Fundamentals of Chemistry 1 (CHEM 1331) Week 1 Notes
Fundamentals of Chemistry 1 (CHEM 1331) Week 1 Notes CHEM 1331
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This 7 page Class Notes was uploaded by Alexis Clowtis on Monday February 1, 2016. The Class Notes belongs to CHEM 1331 at University of Houston taught by Thomas Teets in Spring 2016. Since its upload, it has received 176 views. For similar materials see Fundamentals of chemistry in Chemistry at University of Houston.
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Date Created: 02/01/16
Chemistry 1331- 12pm MWF SEC 100 Professor Thomas Teets Units and Significant Figures 01-20-2016 Units: required to make a measurement have meaning; ALL MEASURED QUANTITES NEED UNITS SI Units to Know: Mass Kilogram (kg) Length Meter (m) Time Seconds (s) Temperature Kelvin (K) *NO degree symbol* Amount of substance Mole (mol) Prefixes to Know: Table R-2 in Textbook When these prefixes are added to units, it changes the size of that unit by the given amount (in multiples of 10 ) Mega (M) 10 6 Kilo (k) 10 3 2 Hecto (h) 10 Deka (da) 10 --------Baseline 1--------- (no prefix) --------Baseline 1--------- Deci (d) 10 -1 -2 Centi ( c) 10 -3 Milli (m) 10 Micro (µ -pronounced “mu”) 10 -6 Nano (n) 10 -9 3 Example: meter kilometer = 10 (1000) meters Uncertainty in Measurement: All measurements have some degree of uncertainty Certain digits: digits not estimated Uncertain digits: digits must be estimated due to limited precision of measuring device Reporting: all certain digits plus one uncertain digit o Assume ±1 error in the uncertain digit Chemistry 1331- 12pm MWF SEC 100 Professor Thomas Teets Analog vs. Digital scale Analog Reported Mass: between 18.6 kg and 18.8 kg Digital Reported Mass: 18.737 kg ± 0.001kg so you still have some uncertainty but it is smaller than the analog scale because the uncertain number is in the third decimal place instead of the first which implies it is a more precise measurement. Accuracy vs Precision “True Value”: accepted value for a measurement accomplished by: 1. Repeated measurement st 2. Determined theoretically from 1 principle Accuracy: how close a measurement is to the true value Precision: how well multiple measurements math with each other Dart Board Analogy- Poor accuracy Poor precision Large Random Errors (“indeterminate error”) Chemistry 1331- 12pm MWF SEC 100 Professor Thomas Teets Poor accuracy Good precision Small random error Large systematic error (“determinate error”) Good accuracy Good precision [End of lecture] Significant Figures “In a measured quantity, all certain and one uncertain digits are considered to be significant” All nonzero figures are significant “Leading zeros”- All zeros BEFORE first nonzero are NOT significant “Captive zeros”- Zeros between nonzero digits are significant “Trailing zeros”- Zeros at the end of the number are significant IF the value has a decimal point Exact Numbers: “Some numbers are assumed to have infinite significant figures.” o “Counted numbers” o Numbers in calculations/formulas o Conversion factors/definitions *All exact numbers are assumed to have infinite significant figures* Rules for Mathematical Operations Multiplication/Division: answer has the same number of significant figures as the number of significant figures in the multiplies value with the LEAST amount of sig figs. Example: 4.06 (3 sig figs) X 2.973 (4 sig figs)= 12.07038 Round to 12.1 (3 sfs) Addition/Subtraction: answer has the same number of decimal places as the calculation with the LEAST amount of decimal places Example: 4.06 (2decimals) + 2.973 (3decimals)= 7.033 7.03 (2 decimals) Chemistry 1331- 12pm MWF SEC 100 Professor Thomas Teets Mathematical Combinations of Measured Quantities 01.22.2016 Multiplication/division: the associated units are also multiplied/divided Example: Area- 8.5in X 11in= 94(in x in)= 94 in 2 Speed- Travel 51 miles to Galveston in .75 hours; speed= distance/time =51miles/.75hours= 68miles/hour Addition/Subtraction: The two quantities you’re adding/subtracting have to have the same units, so answer is those same units Dimensional Analysis/Unit Factor Method Units factors: ration of two equivalent quantities; can be used to convert one unit into another 2.54 ???????? 1 ???????? 1 ???????? 2.54 ???????? Why do they equal 1? 2.54 ????????= 2.54 ???????? 1 = 2.54 ???????? 1 ???????? Conversions to Know: 1 inch 2.54 cm 1 minute 60 seconds 60 minutes 1 hour Any with prefixes- ex. 10 J= 1KJ Examples of Unit Conversions: Convert 3.8 miles into feet. (1 mile=5280 feet) 4 = 20,064 feet= 2.0 X 10 ft Convert that distance into meters. = 6.1 X 10 meters Compound conversions- Converting value with multiple units Chemistry 1331- 12pm MWF SEC 100 Professor Thomas Teets Convert 65 miles per hour into meters per second = 29 meters/second Temperature Temperature is a measure of heat content Celsius Kelvin Fahrenheit Abbreviation ºC K ºF Freezing point H2O 0 273.15 32 Boiling point 2 O 100 373.15 212 Relative Size of 1 1 5/9 degree Unit Conversions: Celsius to Kelvin: k =c +273.15 same degree size, just different “0” points o 0 Kelvin is absolute 0 Celsius to Fahrenheit: T =(9ºF/5ºC)T +32ºF F c Fahrenheit to Celsius: c =(5ºC/9ºF)(TF- 32ºF) o Different degree size, different 0 point Examples: Convert -23.6ºC to Kelvin and Fahrenheit K=-23.6ºC + 273.15= 249.6 K Only one decimal place because given amount has 1 ºF= -23.6ºC(9ºF/5ºC) + 32ºF = -10.5ºF Density The mass of a substance per unit of volume Can be used to identify a pure substance Can be used to separate mixtures (layers) 3 Units: 1cm (solids)=1mL (liquids, gases) Density of water= 1g/mL=1g/cm 3 Chemistry 1331- 12pm MWF SEC 100 Professor Thomas Teets Example 1 An empty, container weighs 86.35 g, and when filled with water it weighs 161.61 g. When an unknown liquid is added, the mass of the filled container is 151.60 g. What is the density of the unknown liquid? 151.60g – 86.35g=65.25g (mass of liquid) ???? ???? = ???? 75.26 ???? M water61.61g-86.35g=75.26g Density of water=1 so 1 =???? ???????? so V=75.26 mL 65.25???? ???? = 75.26???????? = 0.8670 ????/???????? Example 2 Gold has a density of 19.32 g/cm3. How many cubic inches does a standard 12.4 kg gold bar occupy? (1 in = 2.54 cm) 3 = 39.2 in Classification of Matter Matter: anything that has mass Pure substance: matter with constant composition Classification of pure substances o Elements: cannot be decomposed to simpler things/substances physically or chemically o Compounds: pure substances with two or more elements Mixture: matter with variable composition, 2+ pure substances combined physically End of Lecture Chemistry 1331- 12pm MWF SEC 100 Professor Thomas Teets From professor’s notes that he did not cover (available on Blackboard)
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