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# Chemistry 110 Chapter 1 Sections 1.3-1.6 Chem 110

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This 6 page Class Notes was uploaded by Mariah J. Bibbs on Saturday September 5, 2015. The Class Notes belongs to Chem 110 at Northern Illinois University taught by M. Leifker in Fall 2015. Since its upload, it has received 27 views. For similar materials see Chem in Chemistry at Northern Illinois University.

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Date Created: 09/05/15

9115 Section 13 The Units of Measurement Chemistry calculates many equations by utilizing units 0 UnitsThe basic quantity of mass volume or whatever quantity is being measured A measurement is considered useless without units Units give meaning and value 0 There are two types of units that are used to calculate formulas English Imperial System Metric System 0 The English Imperial System is a collection of functionally unrelated units Metric System Units 0 There are two imperative units in the metric system Time and mass grams 0 Time is measured in seconds in the metric system 0 Mass is the quantity of matter in an object Not similar to weight Standard unit of grams lb pound is the common English unit 0 Mass is NOT the same as weight Mass us the quantity of matter in an object Weight is the force of gravity on an object o The formula for weight is WeightMass Acceleration 0 Length is the distance between two points The standard unit is the meter m The yard is the common English unit 0 Volume is the space occupied by an object The standard unit is the liter The quart is the common English unit 0 Several equations to calculate volume is Lwh 1m1m1m1mquot3 1m cubed is the same as 1L Liter Metric Unit Pre xes 0 Basic units are the units of a quantity without any metric prefix Prefix Abbreviation Meaning Decimal Equiv Equality wmetric unit mega M 10quot6 1000000 1 MX 10A6X kilo ka 10A3 1000 1 kX 10A3X deka da 10quot1 10 1 dax 10A1x deci d 10quot1 01 1 dx IDA1X centi c 10A2 001 1 cx 10A2x milli m 10A3 0001 1 m IDA3X Liters l uX IDA6X l nX IDA9X 10A6 10A9 0000001 0000000001 micro u nano n Example of Metric Equivalences 1 million grams 1 Mg 1000 g lkg 10 g l dag lg 10 dg l g 100 cg l g 1000 mg lg 1 million ug l g 1 billion 10 sixthL lML 10 thirdL l kL 10 L l dal l L 10 dL 1 L 100cl 1 L 10 third mL 1 L 10 sixth uL l L 10 ninth nL 14 The Numbers of Measurement 0 Significant figures are bearing digits or figures in a number and tell how accurate or precise a measurement is 0 The degree of uncertainty associated with a measurement is indicated by the number of figures to represent the information The less number you have the more uncertain a number is because more numbers indicate more precision and accuracy Recognition of Significant Figures All nonzero digits are significant 7314 has 4 sig figs The number of sig digits are independent of the position of the decimal point 7314 has 4 sig figs Zeros between nonzero digits are significant 60052 has 5 sig figs Zeros at the end of a number also known as trailing zeros are significant if the number contains a decimal point 470 has 3 sig figs Trailing zero are insignificant if the number does NOT contain a decimal point 100 has 1 sig fig 100 has 3 sig figs 0 Any zeros to the left of the first nonzero integer are not significant 00032 has 2 sig figs Scienti c Notation 0 Scientific Notation is used to express very large or very small numbers easily and with the correct number of sig figs Represents a number as a power of 10 Example A 4300 43 1000 43 10 third Uncertainty 0 Error is the difference between the true values and our estimation There are two types of errors that could occur during the calculation of an equation Random Systematic 0 Random error is data from multiple measurements is scattered 0 Systemic error is data uniformly smaller or larger that the accepted value 0 Outliers are random data points that effect measurements Types of Uncertainty 0 Two types of uncertainty are accuracy and precision Accuracy is the degree of agreement between the true value and the measured value Precision is the measurement of the agreement of replicate measurements Signi cant Figures in Calculation of Results Rules for Addition and Subtraction 0 The result in a calculation that CANNOT have greater significance than any of the quantities that produced the result 0 The following is an example of addition using significant figures 3768 liters 671862 liters 10842 liters 15282664 liters The correct answer is 15283 liters 0 There are two ways to solve addition and subtraction scientific notation using significant figures 947 x10 6 93 x10 5 One solution is converting both numbers to standard form and add or subtract 947 x10 6 93 x10 5 The second solution is to change one of the exponents so that both have the same power of 10 then add Rules for Multiplication and Division 0 The answer can be no more precise than the least precise number from which the answer is derived 0 The least precise number is the one with the fewest significant figures Exact and Inexact Numbers 0 Inexact numbers have uncertainty by definition 0 Exact numbers are a consequence of counting A set of counted items like beakers on a shelf for example has no uncertainty Exact numbers by definition have an infinite number of significant figures 9315 15 Unit Conversions 0 You must be able to convert between units with the metric system 0 The factor used for this method is the Factor Label Method also known as Dimensional Analyzes 0 Here is an example Convert 12 gallons to quarts The solution is A First write the starting quantity B The multiply by the conversion factor so each unit of the starting quantity is in the denominator C The gallons cancel themselves out leaving you with quarts 12 gal x 4qtlgal 48qt 0 Another example is Convert 360 feet to miles The solution is A Write the starting quantity Finish examples 16 Additional Experimental Quantities Temperature Conversions between Fahrenheit and Celsius ToC ToF 32 18 Conversion from F to C ToF 18 x ToC 32 Conversion from C and F The Kelvin Scale 0 The Kelvin Scale is an another important scale because it is related to molecular motion o The Kelvin temperature proportionately increases as the molecular speed increases Conversion from Celsius to Kelvin TK ToC 273 Energy 0 Energy is the ability to do work 0 There are two types of energy kinetic and potential energy also known as stored energy Kinetic energy is when energy is in motion Potential or stored energy is referring the position of energy 0 There are 5 forms that energy is categorized by Light Heat Electrical Mechanical Chemical Characteristics of Energy 0 Energy cannot be created or destroyed 0 Energy can be converted from one form to another 0 Energy conversion always occurs with less than 100 efficiency 0 All energy with a chemical reaction can have a gain or a loss in energy Units of Energy 0 Basic units of energy are Calorie Joule A l calorie cal 4184 joules J o A kilocalorie kcal is also known as the large Calorie This is equalivant to Calorie as food Calories l kcal l Calorie 1000 calories 0 l calorie the amount of heat energy required to increase the temperature of 1 gram of water loC Concentration 0 Concentration is the number of particles of a substance the mass of those particles and are contained in a specific volume 0 Concentration is used to represent the mixtures of different substances Concentration of oxygen in the air Pollen counts Proper dose of an antibiotic Density and Specific Gravity 0 Density is the ratio of mass to volume An extensive property Can be used to characterize a substance as each substance has a unique density Units for density include A gmL B gcmquot3 C gcc o The formula for density is D mass volume Calculations using Density 0 An example of this includes A 200 cm3 sample of aluminum is found to weigh 540g Calculate the density in gcmquot3 and g mL 0 Here is the solution Substitute the information given in the problem into the density formula d 540 g 270 gcm3 200 cm3 Since 1 cm3 1 mL the final answer is 270 gmL

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