Chapter 3 1 The Atom w Summarize the
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Chapter 3 -1: The Atom w Summarize the five essential points of Dalton’s atomic theory w Explain the relationship between Dalton’s Atomic Theory and the laws of conservation of mass and definite composition w Explain the law of multiple proportions.
Early Ideas of the Atom Democritus w Greek Philosopher w 460 -370 B. C. w Stated – Matter could be divided into smaller & smaller particles until it could no longer be divided. w Called these Particles – Atomos (indivisible) w Not based on any physical evidence, only thought
Dalton’s Atomic Theory Late 1700’s – John Dalton – School Teacher(england) Dalton’s Atomic Theory – Summarized 1. All matter is composed of extremely small particles called atoms. 2. Atoms of a given element are identical in size, mass, and other properties. (Isotopes - atoms same element with different mass. ) 3. Atoms cannot be subdivide, created, or destroyed. 4. Atoms of different elements can combine in simple, wholenumber ratios to form chemical compounds. 5. In chemical reactions, atoms are combined, separated, or rearranged.
Law of Conservation Atoms and the conservation of mass If atoms are indivisible w mass must be conserved A + B AB + 1 a. u. + 3 a. u. 4 a. u. Law of Definite Composition Law of Constant Composition A compound always contains the same elements in the same proportions by mass. The mass ratio of A to B will always be the same. In this case 1: 3 or 25% to 75%
Law of Multiple Proportions w If two or more different compounds are composed of the same two elements, then the ratio of the masses of the second element combined with a certain mass of the first element is always a ratio of small whole numbers. w CO 2; 1 g to 2. 66 g CO; 1 g to 1. 33 g w Ratio of oxygen would 2: 1.
Law of Definite Proportions Calculations w Hydrogen and Oxygen have a mass ratio of 1: 16. What is the mass of oxygen needed to form with 14 g of Hydrogen? What is the total mass of this compound?
Law of Definite Proportions Calculations w Magnesium and Oxygen have a mass ratio of 3: 2. What is the mass of oxygen needed to form with 20 g of Magnesium? What is the total mass of this compound?
3. 2 Discovering Atomic Structure w Particle arrangement of the atom. n n n Protons (+) Electrons (-) Neutrons (0) w Michael Faraday n Made the connection between the atom and electricity.
J. J Thomson w Used a Cathode Ray Tube (CRT) n Discovered electrons Charge and mass l Electric current passed through low pressure gas in a glass tube. l
Robert Millikan w Millikan’s Oil-Drop Experiment Proved the mass and charge of all electrons are identical. n Used an electric plate created a resistant force acting on falling particles n Voltage Controlled n By varying the charge on different drops, he noticed that the charge was always a multiple of -1. 6 x 10 -19 C. n
Ernest Rutherford w Used the Gold Foil Experiment n Marsden and Gieger proved the existence of the nucleus w Bombarded thin sheets of metal with (+) charged particles. w Results of the Particle deflections n Most traveled through n Some Deflected in Both Directions n Very Few Deflected Back
Rutherford vs. Thomson
Rutherford results w Atom n Nucleus Central part of the atom. l Very small and very massive = Very dense. l w All the mass of an atom. l n Positively charged, location of protons. Electrons Surround the nucleus, gives size to the atom. l Negatively charged, atom itself is electrically neutral. l
E. Goldstein w Used canal rays in a cathode ray tube to prove the existence of protons. n Positively charged particles that had significant mass moved towards the cathode.
Sir James Chadwick w Using kinematics, Chadwick was able to determine the velocity of the protons. Then through conservation of momentum techniques, he was able to determine that the mass of the neutral radiation was almost exactly the same as that of a proton. . n Neutrons, were neutrally charged particles found in the nucleus to add mass to the atom and to act as “nuclear glue. ”
Summary of the Atom w Nucleus - very dense, (+)charged, center of an atom. n Protons l l n Positively charged particle found in the nucleus Gives mass to the atom. Neutrons l l Neutrally charged particle found in the nucleus. Gives mass to the atom and acts as “nuclear glue” w Electrons n Negatively charged particles that give the atom its size.
3. 3 Modern Atomic Theory Properties of Subatomic Particles Particle Relative Mass Symbols Charge Number Relative Mass Actual Mass Electron e- -1 0 0. 00055 u 9. 11 x 10 -28 g Proton p+ +1 1 1. 00728 u 1. 67 x 10 -24 g Neutron n 0 0 1 1. 00866 u 1. 68 x 10 -24 g
Atomic Mass Units (amu) w the equivalent mass of an atom. n n n Protons = 1 amu Neutrons = 1 amu Electrons = 0 amu
Subatomic calculations w Atomic Number (Z) –the number of protons within an atom. (Atomic # = p+) n n n Every element has its own unique atomic number. p+ = e- , in an atom Mass # = p+ + n 0
Isotope Abbreviations w Atomic Symbol Notation: n X = atomic symbol w. Mass Number form: n name of element – mass#
Isotope Form Practice w Write the element uranium with a mass # of 235 in atomic and mass form: n Mass # = n Atomic form =
Subatomic Particle Practice w How many protons, neutrons and electrons do the following contain? Carbon – 13 n. Carbon lp+ = le- = lno = n Gold p+ = e- = n 0 =
Symbolic Mass Form Atomic Mass p+ # # n 0 e- 33 30 Zinc-63 30 Osmium-191 76 115 76 Tungstun-180 74 106 74 Barium-139 56 83 56 Manganese-53 25 28 25
Isotopes w Atoms of the same element with different mass, due to the number of neutrons. n n n The number of protons determines the identity of an atom !!!!!!! The number of neutrons can be different for the same type of element. Hydrogen
Hydrogen Isotope Names w 3 isotopes of hydrogen n Protium, Deuterium, Tritium, Hydrogen – 1 Hydrogen – 2 Hydrogen – 3
Relative Atomic Masses w Although we know actual masses of atoms it is more practical to use their relative masses w The arbitrary point – Carbon-12 w 1 atomic unit (a. u. ) = Exactly 1/12 of C-12 w All other atoms, weighed based upon Carbon-12
Atomic Mass Average mass of all naturally occurring isotopes of an element. w Calculation Atomic mass = (mass of each isotope)(% abundance)/100 l Add the answers to each of the isotopes to get the average mass. l
Modern Atomic Theory w Determine the atomic mass of Oxygen. Isotope Mass (amu) % abundance Oxygen – 16 15. 994915 99. 762 Oxygen – 17 16. 999131 . 038 Oxygen – 18 17. 999160 . 200
Modern Atomic Theory
Modern Atomic Theory w Determine the atomic mass of Uranium. Isotope Mass (amu) % abundance Uranium-235 235. 043924 . 720 Uranium - 238. 050784 99. 280 Uranium =
Relating Mass to Particles w The Mole w The number of atoms of an element equal to the number of atoms in exactly 12. 0 g of carbon-12. w Referred as a counting number. w 1 mole of any element = 6. 02 x 10 23 atoms w Avogadro’s number, there are exactly 6. 02 x 1023 atoms of any element in 1 mole of that element. 1 mol C = 1 mol H = 1 mol S =
Molar Mass w Mass in grams of 1 mole of any substance. n Equivalent to the formula mass of a compound and to the atomic mass of an element. 1 mol of S = 32 g S l 1 mol of C = 12 g C l
Mole Conversions w 1 mol = Formula mass (or Atomic Mass) n Or the Molar Mass w 1 mol = 6. 02 x 10 23 particles w The mole is the central unit in converting the amount of substances in chemistry.
Mass – Mole Conversions w Use 1 mol = Formula Mass conversions. w Mass to Moles Conversion n n How many moles are in 250. g of Na? Moles to Mass Conversion – How many grams are in. 55 moles of C?
Practice Mass-Mole w Determine the number of moles in the following substances: 1) 2) n 6. 60 g N 100. g Fe Determine the mass for the following substances: 3) 6. 25 mol Cu 4). 650 mol P
Mole – Particle Conversion w Use 1 mol = 6. 02 x 10 23 particles w Particle to Mole Conversion n n How many moles are equivalent to 550. atoms of S? Mole to Particle Conversion – How many atoms are in. 525 moles of Ca?
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