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General Chemistry 1: Week 2

by: David Ruin

General Chemistry 1: Week 2 General Chemistry 1113

David Ruin

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About this Document

Notes on Chapter 2 of the General Chemistry textbook or Week 2 in class.
General Chemistry
Class Notes
General Chemistry




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This 9 page Class Notes was uploaded by David Ruin on Wednesday August 31, 2016. The Class Notes belongs to General Chemistry 1113 at University of Colorado at Boulder taught by in Fall 2016. Since its upload, it has received 12 views. For similar materials see General Chemistry in Chemistry at University of Colorado at Boulder.

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Date Created: 08/31/16
Wednesday, January 20, 2016 Elements, Matter, Dalton’s Atomic Theory Chapter 2 Review of Definitions - Atom: smallest quantity of matter • Concept of atom introduced by Greek philosopher Democritus - Element: simplest type of substance with unique physical + chemical properties - Compounds: a substance composed of two or more different elements which are chemically combined (can be broken down by chemical means) • H2O —> H2 and O2 CO2 —> CO2 and O2 CH4 - Mixture: consist of two or more substances that are physically intermingled NOT chemically combined • You can homogenous mixtures and heterogenous mixtures - Homogenous: consistent composition; no visible boundaries - Heterogeneous: visible boundaries - Molecule: consists of two or more atoms that are chemically bound together • Common diatomic elements: Br I N Cl H O F - Br2 I2 N2 Cl2 H2 O2 F2 • Behaves as an independent unit - All covalent compounds are molecules but not all molecules are compounds • Because we have for example the diatomic gases which are elemental - CQ: Which are pure elements? Na, O2, NH3 or NaCl dissolved in water Na and O2 • Chemistry Laws - Law of Mass Conservation • The total mass of substances present does to change doing a chemical reaction 1 Wednesday, January 20, 2016 - What you start off with you will end up with • Reactant 1 + Reactant 2 = Product Total Mass = Total Mass Calcium oxide + carbon dioxide —> calcium carbonate CaO + CO2 —> CaCO3 56.08 g + 44.0 g —> 100.08 - Matter cannot be created or destroyed. Mass of Reactants = Mass of Products - Law of Definite (or Constant) Composition • No matter what the source of a compound, its elements occur in the same proportion by mass - Calcium carbonate (can come from eggshells, coral, from anything but the elements are the same) • Analysis by Mass (grams/ 20.0g) Mass Fraction (parts/ 1.00 part) - 8.0g Calcium — /20 —> .40 Calcium - 2.4g Carbon — /20 —> .12 Carbon - 9.6g Oxygen — /20 —> .48 Oxygen - You can take all this and the mass percentage —> from the mass fraction - Mass fraction = mass of element x compound A / mass of compound A - Analysis shows that 84.2 g of pitchblende contains 71.4 g of uranium with oxygen as the only other element. How many kg of oxygen are combined in a small of pitchblende that contains 2.3 kg uranium • Mass oxygen= 84.2 g pitchblende - 71.4 g U = 12.8 g Oxygen - 2.3 kg U —> 2300 g U - Mass oxygen/ 2300 g U = 12.8 g O/ 71.4 g U • = 412.3 g —> 0.41 kg oxygen - Law of Multiple Proportions 2 Wednesday, January 20, 2016 • You can have two different elements that combine in two different whole number ratios to get two different compounds Dalton’s Atomic Theory 1. All matter consists of atoms; tiny indivisible parties of an element that cannot be created destroyed • Not true because you can break an atom down to more like protons, neutrons, electrons - Subatomic Particles • Protons + p^+ • Neutrons neutral n • Electrons - e^- • What makes up the nucleus of an atom? - Protons and Neutrons 2. Atoms of one element cannot be converted into atoms of another element • False (Nuclear transmutation) - By bombarding atoms together you are colliding their nucleus change 3. Atoms of an element are identical in mass and other properties and are different from the atoms of any other element • Not true, because of isotopes 4. Compounds result from the chemical combination of a specific ratio of atoms of different elements • TRUE Landmark Discoveries - The people and their experiments! • JJ Thomson —> Cathode Rays - Discovered existence of electrons 3 Wednesday, January 20, 2016 - Also discovered the charge to to mass ratio of electrons • Robert Millikan —> Oil drop experiment - Determined exact charge of an electron as 1.602 x 10^-19 C - C —> coulomb: A derived SI unit of electric charge • Rutherford —> Gold foil experiment - Most of the atom is empty space - Protons are concentrated at the center, called the nucleus • Chadwick —> Bombardment of Be with alpha particles (physics) - Discovers the neutron Subatomic Particles: The Electron - Radiation: refers to emission and transmission of energy in the form of waves - Cathode ray tube - evacuated glass chamber with two metal plates sealed inside • When the metal plates are connected to a high voltage source, radiation known as cathode rays are emitted from the negatively charged plate (called cathode) - J.J Thomspon: Cathode-ray Tubes • Thompson: rays are streams of negatively charged particles called electrons • Ray bends in magnetic field —> consists of charged particles • Ray bends toward positive electric field —> neragitce particles • Ray is id to any ray —> matter is made up of atoms Subatomic Particles: The Proton - Alpha particles emitted by the source and strike through the gold foil, had no deflection most of the time Minor deflection seen occasionally • - But mostly as the alpha particle hits the gold foil it does deflect it goes back - Mass of atom is heavy because of the nucleus 4 Wednesday, January 20, 2016 Subatomic Particles: The Neutron - Electrons and proton alone could not account for the composition of an atom - Offset in mass: • Hydrogen has 1 proton, helium has 2 • But the hydrogen atom only has 1/4 the mass of a helium atom - Discovery of the neutron • James Chadwick • Similar mass as proton Nuclear Model of the Atom - In actuality, neutrons and prions are mad cup of even smaller elemearty particles called quakes • These particles are together by “particles” of force Isotopes - Not all atoms of the same element are created identical • Isotopes: atoms that have the same atomic number (Z), but different mass number - Means that the number of protons stay the same - BUT, the number of neutrons is different - Carbon, carbon os the most abundant element in nature, used every. Carbon (12) is the most abundant, Carbon (13) is very radioactive • Isotopes are naturally occurring 5 Wednesday, January 20, 2016 - Isotopic Symbol - What is the correct symbol for a neutral element that has 20 protons and 20 neutrons? • Ca —> 40 mass number (protons +neutrons) and 20 protons - Average atomic mass • Mass number is equal to the sum of protons and neutrons an element has - Most convient unit to repot mass number is the atomic mass unit (amu) • 1 am = 1.66 x10^-27 • Br - 79: 78.918337 amu • Br - 81: 80.9162906 amu Average: 78.918337 + 80.9162906 / 2 = 79.917 • - But on the 79.904, why the difference? • Because in reality not all isotopes will occur in the same ratio, so we don't take the average but, the Weighted Average • In reality its not going to be 50/50, so you need to do a Percent Abundance - How much of each isotope is naturally occurring? • Br - 79: 50.69% (given) • Br - 81: 100 - 50.69 = 49.31% • Always has to add up to 100 • Calculating a weighted average: - Br - 79: 50.69% Abundance; 78.918337 amu - Br - 81: 49.31% Abundance; 80.9162906 amu 1. Convert each % to decimal Br - 79: 50.69/100 = 0.5069 Br - 81: 49.31/100 = 0.4931 2. Multiply each contribution by its corresponding mass in amu 6 Wednesday, January 20, 2016 Br - 79: (0.5069)(78.918337) = 40.00370503 amu Br - 81: (.4931)(80.9162906) = 39.8998229 amu 3. Add them together • 40.0037…. + 39.899….. = 79.9…. amu - (% of IS#1/100)(Mass IS#1) + (% of IS#2/100)(Mass IS#2) = weighted average - Mass of Particular Isotope • Gallium has two naturally occurring Isotopes, Ga-69 (68.92588; 60.108%) and Ga-71. The average atomic mass of Ga is 69.723 amu - Find the average mass of Ga-71 and its percent abundance - Ga-69= (0.60108)(68.92558) + (0.39892)X = 69.723 amu X = 70.925 amu - Finding Percent Composition of Each Isotope (Percent abundance is NOT given) • Potassium has two naturally occurring isotopes, K-39 (38.9637) and K-41 (40.9618 amu). What is the percent abundance of each isotope? • Average mass of K = 39.0983 amu (From PT) x = % abundance of K-39 y = % abundance of K-41 38.9631(x) + (40.9618)(y) = 39.0983 x + y = 1 —> y= 1-x 38.9637(x) + 40.9618 (1-x) = 39.0983 x= 0.9326 x 100 = 93.26% K-39 y = .0674 x 100 = 6.74% K-41 - Finding Average Atomic Mass • Average Atomic Mass = *sum of ( isotopes mass x % abundance / 100%) - How can you tell which isotope is more abundant? - Since dealing with wighted averages whichever isotope is closet to the average atomic mass, that isotope is most abundant 7 Wednesday, January 20, 2016 - Two stable isotopes of boron have aseptic masses of 10.0139 amy and 11.0093 amu. The cited atomic mass of the element is 10.81. Each isotope is more abundant • Find the difference from the amu Periodic Table - Main Group —> 1A and 2A and 13-14 - Transition metals —> 3-12 - Group 1A: Alkali metals (except hydrogen) - Group 2A: Alkaline earth metals 8 Wednesday, January 20, 2016 - Group 7A: Halogens - Group 8A: Noble Gases - 9


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