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Unit 3 Atomic Structure

Unit 3 Atomic Structure. Chemistry I Mr. Patel SWHS. Topic Outline. Learn Major Ions Defining the Atom (4.1) Subatomic Particles (4.2) Atomic Structure (4.2) Ions and Isotopes (4.3) Nuclear Chemistry (25.1). Defining the Atom.

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Unit 3 Atomic Structure

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  1. Unit 3Atomic Structure Chemistry I Mr. Patel SWHS

  2. Topic Outline • Learn Major Ions • Defining the Atom (4.1) • Subatomic Particles (4.2) • Atomic Structure (4.2) • Ions and Isotopes (4.3) • Nuclear Chemistry (25.1)

  3. Defining the Atom • Atom – the smallest particle of an element that retains its identity • Can not see with naked eye • Nanoscale (10-9 m) • Seen with scanning tunneling electronmicroscope

  4. Democritus • Democritus was a Greek to first come up with idea of an atom. • His belief: atoms were indivisible and indestructible. = WRONG! • Atom comes from “atmos” - indivisible

  5. Dalton’s Atomic Theory • 2000 yrs later, John Dalton used scientific method to transform Democritus’s idea into a scientific theory • Dalton put his conclusions together into his Atomic Theory (4 parts)

  6. Dalton’s Atomic Theory • All elements are composted of tiny, indivisible particles called atoms.

  7. Dalton’s Atomic Theory • Atoms of the same element are identical. Atoms of different elements are different

  8. Dalton’s Atomic Theory • Atoms of different elements can physically mix or chemically combine in whole number ratios.

  9. Dalton’s Atomic Theory • Chemical reactions occur when atoms are separated, joined, or rearranged. Atoms of one element can never be changed into atoms of another element due to a chemical reaction.

  10. The Electron • Particle with negative charge • Discovered by J.J. Thomson • Used cathode ray (electron) beam and a magnet/charged plate. • Millikan found the charge and mass

  11. The Proton and Neutron • An atom is electrically nuetral • If there is a negative particle then there must be positive particle • Proton – particle with positive charge • Chadwick discovered neutron – neutral charge

  12. Thomson’s Atomic Model • Electrons distributed in a sea of positive charge • Plum Pudding Model

  13. Rutherford’s Atomic Model • Performed Gold-Foil Experiment • Beam of Alpha particles with positive charge shot at thin piece of gold foil • Alpha particles should have easily passed through with slight deflection due to positive charge spread throughout. • Results: Most particles went straight through with no deflection. Some were deflected at large angles.

  14. Rutherford’s Atomic Model • The nucleusis the central part of the atom containing protons and neutrons • Positive charge • Most of the mass • Electrons are located outside the nucleus • Negative charge • Most of the volume

  15. Atomic Number • An element is defined only by the number of protons it contains • Atomic Number – number of protons • Number of protons = number of electron • For a neutral element

  16. Identify the number of Protons • Zinc (Zn) • Iron (Fe) • Carbon (C) • Uranium (U) • 30 • 26 • 6 • 92

  17. Mass Number • Nucleus contains most of the mass • Mass Number – total protons and neutrons Number of neutron = Mass # – Atomic #

  18. Identify # of Subatomic Particles • Lithium (MN = 7) • Nitrogen(MN = 14) • Fluorine(MN = 19)**MN = Mass Number • 3 p+ , 3 e-, 4 n0 • 7 p+ , 7 e-, 7 n0 • 9 p+ , 9 e-, 10 n0

  19. Differences in Particle Number • Different element: different number of protons • Ions – same number of proton, different number of electrons • Isotope – same number of proton, different number of neutrons • Different Mass Numbers

  20. Two Notations for Atoms • Nuclear Notation • Write the element symbol • On left side, superscript = Mass Number • On left side, subscript = Atomic Number • Isotope –Hyphen Notation • Write full name of element • On right side, put a dash • On right side put Mass Number after dash Hydrogen - 3

  21. Ex: Three isotopes of oxygen are oxygen-16, oxygen-17, and oxygen-18. Write the nuclear symbol for each.

  22. Ex: Three isotopes of chromium are chromium-50, chromium-52, and chromium-53. How many neutrons are in each isotope?

  23. Ex: Calculate the number of neutrons for 9942Mo.

  24. Atomic Mass • Atomic Mass Unit (amu) – one-twelfth of the mass of the carbon-12 atom • Different isotopes have different amu (mass) and abundance (percentage of total) • Atomic Mass – weighted average mass of the naturally occurring atoms. • Isotope Mass • Isotope Abundance

  25. Atomic Mass • Because abundance is considered, the most abundant isotope is typically the one with a mass number closest to the atomic mass. • Example, Boron occurs as Boron-10 and Boron-11. Periodic Table tells us Born has atomic mass of 10.81 amu. • Boron-11 must be more abundant

  26. Calculating Atomic Mass • Convert the Percent Abundance to Relative Abundance (divide by 100) • Multiple atomic mass of each isotope by its relative abundance • Add the product (from step above) of each isotope to get overall atomic mass.

  27. Ex: Calculate the atomic mass for bromine. The two isotopes of bromine have atomic masses and percent abundances of 72.92 amu (50.69%) and 80.92 amu (49.31%).

  28. Ex: Calculate the atomic mass for X. The four isotopes of X have atomic masses and percent abundances of 204 amu (1.4%), 206 amu (24.1%), 207 amu (22.1%), and 208 amu (52.4%).

  29. Ex: Calculate the atomic mass for H. The three isotopes of H have atomic masses and percent abundances of 27 amu(85%), 26 amu (10%), and 28 amu (5%).

  30. Nuclear Radiation • Radioactivity – nucleus emits particles and rays (radiation) • Radioisotope – a nucleus that undergoes radioactive decay to become more stable • An unstable nucleus releases energy through radioactive decay.

  31. Nuclear Radiation • Nuclear force – the force that holds nuclear particles together • Very strong at close distances • Of all nuclei known, only a fraction are stable • Depends on proton to neutron ratio • This region of stable nuclei called band of stability

  32. Half Life • Half Life – the time required for one-half the sample to decay • Can be very short or very long

  33. Ex: The original amount of sample was 100 g. The amount currently remaining is 25 g. How many half-lives has gone by?

  34. Ex: The original amount of sample was 100 g. The amount currently remaining is 25 g after 30 minutes. What is the half life?

  35. Ex: The original amount of sample was 100 g. The amount currently remaining is 6.25 g. The half life is 50 years. How much time has passed?

  36. Types of Radiation • Alpha Radiation (Helium Atom) • Low penetrating power • Paper shielding • Beta Radiation (Electron) • Moderate penetrating power • Metal foil shielding • Gamma Radiation (Pure energy) • Very high penetrating power • Lead/concrete shielding

  37. Nuclear Decay Equations • Transmutation – conversion from one element to another through a nuclear reaction • Only occur by radioactive decay • Only when nucleus bombarded with a particle • Emissions – given off • Alpha Emission, Beta Emission, Positron Emission • Positron = beta particle with a positive charge • Captures – taken in • Electron Capture

  38. Ex: Show a Beta Emission of Copper-66.

  39. Ex: Show an Electron Capture of Nickel-59.

  40. Ex: Show a Positron Emission of Boron-8.

  41. Ex: Show an Alpha Emission of Thorium-232.

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