7. Describe the structure of a typical atom. Identify where each subatomic particle is located. A typical atom consists of a central, small, dense nucleus containing protons and neutrons. The nucleus is surrounded by a cloud of negatively charged electrons.

Direct vs. Indirect Evidence Direct Evidence: You can use your senses to know what is happening. You can see it! Proves the existence of a particular fact without any assumptions. Indirect Evidence: You can not see what is happening. You must use reasoning powers to make conclusions. [Circumstantial evidence]

9. Evaluate the experiments that led to the conclusion that electrons are negatively charged particles found in all matter. The deflection toward positively charged plates demonstrated the negatively charged nature of electrons; the fact that changing the type of electrode or the type of gas used in the cathode-ray tube did not affect the ray -led to the conclusion that electrons are present in all matter.

Thomson’s Experiment Voltage source + - Passing an electric current makes a beam appear to move from the negative to the positive end By adding an electric field he found that the moving pieces were negative

8. Compare and contrast Thomson’s plum pudding atomic model with Rutherford’s nuclear atomic model. Thomson’s plum pudding model describes atoms as spherical particles with uniformly distributed positive charge in which individual, negatively charged electrons are located in fixed positions. In contrast, Rutherford’s model states that an atom is mostly empty space, with a small, dense, central nucleus containing all of an atoms positive charge and most of its mass. The negatively charged electrons move through the empty space and are held in the atom by their attraction to the positively charge nucleus.

Thomson’s Model Found the electron Couldn’t find positive (for a while) Said the atom was like plum pudding A bunch of positive stuff, with the electrons able to be removed

Rutherford’s Gold Foil experiment Ernest Rutherford English physicist. (1910) Believed in the plum pudding model of the atom. Wanted to see how big atoms are Used radioactivity Alpha particles - positively charged pieces given off by uranium Shot them at gold foil which can be made a few atoms thick

Florescent Screen Lead block Uranium Gold Foil

What he expected: He expected the alpha particles to pass through without changing direction very much Because he thought the positive charges were spread out evenly. Alone they were not enough to stop the alpha particles

What he got:

How he explained it + Atom is mostly empty space Small dense, positive center (nucleus) Alpha particles are deflected by it if they get close enough +

+

Modern View The atom is mostly empty space Two regions: Nucleus- protons and neutrons Electron cloud- region where you might find an electron

10. Compare the relative charge and mass of each of the subatomic particles Relative Mass Electron -1 1/1840 (~0) Proton +1 1 Neutron

- by the atomic number ( # of protons) The # of protons 20. Explain how the type of an atom is defined - by the atomic number ( # of protons) 21. Recall Which subatomic particle identifies an atom as that of a particular element? The # of protons

Average Mass Activity    Find the average mass of all the spheres in the tray.   You may: -Make no more than 3 mass measurements. -Mass only 1 sphere at a time Record all measurements and show your work. Explain what you did and all assumptions.

22. Explain how the existence of isotopes is related to the fact that atomic masses are not whole numbers Atomic masses are not whole numbers because they represent weighted averages of the masses of the isotopes of an element.

Isotopic Notation: Symbolic Written Form p + n gain/loss of e- Isotopes – atoms of the same element (same atomic #), but with different # neutrons (so different mass #). Isotopic Notation: Symbolic Written Form p + n gain/loss of e- atomic mass charge Symbol Atomic number Element – atomic mass p ___ n

#n = (mass # - atomic #) = 14 - 6 = 8 Example: Find the number of protons, neutrons, and electrons 14 C 6 p + n Carbon-14 p # p = atomic # = 6 #n = (mass # - atomic #) = 14 - 6 = 8 # e- = # p = atomic # = 6

Practice Problem  Determine the number of protons, neutrons, & electrons in each of the isotopes given: 16 2- 80 235 64 2+ O Br U Cu 8 35 92 29   8 p+ 35 p+ 92 p+ 29 p+ 8 n 45 n 143 n 35 n 10 e- 35 e- 92 e- 27 e-

23. Calculate Copper has two isotopes: Cu-63 (abundance=69.2%, mass=62.930 amu) and Cu-65 (abundance= 30.8%, mass=64.928 amu) Calculate the atomic mass of copper. .692 x 62.930 = 43.5 .308 x 64.928 = 20.0 63.5 amu

24. Three magnesium isotopes have atomic masses and relative abundances of 23.985 amu (78.99%), 24.986 amu (10.00%), and 25.982 (11.01%). Calculate the atomic mass of magnesium 23.985 x .7899 = 18.95 24.986 x .1000 = 2.499 25.982 x .1101 = 2.861 24.31 amu

Silver is found in two isotopes with atomic Practice Problem  Silver is found in two isotopes with atomic masses 106.9041 amu and 108.9047 amu, respectively. The first isotope represents 51.82% and the second 48.18%. Determine the average atomic mass of silver. 106.9041(0.5182) + 108.9047(0.4818) = 107.87 amu 55.40 + 52.47

Do Problems pgs 128-129 36, 39, 44, 46, 58-62, 64 & 65