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Atomic Structure Atoms and their structure Mr. Bruder Evidence for Atoms Law of Conservation of Mass – Mass is not gained or lost in a chemical reaction. Proposed by Antoine Lavoisier in 1787. What would happen to the mass reading if the reaction was done without the balloon (an open system)? Figure 2.2 2-2 Evidence for Atoms Law of Definite Proportions – Proposed by Joseph Proust between 1797 and 1804 – A compound always has the same relative amounts of the elements that compose it. – For example, when water is broken down by electrolysis into oxygen and hydrogen, the mass ratio is always 8 to 1. Figure 1.2 2-3 Dalton’s Atomic Theory John Dalton (1766-1844) had four theories 1. All elements are composed of submicroscopic indivisible particles called atoms Atoms of the same element are identical. The atoms of anyone element are different from those of any other element Atoms of different elements can physically mix together or can chemically combine w/ one another in simple whole-number ratios to form compounds Chemical reactions occur when atoms are separated, joined, or rearranged. However, atoms of one element are never changed into atoms of another elements as a result of a chemical reaction 2. 3. 4. Atoms & Subatomic Particles Atom- smallest particle of an element that retains the properties of that element Proof of Atoms: STM Image of Gold The scanning tunneling microscope, invented in 1981, allows us to create images of matter at the atomic level. Figure 2.4 2-6 Electron J.J Thomson (1856-1940) – discovered the electron in 1897 Electron is the negative charged subatomic particle An electron carries exactly one unit of negative charge & its mass is 1/1840 the mass of a hydrogen atom Cathode Ray The Cathode Ray tubes pass electricity through a gas that is contained at a very low pressure Thomson’s Experiment Voltage source - + Thomson’s Experiment Voltage source - + Thomson’s Experiment Voltage source - + Thomson’s Experiment Voltage source + Passing an electric current makes a beam appear to move from the negative to the positive end Thomson’s Experiment Voltage source + Passing an electric current makes a beam appear to move from the negative to the positive end Thomson’s Experiment Voltage source + Passing an electric current makes a beam appear to move from the negative to the positive end Thomson’s Experiment Voltage source + Passing an electric current makes a beam appear to move from the negative to the positive end Thomson’s Experiment Voltage source By adding an electric field Thomson’s Experiment Voltage source + By adding an electric field Thomson’s Experiment Voltage source + By adding an electric field Thomson’s Experiment Voltage source + By adding an electric field Thomson’s Experiment Voltage source + By adding an electric field Thomson’s Experiment Voltage source + By adding an electric field Thomson’s Experiment Voltage source + By adding an electric field he found that the moving pieces were negative Thomson’s Atomic Model Thomson’s Atomic Model Thomson though electrons were like plums embedded in a positively charged “pudding”, so his model was called the “plum pudding” model Thomsom’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 Mass of Electron In 1909 Robert Millikan determined the mass of an electron with his Oil Drop Experiment He determined the mass to be 9.109 x 10-31 kg The oil drop apparatus Millikan’s Experiment Atomizer + - Oil Microscope Millikan’s Experiment Atomizer Oil droplets + - Oil Microscope Millikan’s Experiment X-rays X-rays give some drops a charge by knocking off electrons Millikan’s Experiment + Millikan’s Experiment - - + + They put an electric charge on the plates Millikan’s Experiment - - + + Some drops would hover Millikan’s Experiment - - - - - - - + + + + + + + + Millikan’s Experiment - - + + Measure the drop and find volume from 4/3πr3 Find mass from M = D x V Millikan’s Experiment - - + + From the mass of the drop and the charge on the plates, he calculated the charge on an electron Proton In 1886 Goldstein discovered the Proton Proton is a positively charged subatomic particle found in the nucleus of a atom Radioactivity Discovered by accident Bequerel Three types – alpha- helium nucleus (+2 charge, large mass) – beta- high speed electron – gamma- high energy light Ernest Rutherford Rutherford (1871-1937) proposed that all mass and all positive charges are in a small concentrated region at the center of the atom He used the Gold-Foil Experiment to prove his theory In 1911 he discovered the Nucleus Nucleus- central core of an atom, composed of protons and neutrons The nucleus is a positively charged region and it is surrounded by electrons which occupy most of the volume of the atom Rutherford’s Experiment Used uranium to produce alpha particles Aimed alpha particles at gold foil by drilling hole in lead block Since the mass is evenly distributed in gold atoms alpha particles should go straight through. Used gold foil because it could be made atoms thin Lead block Uranium Florescent Screen Gold Foil What he expected Because Because, he thought the mass was evenly distributed in the atom What he got How he explained it Atom is mostly empty Small dense, positive piece at center Alpha particles are deflected by it if they get close enough + + Nuclear Atom Viewed in Cross Section Copyright © Cengage Learning. All rights reserved 46 The Neutron Because the protons in the atom could account for only about half the mass of most atoms, scientists knew there was another heavy particle in the nucleus. Neutrons were proposed by Ernest Rutherford in 1907 (to account for a mass discrepancy in the nucleus) and discovered in 1932 by James Chadwick. The neutron has about the same mass as a proton but with no charge. 2-47 Bohr Model 1. 2. Bohr changed the Rutherford model and explained how the electrons travel. Bohr explained the following in his model: Electrons travel in definite orbits around the nucleus His model was patterned after the motion of the planets around the sun. It is often called the Planetary model. 48 Bohr’s Model Nucleus Electron Orbit Energy Levels The Quantum Mechanical Model Schrodinger Things that are very small behave differently from things big enough to see. The quantum mechanical model is a mathematical solution It is not like anything you can see. Modern View The atom is mostly empty space Two regions Nucleus- protons and neutrons Electron cloud- region where you have a chance of finding an electron Neutron James Chadwick (1891-1974) – discovered the neutron in 1932 Neutron is a subatomic particle with no charge but their mass is nearly equal to that of a proton Quark Protons & Neutrons can still be broken down into a smaller particle called the Quark The Quark is held together by Gluons Density and the Atom Since most of the particles went through, it was mostly empty. Because the pieces turned so much, the positive pieces were heavy. Small volume, big mass, big density This small dense positive area is the nucleus Subatomic Particles Protons and neutrons are located at the center of an atom called the nucleus. Electrons are dispersed around the nucleus. EOS Chapter 2: Atoms, Molecules, and Ions 55 Atomic Particles Particle Charge Mass (kg) Location Electron -1 9.109 x 10-31 Electron cloud Proton +1 1.673 x 10-27 Nucleus Neutron 0 1.675 x 10-27 Nucleus Subatomic particles Relative Actual mass (g) Name Symbol Charge mass Electron e- -1 1/1840 9.11 x 10-28 Proton p+ +1 1 1.67 x 10-24 Neutron n0 0 1 1.67 x 10-24 Symbols Contain the symbol of the element, the mass number and the atomic number Mass number Atomic number X Sub-atomic Particles Z - atomic number = number of protons determines type of atom A - mass number = number of protons + neutrons Number of protons = number of electrons if neutral Symbols A X Z 23 Na 11 Atomic Structure Symbols Proton = p+ Electron = e Neutron = n0 Atomic # - Subscript Mass # - Superscript 235 92 U Rules for Atomic Structure 1. 2. 3. Atomic # = # of Protons # of Protons = # of Electrons Mass # = # of Protons + # of Neutrons # of Neutrons = Mass # - # of Protons If you know the Mass # & Atomic # you know the composition of the element Symbols Find the –number of protons –number of neutrons –number of electrons –Atomic number –Mass Number 80 35 Br Symbols if an element has an atomic number of 34 and a mass number of 78 what is the –number of protons –number of neutrons –number of electrons –Complete symbol Symbols if an element has 78 electrons and 117 neutrons what is the –Atomic number –Mass number –number of protons –Complete symbol Example Element Atomic Mass # Protons Electro Neutro K # 19 ns 19 11 16 5 17 46 35 ns 19 23 35 Isotopes Isotope- atoms that have the same number of protons but different number of neutrons Since isotopes have a different number of neutrons the isotope has a different mass number. Isotopes are still chemically alike because they have the same number of protons and electrons Two Isotopes of Sodium Copyright © Cengage Learning. All rights reserved 68 Examples of Isotopes Heavy Water One ice cube is made with water that contains only the hydrogen-2 isotope. The other ice cube is composed of water with normal water which contains mostly hydrogen-1. Which is which? Figure 2.13 2-70 Isotopes Isotopes are atoms of the same element with different masses. Isotopes have different numbers of neutrons. 11 C 6 12 C 6 13 C 6 14 C 6 © 2009, Prentice-Hall, Inc. Naming Isotopes Put the mass number after the name of the element carbon- 12 carbon -14 uranium-235 Electrical Charges Electrical charges are carried by particles of matter Atoms have no net electrical charges Given the number of negative charges combines with the number of positive charges = Electrically Neutral All elements are Electrically Neutral Atomic Mass vs. Atomic Weight Atomic Mass is for a single element Most elements are Isotopes How do we find their mass? We use Atomic Weight Atomic Masses An atomic mass unit (amu) is defined as exactly one-twelfth the mass of a carbon-12 atom 1 u = 1.66054 × 10–24 g The atomic mass of an element is the weighted average of the masses of the naturally occurring isotopes of that element EOS Chapter 2: Atoms, Molecules, and Ions 75 Measuring Atomic Mass Unit is the Atomic Mass Unit (amu) One twelfth the mass of a carbon-12 atom Each isotope has its own atomic mass. We need the average from the percent abundance Each isotope of an element has fixed mass and a natural % abundance You need both of these values to find the Atomic Weight Calculating Atomic Weight 1. 2. 3. Cl-35 34.969amu and 75.77% abundance Cl-37 36.966amu and 24.23% abundance To solve for Cl-35 AMU x Abundance 34.969 x .7577 = 26.496 You solve for Cl-37 Atomic Weight Cont. 1. 2. 3. Cl-37 AMU x Abundance 36.966 x .2423 = 8.957 Now you combine your two answers 26.496 + 8.957= 35.453 Look at Cl on the table. What is the Atomic Weight? Example Calculate the atomic weight of copper. Copper has two isotopes. One has 69.1% and has a mass of 62.93 amu. The other has a mass of 64.93 amu. What is the atomic weight??? Atomic Weight & Decimals Atomic Weight- of an element is a weighted average mass of the atoms in a naturally occurring sample of an element Atomic Weights use decimal points because it is an average of an element