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Ch. 3: Atomic Structure The Theory of the Atom • ________________, Democritus a famous Greek teacher who lived in the 4th Century B.C., first suggested the idea of the atom. • ________ John __________ Dalton came up with his atomic theory based on the results of his experiments. (See p. 56) The Atom • element The smallest particle of an ________________ is an atom. • subatomic particles. The atom is made up of three ________________ 1897 by J. J. Thomson by using a cathode (1) The electron was discovered in _______ (−) charge. It’s mass is much smaller than the ray tube. The electron has a _______ ignored other 2 subatomic particles, therefore it’s mass is usually ______________. Cathode Ray Tube (+) charge, and it was discovered in (2) The proton has a ______ 1886 _________ by E. Goldstein. (3) The neutron does not have a charge. In other words, it is neutral It was discovered in 1932 ________. ____ by James Chadwick. The neutron has about the same _________ mass as the proton. visible matter • These three particles make up all the ____________________ in the Universe! • There are other particles such as neutinos, positrons, and quarks, but are typically left for 2nd year chemistry courses. Nuclear Atomic Structure • The atom is made up of 2 parts/sections: (1) The ______________ --- (in the center of the atom) nucleus electron _________ cloud --- (surrounds the nucleus) (2) The ____________ (p+ & n0) e− cloud The Nucleus • Discovered by Ernest ________________ in ________. Rutherford 1911 • He shot a beam of positively charged “alpha particles”, which are ___________ nuclei, at a thin sheet of ______ helium gold_____. foil • 99.9% of the particles went right on through to the ______________. detector • • Some were slightly deflected. Some even ____________ ________ bounced back towards the source! • This would be like shooting a cannon ball at a piece of tissue paper and having it bounce off. Rutherford’s Experiment Conclusions about the Nucleus empty ___________. space (1) Most of the atom is more or less _________ tiny (2) The nucleus is very _________. (Stadium Analogy) dense (3) The nucleus is very ___________. (Large Mass ÷ Small Volume) (4) The nucleus is ______________ positively charged. Counting Subatomic Particles in an Atom protons in the nucleus. • The atomic # of an element equals the number of ____________ protons neutrons • The mass # of an element equals the sum of the _____________ and _____________ in the nucleus. electrons • In a neutral atom, the # of protons = # of ______________. subtract mass • To calculate the # of neutrons in the nucleus, ______________ the ___________ # atomic #. from the __________ Practice Problems (1) Find the # of e-, p+ and n0 for sodium. (mass # = 23) Atomic # = 11 = # e- = # p+ 2) # neutrons = 23-11 = 12 Find the # of e-, p+ and n0 for uranium. (mass # = 238) Atomic # = 92 = # e- = # p+ # neutrons = 238-92 = 146 3) What is the atomic # and mass # for the following atom? # e- = 15; # n0 = 16 Atomic # = 15 = # e- = # p+ Mass # = p+ + n0 = 15+16 =31 The element is phosphorus! Isotopes protons • An isotope refers to atoms that have the same # of ___________, but they have a neutrons different # of ___________. mass • Because of this, they have different _________ #’s (or simply, different masses ___________.) • Isotopes are the same element, but the atoms weigh a different amount because neutrons of the # of ______________. Examples---> (1) Carbon-12 & Carbon-13 (2) Chlorine-35 & Chlorine-37 (The # shown after the name is the mass #.) atomic • For each example, the elements have identical ___________ #’s, (# of p+) but mass #’s, (# of n0). different _________ • Another way to write the isotopes in shorthand is as follows: 12 C 6 35 17 Cl The top number is the ________ mass #, and the bottom # is the __________ atomic number. subtracting the #’s! Calculating the # n0 can be found by _____________ Figure 3.10: Two isotopes of sodium. More Practice Problems (1) Find the # e-, p+ and n0 for Xe-131. Atomic # = 54 = p+ = e− 2) Find the # e-, p+ and n0 for n0 = 131 − 54 = 77 63 29 Cu Atomic # = 29 = p+ = e− n0 = 63 − 29 = 34 3) Write a shorthand way to represent the following isotope: # e- = 1 # n0 = 0 # p+ = 1 Atomic # = p+ = e− = 1 H-1 or mass # = n0 + p+ = 1+ 0 = 1 1 1 H Ions • An atom can gain or lose electrons to become electrically charged. • Cation = (___) + charged atom created by ___________ losing e-’s. – Cations are ______________ than the original atom. smaller – _____________ generally form cations. Metals • Anion = (___) e-’s. − charged atom created by _____________ gaining – larger Anions are ____________ than the original atom. – _______________ generally form anions. Nonmetals Practice Problems: Count the # of protons & electrons in each ion. a) Mg+2 12 10 p+ = _____ e− = ______ b) F−1 9 10 p+ = _____ e− = ______ Atomic Mass 12 • Based on the relative mass of Carbon-12 which is exactly _______. 1 amu 1 atomic mass unit (amu) 1 n0 ≈ __ • 1 p+ ≈ __ 0 amu 1e- ≈ __ • The atomic masses listed in the Periodic Table are a “weighted average” of all the isotopes of the element. Weighted Average Practice Problems: (1) Mr. Turkowski’s Algebra 1 semester grades are calculated using a weighted average of three category scores: Major Grades= 60% of your grade Minor Grades= 30% of your grade Semester Exam=10% of your grade • If a student had the following scores, what would they receive for the semester? Major= 80 (B − ) Minor= 60 (D −) Semester Exam=65 (D) Weighted Average Step (1): Multiply each score by the % that it is weighted. Step (2): Add these products up, and that is the weighted average! 60% x 80 = 48.0 30% x 60 = 18.0 10% x 65 = + 6.5 Add them up!! 72.5 (C−) A “normal average” would be calculated by simply adding the raw scores together and dividing by 3… 80 + 60 + 65 = 205 ÷ 3 = 68.3 = D Average Atomic Mass Practice Problems: (2) In chemistry, chlorine has 2 isotopes: Cl-35 (75.8% abundance) Cl-37 (24.23 % abundance) What is the weighted average atomic mass of chlorine? 35 x 0.758 = 26.53 37 x 0.2423 = + 8.9651 Add them up!!! 35.4951 amu (3) Oxygen has 3 isotopes: O-16 (99.76%) O-17 (0.037%) Estimate oxygen’s average atomic mass. Barely over 16.0 amu O-18 (0.2%) Average Atomic Mass (4) Copper has an average atomic mass of 63.546 g/mole. It contains only two natural isotopes, which are Cu-63, with an isotope mass of 62.940 and Cu-65 with an isotope mass of 64.928. What are the percent of the two isotopes in naturally occurring copper?