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Lectures notes 1/7/16 Chapter 1

by: Kennedy Connolly

Lectures notes 1/7/16 Chapter 1 CHM 111

Marketplace > Oakland University > Chemistry > CHM 111 > Lectures notes 1 7 16 Chapter 1
Kennedy Connolly
Oakland University
GPA 3.98
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About this Document

These notes cover everything we did in class on 1/7. Professor Felton's hand written notes can be a bit hard to understand sometimes, so these may be a little easier. Please let know if there is an...
General Chemistry 1
Greg Felton
Class Notes




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This 3 page Class Notes was uploaded by Kennedy Connolly on Thursday January 21, 2016. The Class Notes belongs to CHM 111 at Oakland University taught by Greg Felton in Winter 2016. Since its upload, it has received 323 views. For similar materials see General Chemistry 1 in Chemistry at Oakland University.

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Date Created: 01/21/16
Thursday 1/7 Lecture Chapter 1 Units of Measurement: Quantity Unit Symbol length meter m mass kilogram kg time second s temperature Kelvin K Amt. of Substance Mole mol. Mass: - Measure of a quantity of matter - Standard unit NOT grams Temperature: - Measure of the amount of kinetic energy of the atoms or molecules in matter - Temperature determines the direction of thermal energy (heat) transfer - Thermal energy transfers from hot —> cold objects - Kelvin scale starts at 0 K - 0 K is the coldest temperature possible (also known as absolute 0) - 0 K = -273.15 degrees C or -459 degrees F (e.g.) 40 degrees C K = 40.00 degrees C + 273.15 = 313.15 K Prefix Multipliers: - Change the unit value by powers of 10 Mega M 1,000,000 10^6 Kilo k 1,000 10^3 Deci d 0.1 10^-1 Centi c 0.01 10^-2 Mili m 0.001 10^-3 Micro μ 0.000001 10^-6 Nano n 0.000000001 10^-9 Significant Figures (SF): - Writing numbers to reflect level of precision - The greater the number of significant figures, the greater the certainty of that value of the measurement - Significant Figures Rules: 1. All non zero digits are significant —> (e.g.) 28.12 = 4 SF’s 2. Interior zeros are significant —> (e.g.) 28.02 = 4 S.F’s 3. Leading zeros are not significant —> (e.g.) 0.0032 = 2 S.F. 4. Trailing zeros are significant —> (e.g.) 45.000 = 5 S.F. • (e.g.) 1200 —> ambiguous because there is no decimal, therefore there could be 2 or 4 SF’s 5. Exact Numbers • Have unlimited SF • Exact count of discreet objects • Defined quantity • Numbers that are part of an equation - In calculations using measured quantities, the result must reflect the precision of the measured quantities - (e.g.) 1.052 x 12.504 x 0.53 = 6.7208 = 6.7 because 0.53 only has 2 SF’s - In multiplication and division, the result carries the same number and the number with the fewest SF’s - (e.g.) 2.0035 / 320. = 0.626044 = 0.626 because 320. has 3 SF’s - In addition and subtraction the result carries the same number of decimal places as the quantity with the fewest decimal places - (e.g.) 2.345 + 0.07 + 2.9975 = 5.4125 = 5.41 because .07 has two decimal places Accuracy: - How the measured value is to the actual value Precision: - How close a series of results are to one another (produces the same result) (e.g.) A B 10.99 9.78 9.79 9.82 9.92 9.75 10.31 9.80 Mean = 10.18 Mean = 9.79 Set A is accurate, not precise Set B inaccurate, but precise Solving Chemical Problems: - Unit conversions - Dimensional analysis - Units should always be used - Unit equation - (2.54 cm = 1 inch) - Conversion factor - A fractional quantity of a unit equation with the units we are converting from on the bottom and units we are converting to on the top - Info given x Conversion factor = Desired information Dimensional Analysis - Unit raised to a power - Raise both the number and the unit to a power - (e.g.) in^2 —> cm^2 - 2.54 cm = 1 in - (2.54 cm)^2 = (1 in)^2 - (2.54)^2 cm^2 = (1)^2 in^2 - 6.45 cm^2 = 1 in^2


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