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Lecture 1

CHM 113 Lecture 1: Test 1 Notes

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Arizona State University
CHM 113
Ron Briggs

Chapter 1- Disciplines within Chemistry - Organic: study of organic compounds (primarily Carbon and Hydrogen) - Inorganic: study if inorganic compounds (substances not composed of primarily Carbon and Hydrogen) - Biochemistry: study of chemical processes that occur in living things - Physical Chemistry: application of physical principles and measurements to understand properties of matter - Analytical Chemistry: separation, identification, and quantification of chemical components Interdisciplinary areas within Chemistry - Environmental Chemistry - Radiochemistry - Geochemistry Classification of Matter Matter-​ anything that has a mass and occupies space Pure substances-​ matter with a fixed composition and distinct properties, can’t be separated by normal physical or mechanical means (grinding, using fingers, filtering, magnets, distillation, etc.) Element-​ a pure substance consisting of one type of atom, listed on the periodic table by atomic symbol Compound-​ atoms of two or more elements bonded together Mixtures-​ two or more pure substances intermingled, can usually be separated by physical/mechanical mechanisms Homogeneous Mixture/Solution- ​uniform composition throughout, but the percentage of each component in that mixture can vary from one component to the next, all components must be in the same phase (liquid, gas, or solid); Ex: Air, Bronze =, Gasoline, Vinegar, and Salt) Heterogenous Mixture-​ non-uniform composition, may contain multiple phases; Ex: Marble, Orange Juice with Pulp Properties of Matter Physical - A change in appearance or phase (solid, gas or liquid), but not the composition - Ex: Bending, Breaking, Phase Changes ( Melting, Vaporization, Condensation, Sublimation, Deposition) - The temperatures at which these phase changes occur are examples of physical properties. Different substances have different sets of physical properties. - Intensive (Bulk) Properties: Does not depend on the amount of material; Ex: color, temperature, density, hardness, etc.) - Extensive Property- depends on the amount of material that is present (volume, mass, & length) ​Chemical - Behaviour when combined with another substance. In chemical change (a reaction), a substance, is transformed into different substances Units of Measurement Different units can be used for the same measurement 1 in = 2.54 cm 1 min = 60 seconds 23 1 mole of atoms= 6.022 x10 atoms - SI System: standard system of measurement used in science consists of 7 fundamental (base) SI units - Other units can be expressed as some combination of these 7 units Metric Prefix What it means Abbreviation 15 Peta 10 P Tera 10 12 T Giga 9 G 10 6 Mega 10 M Kilo 10 3 K −1 Deci 10 d Centi 10 −2 c Mili 10 −3 m −6 Micro 10 μ Nanon 10 −9 n −12 Pico 10 p Femto 10 −15 f Atto 10 −18 a −21 Zepto 10 z Ex: Converting 15.5 mm to ​n​m 9 15.5mm ( 10 mm )( 11mm) = 1.55x10 n​m - Volume (V): amount of space occupied by a substance. Units: L, mL,cm ,m , etc. 3 g g kg - Density (D): the mass per unit volume of a substance. Units: mL , cm 3, m3,etc. - Ex: What is the mass of 13.4cm of Pb if Pb has a density of 11.4 g ? cm3 Conversion Factor Method 11.4g 2 13.4 cm 3 ( cm3 ) = 153g = 1.53x 10 g Equation Method m g 3 2 D = V → m = DV =(11.4 cm3 )(13.4cm ) = 1.53x 10 g Dalton’s Atomic Theory (1803-1807) - Elements are composed of very small particles (atoms) - All atoms of the same element are identical - Atoms of one element can’t be changed into atoms of a different element - Compounds are founded when atoms of more than one element are combined - Compounds always have the same relative number and kind of atoms (Law of Constant Composition) - Atoms can’t be created nor destroyed in a chemical reaction (Law of Conservation of Matter) Atomic Structure - J.J. Thompson (1897): discovered a stream of electrons (cathode ray) by applying a potential a potential through two electrodes in an evacuated glass tube - Robert Millikan (1909): used Thompson’s data and an Oil Drop Experiment to determine the charge and mass of an electron 1.60x101c −28 - Electron mass= 1.76x108gc= 9.10x10 g - Henri Becquerel Pierre and Marie Curie (1895-1911): Discovered and characterized radioactivity (atomic level disintegration) - Ernest Rutherford (1910): Alpha Particle Scattering Experiment proved the existence of a dense, positively charged nucleus (and disproved Thomson's “Plum Pudding model”) - Ernest Rutherford (1919): Discovered the proton - James Chadwick (1932): Discovered the neutron - Elementary (Subatomic) Particles: Particle Symbol Actual Charge Relative Charge Mass (amu) Electron e− -1.602x10 −4 -1 5.486x10 −4 Proton p+ +1.602x10 −4 +1 1.0073 Neutron n0
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