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CHEM1030 Chapter 1 Lecture Notes

by: Caroline Smith

CHEM1030 Chapter 1 Lecture Notes CHEM 1030 - 002

Marketplace > Auburn University > Chemistry > CHEM 1030 - 002 > CHEM1030 Chapter 1 Lecture Notes
Caroline Smith
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These notes all class lecture notes for chapter 1 in our assigned textbook.
Fundamentals Chemistry I
Dr. Rik Blumenthal
Class Notes
Chemistry, chem1031




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This 8 page Class Notes was uploaded by Caroline Smith on Thursday September 1, 2016. The Class Notes belongs to CHEM 1030 - 002 at Auburn University taught by Dr. Rik Blumenthal in Fall 2016. Since its upload, it has received 149 views. For similar materials see Fundamentals Chemistry I in Chemistry at Auburn University.


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Date Created: 09/01/16
CHEM 1030 – Rik Blumenthal Chapter 1 Lecture Notes  Chemistry: The Science of Change  Expressing Scientific Quantities  Chemistry is an experimental science which requires specific ways to express  quantitative (numerical) data.  Some quantities are exact  CERTAINTY of value  Ex 1. Integers (whole numbers)  Ex 2. Π = 3.1415…  Ex 3. e = 2.71…  Some quantities are inexact  UNCERTAINTY of value  Ex 1. Distance from Auburn to Tuscaloosa is about 156.4 miles. There is assumed error of ± 0.1 miles because distance depends on your exact location in Auburn (uncertainty).  Ex 2. Population of U.S. is about 324,000,000 people. There is assumed error of ± 1,000,000 people because the population is constantly changing (uncertainty).  Quantities can be precise, accurate, neither or both.  Precision: the closeness of repeated measurements to one another.  Accuracy: the closeness of measurements to an actual value.  To express the precision of a measurement or measuring tool, significant  figures are used.  Significant Figures include all known values plus one estimated  value. RULES FOR SIGNIFICANT FIGURES 1. Nonzero numbers are significant. 2. Zeros between nonzero numbers are significant. 3. Zeros to left of the first nonzero number are not significant. 4. Zeros to right of last nonzero number are significant IF A DECIMAL IS PRESENT. 5. Zeros to right of last nonzero number are not significant IF NO DECIMAL IS PRESENT. RULES FOR CALCULATING MEASURED VALUES 1. Addition and Subtraction: the sum or difference cannot be more precise than least precise measurement. 2. Multiplication and Division: product or quotient cannot have more significant figures than the measurement with the least number of significant figures.  Scientific Measurement  The universal system of measurement is the SI Unit System.  SI units use base­10 nature  Ex 1. 1 kilogram (kg) = 10 ^2 grams (g)  Ex 2. 1 centimeter (cm) = 10 ^­2 meters (m)  Ex 3. 1 m = 10 ^2 cm Seven Common SI Units Measurement Unit Symbol Length Meter M Mass Kilogram Kg Time Second S Electric Current Ampere A Temperature Kelvin K Amount of ubstance Mole mol Luminous Intensity Candela cd *Length, mass and time most important for chapter 1.  Derived units are combinations of base units.  Volume  SI Units = m^3  Conversion Factors: HELPFUL CONVERSION o 1 L = dm^3 FACTOR FOR REAL WORLD o 1 mL = cm^3 APPLICATION  Density 1 m = 1 yd  SI Units = kg/m^3 1 in = 2.54 cm 1 L = 1 quart 2 L = ½ gallon 2.21 lb = 1 kg 1 kg = mass of 1 L of water  Dimensional Analysis  In problems using conversion factors, dimensions must be tracked throughout the  calculation.  Ex 1. Use dimensional analysis to convert 12.00 inches into meters.    12.00 in * 2.54 cm / 1 in * m / 100 cm = 0.3048 m  Dimensional analysis will NOT tell you if an equation is right.         Ex 1. V = RT / M     Dimensions for velocity (V) are (m/s). TEST FOR WRONGNESS V = (8.324 J/K*mol)(300 K) / (800 kg/mol) Simplify units – ignore numbers for analysis. J = kg*m^2/s^2 V = (J/K*mol)(K) / (kg/mol) = (J/mol) / (kg/mol) = J/kg =  (kg*m^2/s^2)/kg = m^2/s^2  BUT V = m/s So resulting equation is V = √ RT/M Actual equation is V =  √ 3RT/M *Conclusion: Dimensional analysis can’t tell you what the right equation  is, so don’t use it for this purpose. PREFIX FOR CONVERSION FACTORS IN CALCULATIONS Tera Giga Meg Kilo Hect Deca Deci Centi Milli Micro Nan Pico a o o 12 9 6 3 2 1 ­1 ­2 ­3 ­6 ­9 ­12 10 10 10 10 10 10 10 10 10 10 10 10 T G M K H D d c m µ n p  Classification of Matter  Matter: a substance or mixture of substances   Substances have a constant, definite compositions (made of the same  stuff).  Ex 1. NaCl – ALWAYS 1 sodium (Na) atom to 1 chlorine (Cl)  atom  Matter comes in 3 different states:  Solids: have fixed neighbor distances and fixed geometric  arrangements (closely packed particles that do NOT take shape of  container they are in).  Liquids: have fixed neighbor distances but NO fixed geometric  arrangement (closely packed particles that take shape of container  they are in).  Gases: have NO fixed neighbor distances but have fixed geometric arrangements (loosely packed particles that take shape of container they are in).  Mixtures: combinations of substances NOT chemically bonded. *Each substance keeps its distinct properties.  There are two types of mixtures: o Homogeneous mixtures have a uniform composition. The  amount of solute in solvent are the same in different areas of the solution.  Ex 1. 1 teaspoon of sugar in 2 cups of water  would be uniformly distributed throughout the  solution. o Heterogeneous mixtures do NOT have a uniform  composition. The amount of solute in solvent is different  in different areas of the solution.  Ex 1. Mud has rocks, sticks, leaves etc. mixed in it,  but the objects are not uniformly distributed  throughout the solution.  Properties of Matter  Matter can change and be changed in physical and chemical ways.  Physical   Physical Properties: can be observed or measured without  changing the chemical composition. o Ex 1. Color can be observed without changing the  chemical composition of a substance. o Ex 2. Density can be measured without changing the  chemical composition of a substance.  Physical Changes: changes the state of matter without changing  the chemical composition of matter. o Ex 1. Freezing water changes the state of water, but the  substance is still water. o Ex 2. Breaking a rock changes the state of the rock, but  the substance is still a rock.  Chemical  Chemical changes change the composition of matter. The original  substance no longer exists. o Ex 1. Baking a cake changes the composition of cake  batter.  Chemical properties CANNOT be observed without a chemical  change occurring. o Ex 1. Rusting of a bike reveals an oxidative chemical  property of metal.  Matter also has extensive and intensive properties.  Extensive properties depend on the amount of matter there is.  Ex 1. Mass depends on how much matter there is in an object.  Intensive properties are independent of (do NOT depend on) the amount  of matter there is.  Ex 1. Density does not depend on the amount of matter because  any amount of a particular substance will always have the same  mass to volume ratio (1 mL of water has same density as 200 mL  of water).


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