  13. J.J. Thomson (1897): Cathode Rays Atoms subjected to high voltages give off cathode rays.

  14. J.J. Thomson: Cathode Rays Cathode rays can be deflected by a magnetic field. Cathode rays are negatively charged particles (electrons). Electrons are in atoms.

  15. J.J. Thomson – The Electron “Plum pudding” model: Negative electrons are embedded in a positively charged mass. Electrons (-) Unlike electrical charges attract, and that is what holds the atom together. Positively charged mass Structure of the atom after Thomson (ca. 1900)

  16. Radioactivity • Radioactivity is the spontaneous emission of radiation by an atom. • First observed by Henri Becquerel (1852-1908). • Marie and Pierre Curie also studied it. • Nobel Prize in 1903 (physics).

  17. Studies of Natural Radioactivity Some atoms naturally emit one or more of the following types of radiation: alpha (α) radiation (later found to be He2+ - helium nucleus) beta (β) radiation (later found to be electrons) gamma (γ) radiation (high energy light) α Alpha particles Electrons (-) γ γ Positively charged mass α Somehow gamma radiation is in there, too. Structure of the atom after Becquerel (early 1900s)

  18. Radioactivity • Three types of radiation were discovered by Ernest Rutherford: •  particles (positive, charge 2+, mass 7400 times of e-) •  particles (negative, charge 1-) •  rays (high energy light)

  19. Ernest Rutherford (1910) Scattering experiment: firing alpha particles at a gold foil

  20. The Nuclear Atom Some alpha particles bounce off the gold foil. This means the mass of the atom must be concentrated in the center and is positively charged! Thompson’s model could not be correct.

  21. Ernest Rutherford The Nucleus and the Proton The mass is not spread evenly throughout the atom, but is concentrated in the center, the nucleus. The positively charged particles in the nucleus are protons. Electrons (-) are now outside the nucleus. Structure of the atom after Rutherford (1910)

  22. James Chadwick – The Neutron In the nucleus with the protons are particles of similar mass but no electrical charge called neutrons. Electrons (-) are now outside the nucleus in quantized energy states called orbitals. (From Niels Bohr and quantum mechanics) The positively charged particles in the nucleus are protons. n + n Structure of the atom after Chadwick (1932)

  23. Structure of the Atom proton (+) neutron electrons -responsible for the volume and size of the atom, negatively charged 10-10 m nucleus - responsible for the mass of the atom, positively charged 10-14 m

  24. Subatomic Particles • Protons and electrons are the only particles that have a charge. • Protons and neutrons have essentially the same mass. • The mass of an electron is so small we ignore it.

  25. Atomic Facts 1 amu = 1 atomic mass unit = 1.66054 x 10-24 g n + n Electrons are outside the nucleus in quantized energy states called orbitals.

  26. Symbols of Elements Elements are symbolized by one or two letters.

  27. Atomic Number All atoms of the same element have the same number of protons: The atomic number (Z)

  28. Atomic Mass The mass of an atom in atomic mass units (amu) is the total number of protons and neutrons in the atom.

  29. Atomic Number Carbon atom • The number of protons in the nucleus is called the atomic number Z. • Z determines the identity of an element. • Saying “the atomic number of an element is 6” is the same as saying “carbon.” • The number of electrons in the atom is also Z (because atoms have no net electric charge). • How many neutrons are in C? - proton - neutron

  30. Isotopes A 12C Z 6 • The number of protons and neutrons (nucleons) in an element is called the mass number A. A = Z + number of neutrons. • An element may have different numbers of neutrons but NOT different numbers of protons. • Atoms of an element with different numbers of neutrons are called isotopes of that element. - proton - neutron How many neutrons are in C? The answer is “it depends on the isotope.”

  31. Isotopes are atoms of the same element with different masses. Isotopes have different numbers of neutrons. 11 6 12 6 13 6 14 6 C C C C Isotopes

  32. Isotopes 12 6 12C or C-12 6 14 6 14C or C-14 6 16 8 16O or O-16 8 238 92 238U or U-238 92

  33. Atomic Masses Atomic masses are based on 12C. The mass of 12C (or C-12) is defined to be exactly 12 amu.

  34. Atomic Masses The mass (weight) shown in the periodic table is the mass of the element as its occurs naturally. If the element has more than one isotope, the mass shown is the weighted average of the masses of the isotopes. Mg has 3 isotopes. 24Mg 78.99% 23.985 amu 25Mg 10.00% 24.986 amu 26Mg 11.01% 25.983 amu weighted average of Mg: 0.7899x23.985 18.946 0.1000x24.986 2.499 0.1101x25.983 +2.861 24.31 amu atomic weight of Mg based on natural abundance: 24.31 amu

  35. Ions • Atoms can gain or lose electrons to become charged particles called ions. • A chemical particle that contains a positive or negative charge • Cationsare positively charged ions. • Formed when an atom loses electrons • Anionsare negatively charged ions. • Formed when an atom gains electrons

  36. + e- 1p Hydrogen atom 1p, 0 n, 1 e- Hydrogen ion (cation) 1p, 0 n, 0 e- 1p e- 1 H+ 1 H Ions Formation of a cation Net charge = 0 Net charge = +1

  37. Ions Formation of an anion 8p 8n 8e- 8p 8n + 2e- 10e- Oxygen atom 8p, 8 n, 8e- Oxygen ion (anion) 8p, 8n, 10e- 16 O2- 16 O Net charge = 0 Net charge = -2

  38. Nuclear Symbols Mass Number Charge Atomic Number X Charge = # p - # e-

  39. 16 8 O Nuclear Symbols • Using nuclear symbols to determine the number of p, n, e, and total charge Mass Number = Atomic Number = 16 8 # protons = atomic number = 8 # neutrons = Mass # - Atomic # = 16 - 8 = 8 # electrons = # protons = 8

  40. 16 8 2- O Nuclear Symbols Mass Number = Atomic Number = 16 8 # protons = atomic number = 8 # neutrons = Mass # - Atomic # = 16 - 8 = 8 # electrons = # protons - charge = 8 - (-2) = 10

  41. 137 56 2+ Ba Nuclear Symbols Mass Number = Atomic Number = 137 56 # protons = atomic number = 56 # neutrons = Mass # - Atomic # = 137 - 56 = 81 # electrons = # protons - charge = 56 - (+2) = 54

  42. Nuclear Symbols - Atoms Example: Write the nuclear symbol for the following atoms: 1) 50 p, 70 n 2) 17 e-, 20 n 120Sn 50 37Cl 17

  43. Nuclear Symbols - Ions Practice writing nuclear symbols from information given: • 53 p, 74 n, 54 e- 53 proton (= atomic number)  I 74 neutrons + 53 proton  mass number = 127 54 electrons (one more than protons)  1- 127I1- 53

  44. 2) 23 e-, 30 n, net charge = +3 # protons? 23 electrons, but charge of 3+ ie 3 more protons than electrons  p= 26  Atomic number = 26  element = Fe 56 3+ Fe 26

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