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Chemistry 1210, Week 1 Notes

by: Grace Campbell

Chemistry 1210, Week 1 Notes CHEM 1210

Marketplace > Ohio State University > CHEM 1210 > Chemistry 1210 Week 1 Notes
Grace Campbell
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These notes cover everything that was gone through in class and in the book readings from 8/22/16 to 8/26/16
General Chemistry I
Class Notes
General Chemistry, Chemistry




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This 7 page Class Notes was uploaded by Grace Campbell on Friday August 26, 2016. The Class Notes belongs to CHEM 1210 at Ohio State University taught by in Fall 2016. Since its upload, it has received 30 views.


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Date Created: 08/26/16
WEEK 1 NOTES 8/22/16 – 8/26/16 Chapter 1: Matter and Measurement A. Chemistry- study of matter and how it changes; deals with the composition, structure, and reactions of matter a. Mass- the measure of the quantity of matter b. Weight- a result of the gravitational attraction between matter; a force but is often used interchangeably with “mass” c. Atoms- “building blocks” of matter d. Molecules- two or more atoms that have joined together e. Composition i. The kinds of atoms that matter contains ii. Can be at the macroscopic (can be seen and handled) or submicroscopic (atoms and molecules) level iii. Qualitative (what substances it is made of) or quantitative (the number of substances it is made of) f. Structure i. Arrangement; how atoms are put together ii. Makes a difference in what the final product is 1. EX: a sandwich layered in the usual order is different than a sandwich layered with the lettuce and cheese on the outside and the bread on the inside g. Reactions i. Occurs when the composition and structure changes ii. Exothermic (gives off heat) and endothermic (absorbs heat) B. Matter- anything that has mass and occupies space; physical material a. States of matter i. Gas (also known as vapor) 1. No definite volume or shape 2. Takes the shape of its container and expands to fill it 3. Very compressible 4. Expands greatly when heated ii. Liquid 1. Unchanging volume 2. Takes the shape of its container 3. A little compressible 4. Expands a little when heated iii. Solid 1 1. Definite volume and shape 2. Not compressible 3. Barely expands when heated b. Pure substance- composition does not change between samples used; definite properties i. Can only be separated by chemical methods c. Elements- the simplest substances that can not be broken down any further; consists of atoms that all contain the same number of protons i. 118 known elements d. Compounds- substances made up of 2 or more elements i. chemically combined ii. can be separated through chemical reactions iii. ionic compounds iv. molecular compounds v. Law of Definite Proportions/ Law of Constant Composition 1. The elements in a compound are always combined in the same way e. Mixtures- 2 or more substances physically, not chemically, combined i. Each substance retains its original chemical identity ii. Composition can vary iii. Homogenous 1. Solutions (can be solid, liquid, or gas) 2. Mixture is uniform throughout 3. Miscible- mix in all proportions iv. Heterogenous 1. Mixture is NOT uniform throughout 2. Parts vary throughout the mixture C. Properties of Matter a. Physical properties i. Observed without changing the substance ii. Physical appearance changes but its composition does not iii. EX: color, odor, density, specific heat, physical state, melting point, boiling point...evaporating water b. Chemical properties i. The way a substance reacts with another substance or transforms into a chemically different substance ii. EX: flammability…hydrogen burning in air c. Intensive properties i. Do NOT depend on the amount of a substance being used 2 ii. Very useful in identifying substances iii. EX: color, melting point, boiling point, density, specific heat, temperature d. Extensive properties i. Depends on the amount of a substance being used ii. EX: mass, volume, heat content e. Physical change i. Physical appearance changes but its composition does not ii. EX: change in state 1. Fusion- solid to liquid 2. Freezing- liquid to solid 3. Condensation- gas to liquid 4. Vaporization- liquid to gas 5. Sublimation- solid to gas 6. Deposition- gas to solid f. Chemical change i. Also called chemical reaction ii. Substance transforms into a chemically difference substance iii. Composition and/or structure changes g. Separation of mixtures i. Filtration- filtering a substance through a filter 1. Used for heterogenous mixtures ii. Distillation- utilizes the different abilities of substances to form gases 1. Used for separating a homogenous mixture 2. EX: boiling salt water…the water evaporates and the salt is left behind iii. Extraction- separation based on solubility differences iv. Chromatography- utilizes the different abilities of substances to adhere to the surface of solids D. Units of Measurement a. Prefixes i. Kilo (k)… 10^3 ii. Deci (d)… 10^-1 iii. Centi (c)… 10^-2 iv. Milli (m)… 10^-3 v. Micro (µ)… 10^-6 vi. Nano (n)… 10^-9 b. Mass = gram c. Length = meter i. 1 inch = 2.54cm d. Volume = m^3, liter (L), milliliter (mL) 3 i. 1dm = 10cm ii. 1mL = 1cm^3 e. Temperature i. 212° F = 100°C 1. boiling point of water ii. 98.6°F = 37°C 1. body temperature iii. 32°F = 0°C 1. freezing point of water iv. Conversion formulas 1. y° C = 5°C/9°F (x° F – 32°F) 2. y° F = 9°F/5°C (x° C) + 32°F 3. 0° C = 273.15 K v. 0 K is the lowest possible temperature f. Density i. D = m/V ii. Solids = g/cm^3 iii. Liquids = g/mL iv. Gases = g/L v. Specific gravity 1. Density of substance (g/mL) / density of water (g/mL) = specific gravity 2. No units 3. Density of water = 1.0 g/mL g. Precision- how close each measurement is to another i. Standard deviation- how much the individual measurements differ from the average ii. h. Accuracy- how close a measurement is to the correct value E. Significant Figures a. When recording from an instrument, report all the digits that you know PLUS one that you estimate b. Rules of Significant figures i. All nonzero digits are significant ii. Zeros between significant digits are significant 1. Captive Zeros 2. EX: 20.006 has 5 significant figures iii. Zeros to the left of the first nonzero digit are NOT significant 1. Leading zeros 2. Make sure to locate the decimal point 3. EX: 0.00004 has 1 significant figure iv. Zeros to the right of the last nonzero digit 4 1. The number ends in zero to the RIGHT OF THE DECIMAL POINT, zeros are significant a. Trailing zeros b. EX: 0.0040 has 2 significant figures c. EX: 4000.0 has 5 significant figures 2. The number ends in zero to the LEFT OF THE DECIMAL POINT, zeros are NOT significant a. Trailing zeros b. EX: 4000 has 1 significant figure c. EX: 4120000 has 3 significant figures c. Scientific notation i. 1 nonzero digit to the left of the decimal point ii. EX: 400 = 4 x 10^2 (1 significant figure) OR 4.0 x 10^2 (2 significant figures) OR 4.00 x 10^2 (3 significant figures) d. Calculations i. The result of a calculation has to reflect the significant figures of the original measurements ii. Multiplication and Division 1. The answer must possess the same number of significant figures as the original number in the problem that had the LEAST amount of significant figures a. EX: If dividing/multiplying 907.54 with 43.9, your answer must have 3 significant figures because 43.9 has 3 significant figures, which is less than 907.54 which has 5 significant figures 2. Rounding Rule a. If the leftmost number that is going to be left out because of rounding is less than 5, round down (the last number that will be remaining doesn’t change) iii. Addition and Subtraction 1. The last place in your answer will be the last place that is common to all of the numbers a. EX: 5 8.35 + 973.489 5 986.839 The correct answer is 987 because the last digit common to all of the numbers is the one right before the decimal point, so that has to be the digit where you answer stops 2. Rounding Rule a. If the leftmost number to be discarded is more than or equal to 5 and is followed by nonzero digits, round up 3. Rounding Rule a. if the number that is going to be left out is 5 or 5 followed by zeros, round even (leave the last digit unchanged if it is an even number, or increase it by 1 if it is an odd number) 6 WEEK 1 NOTES CONTINUED End of Chapter 1 A. Dimensional Analysis (Factor Unit Method) a. Conversion factor- a number that has 2 or more units associated with it i. Equal to 1 ii. Are exact by definition b. To get the same number but in a different unit, multiply the given information of one type of unit by a conversion factor c. EX: Convert 1.54cm to meters i. 1cm = 10^-2 m (this is your conversion factor) ii. 1.54 cm x 10^-2 m = 0.154 m = 0.2 m 1 1 cm


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