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Chemistry 107 "General Chem." - Chapter 1 Notes

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by: Kalia Notetaker

Chemistry 107 "General Chem." - Chapter 1 Notes CHEM 107

Marketplace > University of Louisiana at Lafayette > Chemistry > CHEM 107 > Chemistry 107 General Chem Chapter 1 Notes
Kalia Notetaker
University of Louisiana at Lafayette
GPA 4.0

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These notes cover the entire chapter one, including some definitions and examples covered in class.
General Chemistry I
Richard Perkins
Class Notes
25 ?




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This 4 page Class Notes was uploaded by Kalia Notetaker on Saturday January 16, 2016. The Class Notes belongs to CHEM 107 at University of Louisiana at Lafayette taught by Richard Perkins in Spring 2016. Since its upload, it has received 155 views. For similar materials see General Chemistry I in Chemistry at University of Louisiana at Lafayette.

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Date Created: 01/16/16
Ch. 1 Notes  Macroscopic – big level Microscopic – small level  Scientific Method – systematic approach to research: Observation, Representation, Interpretation…Observation… o Hypothesis – explanation of what’s going on o Law – summary of many observations, concise statement of a relationship between phenomena that is always the same under the same conditions o Theory – picture, explanation of what’s going on, unifying principle explains body of facts and/or those laws that are based on them  Chemistry – study of matter and changes it undergoes o Matter – anything that occupies space and has mass o Substance – definite composition and distinct prop.. iron, liquid nitrogen, silicon crystals  3 states of matter – solid, liquid, and gas o Mixture – combination of two or more substances in which the substances retain their distinct identities  Homo- same throughout... soft drink, milk, solder  Hetero- compostion if not uniform throughout… cement, iron filings in sand, container with water with iron nail sticking out the top of contatiner (h2o .. can be broken down & iron can’t  Physical means can be used to separate mixture.. no chemistry going on.. distillation, magnet o Element – substance that can’t be separated into simpler substances by chemical means  Hydrogen & oxygen, water is a compound, can’t be broken down – elements  114 identified  82 occur naturally on earth – gold, aluminum, lead, oxygen, carbon, sulfur  32 have been created by scientists.. technetium, americium, seaborgium  Break mixture down, get substance, break down get elements o Compound – substance composed of atoms of two or more elements chemically united in fixed proportions  Only be separated into their pure components (elements) by chemical means.. lithium fluoride, quartz, dry ice (co2) o Classification of Matter: Matter, Mixtures (Homogeneous {solution} &heterogeneous), separated by physical methods, Pure Substances (compounds .. chemical methods.. elements) PG. 7 o Types of Changes: Physical – does not alter the composition or identity of a substance… ice melting, sugar dissolving in water Chemical – alters the comp. or identity of substance involved… hydrogen burns in air to form water o Extensive & Intensive Prop. – E: depends upon how much mater is being considered… mass, length, volume I: doesn’t depend upon matter consumption.. color, density, temperature.. breaking red temperature stick o Mass – measure of quantity of matter SI unit of mass is the kilogram 3 (kg) 1 kg = 1000 g = 1x10 g o Weight – force that gravity exerts on an object weight = c x mass Earth: c= 1.0 Moon = 0.1 A 1kg bar o Know Periodical Table and SI Base Units and Prefixes o Volume – SI derived unit for volume is cubic meter (m ) 1 cm = mL  Gal = about 4 L o Density – mass/volume (mass divided by volume) d=m/V, is an intensive property, can be used to characterize something, iron is heavy – has high density  Ex. a liquid has a density of 1.21 g/mL. What is the mass of 50.0 mL of the liquid? V(d=m/V)V v’s cancel out V x d = m m=d x V =1.21 g/mL x 50.0 mL = 60.5 g (mass)  Ex. What volume of acetone has a mass of 100.0 g? The density of acetone is 0.788 g/mL. V(d=m/V)V V x d = m (divide both sides by d) d cancels out on left side ….. V= m/d 100.0/.788 = 127 mL = V o Comparison of Temperature Scales – from Celsius to Kelvin (add 273) reverse subtract 273  Celsius 100 Boiling point of water 212 Fahrenheit 373 Kelvin  37 Body Temperature 98.6 Fahrenheit 310 Kelvin  25 Room Temperature 77 Fahrenheit 298 Kelvin  0 Freezing point of water 32 Fahrenheit 273 Kelvin  Ex. Mercury, the only metal that exists as a liquid at room temp., melts at -39.9 C. Convert its melting point to kelvins. (-38.9 =273 = 234.3 K) o Scientific Notation – 10 to the 0 is equal to one (10 = 1) make number between 1&10  602, 200, 000, 000, 000, 000, 000, 000 = 6.022 x 10 (move to3 left, positive subscript)  0.0000000000000000000000199 x 10 1.99 x 10 0 -2(move to right, negative sub) move to left, positive8  568.762 2 to left = 5.68762 x 10 2  0.00000772 x 10 0 move to right = 7.72 x 10 -6  Addition or Subtraction – 1.write each quantity with the same exponent, n 2.Combine N1 and N2 3.The exponent, n, remains the same  4.32 x 10 + 3.9 x 10 = 4.31 x 10 + .39 x 10 4  Multiplication – 1. Multiply N1 and N2 2. Add exponents on n1 and n2 -5 3 -5+3 -2  4.0 x 10 x 7.0 x 10 = (4.0 x 7.0) x (10 ) = 28.0 x 10  Division – 1. Divide N1 and N2 2. Subtract exponents n1 and n2 4 9 4-9 -5  8.5 x 10 / 5.0 x 10 = 8.5/5.0 x 10 = 1.7 x 10 o Significant Figures – any digit that is not zero is significant 1.234 =4  Zeros between nonzero digits are significant 606 =3  Zeros to the left of first nonzero digit are not significant 0.08 = 1  Zero immediately behind is significant  If a number is greater 1, then all zeros to the right decimal point are significant 2.0 = 2  If a number is less than 1, the only zeros that are at the end are significant 22  Ex. 393 = 3 5.03 = 3 0.714 = 3 0.052 = 2 2.720 x 10 = 4 3000 = ambiguous case, mat be 4, 3, 2, or 1  Add/Sub SFs – the answer can’t have more digits to the right of the decimal point that any of the original numbers 89.332 + 1.1 = 90.4 3.70 – 2.9233 = 0.7867 = 0.79  Mulit/Div. – the number of SFs in the result is set by the original number that has the smallest number of SFs  Exact #s - #s from definitions or numbers of objects are considered to have an infinite number of SFs o Accuracy & Precision – A: how close a measurement is to the true value P:how close a set of measurements are to each other  Dimensional Analysis Method of Solving Problems – given quantity x conversion factor = desired quantity o Ex. An object has a mass of 85 g. What is this in kg? (1000 g= 1 kg) 85 g x 1 kg/1000 g grams cancel = .085 kg o Ex. a liquid helium storage tank has a volume of 275 L. What is the 3 3 volume in m ? 275 L x 1000 mL/L x 1cm /100 cm


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