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CHEM1110 Week 1 Lecture Notes Dr. Lee

by: Madison Stewart

CHEM1110 Week 1 Lecture Notes Dr. Lee Chem 1110 04

Marketplace > University of Tennessee - Chattanooga > Chemistry > Chem 1110 04 > CHEM1110 Week 1 Lecture Notes Dr Lee
Madison Stewart
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About this Document

I compiled notes I took during our first two classes. It includes most everything I thought was important... except conversions, haven't gotten to those yet
General Chemistry I
Class Notes
Chemistry, General Chemistry, significant figures, significant, figures, properties




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This 5 page Class Notes was uploaded by Madison Stewart on Friday August 26, 2016. The Class Notes belongs to Chem 1110 04 at University of Tennessee - Chattanooga taught by Lee in Fall 2016. Since its upload, it has received 61 views. For similar materials see General Chemistry I in Chemistry at University of Tennessee - Chattanooga.

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Date Created: 08/26/16
#1 lecture notes MON 8/22/16     Science and Chemistry  ­ Science is super tough because it’s like a foreign language, it helps to  approach it as such  ­ Chemistry is a central science, without it, there would be no physics  or biology    The Study of Chemistry  ­ “Matter is basically everything, we just study specific bits of it” ­Dr.Lee  ­ sub­microscopic/ molecular  ​ = too small to see with even a  microscope  ­ Macroscopic​ ­ you can see with your eyes     Classification of Matter  ­ This slide just describes the difference between a solid a liquid and a  gas, pretty self explanatory lololol    Classification of Matter  ­ Substance ​= elements and/or molecules/compounds  ­ Mixtures:  ­ Homogeneous  ​ = uniform throughout (also a solution)  ­ Definite composition wherever sample is taken  ­ Ex.​  saltwater, apple juice  ­ Heterogeneous  ​ = NOT uniform throughout   ­ ​ex​. Trail mix, chicken noodle soup  The properties of Matter  ­ Quantitative ​properties ­ (math comes in here) these are the ones  you can measure with numbers (more common in classwork) ​ex.  Mass,   ­ Qualitative ​properties ­ (adjectives come in here) these are the ones  you describe (more common in lab work) ​ex. ​  Color, smell    ­ Physical property​ ­ can be observed observed and measured  without changing the identity of the substance (colors, melting point,  boiling)  ­ Chemical property​ ­ describes a quality that exhibits as it interacts  with another substance    The Properties of Matter  ­ Chemical change ­ ​  results in the change of composition, the original  doesn’t exist anymore ​ex.  ​ Chemical reaction i.e. burning  ­ Physical change  ​ ­ composition stays the same e ​ x.​ Changing  properties (liquid to solid), chopping up    Scientific Measurement  ­ Units are very important!! Final answers should include a metric unit  with it  ­ Ex. meters, liters    SI base units  ­ Meter, kilogram, second, kelvin, mole  ­ When writing equations on paper use normalized (1.0 x 10^6) not the  “E” form (that's for the calculator only) both are common when  describing molecular length  ­ Angstrom ­ ​  (non SI unit: 1.0 x 10^­10)    Mass  ­ A base unit, always the same (even if you’re on the moon), unlike  weight (it depends on gravity)  ­ The gram is most commonly used in chemistry  ­ Atomic mass unit (amu)​ ­ used to express the masses of atoms and  other super teeny tiny stuff  ­ 1 amu = 1.661 x 10^­24    Temperature  ­ Celsius​: based on water  ­ 0 degrees = freezing ­ 100 degrees = boiling    ­ Kelvin​: based on absolute zero  ­ 0 degrees K = absolute zero (can’t go colder)  ­ K = degrees C + 273.15  ­ Ignore fahrenheit in this class    Properties of Matter  ­ Extensive ​property:changes depending on the amount of matter  ­ Ex. the ​volume ​and m ​ ass  ​ of an object change depending on  how much of it you have  ­ ​Intensive ​property: stays constant no matter how much of it you may  have  ­ temperature and density don’t change if you add or take away any               Matter    **​no matter if you have one atom of a substance or a brick of the  substance, the density will always be the same.* ​ *           #2 Lecture Notes WED 8­24­16    Uncertainty in measurement  ­ Exact numbers​: defined values, numbers that are c ​ ounted.  ​ ​example.  ​ 1 kilogram =  1,000 grams, 2 dozen = 24, 30 students in a class ( ​ unlimited amount of significant digits)  ­ Inexact numbers​: measured any other way than counting (ex. If you use a ruler or  scale) ​example​. Length, mass, time, speed     ­ Significant figures/digits : meaningful digits in a reported number  ­ Uncertain digit : last digit in a measured number (+ or ­ 1 in the last digit)    Accuracy and Precision  ­ Accuracy​: closeness to the true value  ­ Precision​: how well an outcome is replicated     SIGNIFICANT FIGURES:  ­ Non zero = significant    ex. 234.5 ­ 4 sigfigs   ­   ­ Sandwiched zeros = significant     ex. 56,007 ­ 5 sigfigs,    20.019 ­ 5 sigfigs  ­   ­ Zeros to the left = NOT significant    ex. 0.0042 ­ 2 sigfigs (4.2 x 10^­3)  ­   ­ Zeros to the right with a decimal = significant   ex. 1.340 (4 sigfigs)  ­   ­ Zeros to the right with NO decimal = significant (at least they are in class) but otherwise,  it can be ambiguous.    Ex. 300 (3 sigfigs)    Calculating with measured numbers  ­ addition/subtraction: use smallest decimal point in answer  ­ multiplication/division: use smallest sigfig  ­ If calculating with an exact number, include all decimal points     ­ When calculating with multiple steps, don’t round any numbers until the end, use the  ANS key on your calculator to use the exact number from the previous step.    Using Units and Solving problems  ­ Conversion factor: a fraction in which the same quantity can be expressed two different  ways  ­ EX. 1 in. = 2.54 cm can be expressed as:  1 in.    2.54 cm      Or     2.54 cm  1 in    Conversion notes to come…. 


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