Atomic Structure Early Theories Democritus: 4 B.C.: atom Dalton: atoms cannot Thomson: Cathode Ray Tubes Rutherford:
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1 Atomic Structure n a well-substantiated explanation of some aspect of the natural world; n an organized system of accepted knowledge that applies in a variety of circumstances to explain a specific set of phenomena; n "theories can incorporate facts and laws and tested hypotheses" Early Theories Democritus: 4 B.C.: atom Believed there were 4 elements: Fire, Air, Water, Earth Dalton: >All elements composed of tiny particles called atoms >Atoms of same element are identical; atoms of different elements are different >Atoms of different elements can physically mix together or chemically combine to form compounds >Chemical reactions cannot change atoms of one type of element to another Thomson: >discovered electrons in 1897 >used a cathode ray tube >the ray produced was deflected by an electrical field (showed that atoms had particles with (-) charge) Cathode Ray Tubes n A cathode ray tube or CRT is a specialized vacuum tube in which images are produced when an electron beam strikes a phosphorescent surface. n TVs, PCs, ATMs, video games, video cameras, and monitors all contain cathode-ray tubes. n Displays millions of colors. Rutherford: >Gold Foil Experiment >Discovered the nucleus Rutherford s Gold Foil Experiment Experiment Shot positively charged alpha particles at gold foil Results 1) Most particles passed through the foil 2) A few were deflected 3) small, dense, positively charged core (nucleus) 4) the rest of the atom is empty space
2 Bohr à planetary model n electrons arranged in concentric circular patterns n paths or orbits around nucleus (energy level) Wave-Mechanical Model à Electron Cloud Model q based on the ideas that orbitals are the area of highest probability where an electron will be found. q Orbitals have a variety of shapes and names (s, p, d, f) Summary- Atomic Models Dalton s Cannonball Thomson s Plum Pudding Rutherford s Nuclear Model-most of atom s mass is small positively charged located in the nucleus Bohr s Planetary Model-electrons travel in defined orbits Sub Atomic Particles 1 amu = 1/12 th mass of a carbon-12 atom Name Symbol Charge Mass Proton (located in nucleus à nucleon) Neutron (located in nucleus à nucleon) Electron (located outside the nucleus in orbitals) p amu n amu e /1836 amu Atomic Number n Equal to the number of protons n Every element has its own atomic number n See Periodic Table 6C Mass Number n Equal to the sum of the protons and the neutrons (whole number) n Can be written as carbon C
3 Number of Protons =Atomic number # of electrons = in a neutral atom, it is equal to the number of protons # of neutrons à if protons + neutrons = mass then, # of neutrons = mass # - # protons 1. What is the nuclear charge of a sulfur atom? 2. If an atom has 5 protons and 6 neutrons how many electrons does it have? What would the mass of this atom be? 3. What must all atoms of the same element have in common? 4. List the relative mass, charge and location of all three subatomic particles 5. How many protons, neutrons and electrons does an atom of Lithium-7 have? 6. What conclusions about the atom where discovered as a result of the Gold Foil Experiment? Ions: n Defined as charged particles n Ions are formed when the number of electrons changes. n If a (+) ion is formed, electrons are lost (called cations). If a (-) ion is formed, electrons are gained (called anions). Example; n Ca 2+ A Ca atom has 20 protons and 20 electrons. A Ca 2+ ion has lost two electrons to have 18. n Cl - A Cl atom has 17 protons and 17 electrons. A Cl - ion has gained one electron to have 18 Isotopes n Definition: Atoms that have the same atomic number (same # of protons) but a different mass number (different # of neutrons) Isotopic Symbols n 2 Isotopes will have the same atomic # (bottom) and a different mass # (top) Mass # X Atomic # Write the isotopic symbol for: n Carbon-14 (write a symbol for a different isotope of carbon) 14 C 6 Hydrogen has three Isotopes H-1,H-2,H-3
4 Why is atomic mass not a whole number? n The atomic mass on the periodic table is a weighted average of the isotopes of the elements. n The weighted atomic mass takes into account the relative abundances of all the naturally occurring isotopes. How do you calculate a weighted average? n To calculate the weighted average you convert each percentage to a decimal by moving it 2 places left. Multiply the decimal by the mass for each isotope and add them all up. n Or you can multiply the percent abundance (without moving the decimal) by the mass for each isotope add them all up and divide by 100 Example of a general weighted average n Your grade in chemistry n 70% exams 85 n 10% quizzes 100 n 10% labs 95 n 10% HW/CW 80 n (0.70)85 + (0.10)100 + (0.10)95 + (0.10)80 = 87 Example 1: n Determine weighted atomic mass n Boron % amu n Boron % amu n (0.1978) (.8022) = amu n How many total electrons does an Al +3 ion have? n If a neutral atom has 10 neutrons and 8 electrons how many protons does it have? n How does an atom of Lithium-7 differ from an atom of Beryllium-9 n Compare a Na +1 ion to a Na atom? n How are 14 N, 15 N and 16 N different and the same? Bohr models How do electrons orbit the nucleus? *Each principal energy level is a fixed distance from the nucleus *can hold a specific number of electrons *has a definite amount of energy The greater the distance from the nucleus the greater the energy of the electrons in it. The orbits are called principal energy levels or shells. Energy levels or shells Energy Level # of electrons 1 2 Closest to nucleus/less energy Furthest from nucleus/more energy
5 Examples: Carbon 1 st shell 2e -, 2 nd shell 4e - Sodium 1 st shell 2e -,2 nd shell 8e -, 3rd shell 1e - Electron configurations indicates number and location of electrons, located on periodic table on key Lewis Dot Diagrams Valence shell: outer most shell of an atom that contains electrons Valence electrons: electrons that occupy the valence shell (last number in electron configuration) Electron dot diagrams or Lewis dot diagrams: show only the valence shell of the atom Ex: Lewis dot for nitrogen: Practice: O,F,C Ne, I, K Ions For ions: remember that ions have gained or lost electrons. Use periodic table to find charge of ion (see table) For dot diagrams of Ions (+)ions à indicate charge no dots around the symbol (-)ions à use brackets, charge, and always 8 dots around the symbol Dot Diagrams for Ions Ground State vs. Excited State n When all electrons in an atom occupy the lowest available orbitals, it is said to be in the ground state. n When electron(s) absorb energy, they have the ability to jump to higher energy levels. n The excited state is when electrons have absorbed energy and no longer occupy the lowest available energy levels. Electron Configuration examples Ground State for Na Excited state(many possibilities) for Na (2-7-2), (2-6-3),(2-5-4)
6 Absorption n When an electron jumps to a higher energy level it absorbs energy. n The excited state is a temporary state. n Examples; electrons go from 2 PE level to PE level 3 or 4, etc Emission n The electron then falls back down to the ground state, emitting energy. n The energy is in the form of light. n This radiant energy has a characteristic color and wavelength that can be determined. n Every electron transition produces a specific wavelength of light and all transitions for an element blend together. n This light can be separated through a prism into its various wavelength components. n Every element has its own unique bright line spectrum that can be used to help identify the presence of that element. n Ex: elements in a star, forensic analysis, flame tests, spectroscopy Light and Atomic Spectra (bright line spectra) Electromagnetic spectrum consists of light that exists as waves. Atomic emission spectra produce narrow lines of color called bright line spectra. n Each line corresponds to an exact wavelength.
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