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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 composed 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 neutral • 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 • Rounded Atomic 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. Ex: Calculate the number of neutrons for 23892U.

  25. Ex: Classify the following atoms2145X 2345X 2045X.

  26. Ex: Classify the following atoms19679X 19580X 19578X.

  27. 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

  28. Atomic Mass • Percent Abundance – the number of desired particles in 100 total particles of sample • Allows for comparison to any sample set • Relative Abundance – the number of desired particles in the sample used • Specific to the sample used; not useful in comparison • Convert % abundance to a decimal = relative abundance Desired particles % Ab = x 100% Total particles in sample

  29. 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

  30. 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.

  31. Ex: If there are 100 black beans, 27 pinto beans, and 173 lima beans in the container, what is the percent abundance of the container by bean? Relative abundance?

  32. 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%).

  33. 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%).

  34. 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%).

  35. 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.

  36. 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

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

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

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

  40. 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?

  41. Nuclear vs. Chemical Reactions Nuclear Reactions Chemical Reactions • Deals with nucleus • Can end up with new atoms/elements • Mass is not strictly conserved • Mass Defect • E = mc2 • Deals with electrons • Atoms/elements remain unchanged – rearranged • Mass is strictly conserved

  42. 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

  43. 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

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