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Welcome to Chemistry! with Mrs. Guirguis Rm. 405 Do Now: Pick up homework reminder on table near the door. Find a seat & fill out Student Info Sheet. HWK: 1. Read syllabus and Frequently Asked Questions on my webpage. (Be prepared to answer questions on these two things on Monday.) 2. Start Intro to WebAssign: due Wednesday 3. Supplies & Signature Sheet Pr. 4- Lab on Day 2 (Mon): Safety Quiz Pr. 8 - Math Assessment on Mon (no need to study )

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What is Chemistry? Your Task: on a piece of paper answer this question... What is Chemistry? Does your answer sound like any of these responses? Does your answer sound like any of these responses?

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What is Chemistry? The definition well work off of this year: Chemistry is the study of matter & of the changes it undergoes Composition Structure Properties Energy changes

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A Quick Demo If we want to describe matter & its changes, there is a certain language we need to become familiar using. There are good observations & there are bad observations. During the demo: Write down what you see happening. Imagine you were trying to explain this to someone who is not present in the room.

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In science we can take two different kinds of observations: Qualitative Quantitative Two Types of Measurements Qualitative (think quality): observations using words Quantitative (think quantity): observations using numbers and units.

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Heres what I am hoping to see Qualitative observations: States of matter Color Texture Smell Viscosity Quantitative observations: Amount of substances present Step by step procedure!

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Heres what I dont want to see Opinionated language I feel I like Non-specific wording sort of, lots of, kinda Descriptions that sound like a kindergartener wrote them It was all bouncy and describing something as chunky

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Taking Measurements in Chemistry Ch. 2 The SI or Metric System

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The SI System Around 1793, scientists all over the world began to agree upon a single measurement system called Le Systeme International d Unites or SI System 7 base units The idea was to create a unifying system of weights and measurements

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QuantityUnitSymbol LengthMeterm MassKilogramkg TimeSeconds TemperatureKelvinK Amount of a substance Molemol Electric current AmpereA Luminous intensity Candelacd Crash Course: Units Wheres volume??

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m D V Combinations of base units Volume: amount of space taken up by an object Derived SI unit is cubic meter, m 3 More often we use cm 3 = mL Density: ratio of mass to volume g/cm 3 of g/mL or g/L Does not change for a given substance Derived Units m D V D = m V D = m V

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Other Derived Units QuantityUnitSymbolDerivation Areasquare meterm2m2 Length x width Molar Massgrams per mole g/molMass / amount EnergyjouleJForce x length

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Metric Prefix SymbolMeaningScientific Notation megaMMillion / 1,000,0001 x 10 6 kilokThousand / 1,0001 x 10 3 hectahHundred / 1001 x 10 2 dekadaTen / 101 x 10 1 Base Unit decidTenth /.11 x 10 -1 centicHundredth /.011 x 10 -2 millimThousandth /.0011 x 10 -3 microMillionth /.000 0011 x 10 -6 nanonBillionth /.000 000 0011 x 10 -9 picopTrillionth /.000 000 000 0011 x 10 -12 Larger quantities

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Worlds Roundest Object Challenging foundations of the SI System The worlds roundest object hopes to solve the longest running problem in measurement how to define the kilogram. Check out this video! (12 min) Check out this video!

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Using SI prefixes: Number Line Method Conversions from one SI prefix to another (within 1 of the 7 base units) can easily be preformed by moving the decimal place of a quantity by 1 space or 3, left or right.

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Practice Problems 1. 5.6 cm to m 2. 56 mg to g 3. 340 mm to cm 4. 1.2 ML to L 0.056 m 0.056 g 34 cm 1,200,000 L

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Using SI prefixes: Factor-Label Method (Dimensional Analysis) Method requires translating two equal quantities into a ratio or conversion factor Ex: 16 oz = 1 lb can be written 16 oz or 1 lb 1 lb 16 oz Notice: a conversion factor can be represented 2 ways! This can be done with any 2 equal quantities 2 grand slams = 8 R.B.I.s 1 fortnight = 14 days 100 cm = 1 m

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Using SI prefixes: Factor Label Method Using the factor label method to solve problems Ex: How many dimes are in 14 dollars? 1. Write the given 2. Write conversion factor 3. Solve, crossing out units that have divided out 14 dollars x 10 dimes = 14o dimes 1 dollar

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Using Factor-Label Method Sample Problems: Converting 9.8 g to kg 9.8 g x 1 kg = 0.0098 kg 1000. g Converting 9.8 kg to g 9.8 kg x 1000. g = 9800 g 1 kg 1 goes in front of larger unit!

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Practice Problems Try these practice problems, but now using the Factor- Label Method (I realize this seems like more work than the number line methodbut theres a reason why we have to learn this) 1. 5.6 cm to m 2. 1.2 L to ML 3. 100 mm to cm 4. 25 kg of water to mL 0.056 m 1.2 x 10 -6 ML 10 cm 2500 mL

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Do Now: Test your Metric System With-it-ness For each of the measurements on your worksheet, decide the appropriate quantity that should be assigned to it.

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Density Practice Density Formula Use Density Pyramid as a short cut m D V m D V D = m V

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Taking Measurements in Chemistry Accuracy vs. Precision

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Accuracy & Precision in Measurements Accuracy: closeness of measurements to correct value Precision: closeness of a set of measurements to each other (assuming theyre made in the same way)

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High accuracy High precision Low accuracy High precision Low accuracy Low precision

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Accuracy vs. Precision Example: A student measures the density of a sample of nickel. The density of nickel is 8.9 g. mL -1 So the results were: Precise, but not accurate Density Result (g.mL -1) Trial 17.8 Trial 27.7 Trial 37.8

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Accuracy & Precision (continued) Some error always exists in measurements Skill of measurer Conditions of measurements Limitation of instruments

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Percentage Error Accuracy of an individual value (or average) can be compared to the correct/accepted value % Error = Experimental Accepted x 100 Accepted

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Percentage Error What is the percentage error for a mass measurement of 17.7 g, given that the correct value is 21.2 g? A volume is measured experimentally as 4.26 mL. What is the percentage error, given that the accepted value is 4.15 mL?

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Taking Measurements in Chemistry Significant Figures

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Exploring Uncertainty and Precision The Paper Clip Activity Measuring always involves some degree of estimation (i.e. uncertainty) Ruler #3 required the least amount of estimation because instrument had greater precision (more markings)

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Significant Figures Certain digits: digits that represent a marking on a scale or non-blinking number of a display Uncertain (estimated) digits: digits that represents the space between the marks on a scale or the blinking number on a display Sig Figs all digits of certainty + 1 estimated

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Sig Figs: Using the Pacific/Atlantic Rule Step 1: Ask yourself: is the decimal point present or absent? Step 2: Determine which way to start counting If the decimal point is present, start counting from the LEFT If the decimal point is absent, start counting from the RIGHT PACIFICPACIFIC ATLANTICATLANTIC resent bsent

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Pacific/Atlantic Rule Step 3: Start counting on Pacific or Atlantic side from the first NON-ZERO number. Count all numbers after the first non-zero number including zeros.

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Pacific/Atlantic Rule Examples: a) 1234 = ________ sig figs b) 1204 = ________ sig figs c) 0.00234 = _______ sig figs d) 1230 = ______ sig figs e) 1234.0 = ______ sig figs 4 4 3 3 5 Absent Present

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Pacific/Atlantic Rule Examples: a) 1234 = ________ sig figs b) 1204 = ________ sig figs c) 0.00234 = _______ sig figs d) 1230 = ______ sig figs e) 1234.0 = ______ sig figs 3 certain digits indicated by lines on measuring device ; 1 estimated digit - in between lines 3 certain ; 1 estimated 2 certain ; 1 estimated (zero is a place holder) 2 certain ; 1 estimated (zeros are place holders) 4 certain ; 1 estimated 5 3 3 4 4

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Using Sig. Figs. In Calculations Addition/Subtraction Rule Answer should contain least # of decimal places Multiplication/Division Rule Answer should contain least sig figs.

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Do Now: Precision of Lab Instruments 1. Record the following quantities to the correct number of decimal places. ________ L ________ mL _______ o C 2. Convert your answer in A to milliliters: ________ mL 3. Add your answer from A & B. Record using correct sig. figs. ________ mL

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Scientific Notation Some numbers are very large or very small, so we need a short hand notation. 602,200,000,000,000,000,000,000 6.022 x 10 23 0.0000000000000000000000199 1.99 x 10 -23 Too large: Too small:

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Scientific Notation N x 10 n N is a number between 1 and 10 n is a positive or negative integer if n is a negative number, the full number is a small decimal if n is a positive number, the full number is a large number 3.69 x 10 -4 ________________ 1.245 x 10 5 ________________

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Taking Measurements in Chemistry According to the Scientific Method

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The Scientific Method Scientific Method: logical approach to solving problems by a. Observing & collecting data b. Formulating hypotheses c. Testing hypotheses d. Formulating theories e. Publishing results

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Remember: observations about matter can be categorized in two groups: Qualitative Data Quantitative Data Two Types of Measurements Qualitative (think quality): observations using words Quantitative (think quantity): observations using numbers and units.

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Studying a System System: specific portion of matter in a given region of space that has been selected for study Microscopic or macroscopic level Variable: any condition that changes during an experiment Independent: value being manipulated Dependent: result

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Studying a System Experimental Control: conditions that remain constant throughout (i.e. dont change) Often many controlled portions of system Model: Explanation of how phenomena occur and how data or events are related Visual Verbal Mathematical Ex: atomic model of matter

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Studying a System Theory: broad generalization that explains a body of facts or phenomena Used to predict results of new experiments Ex: kinetic molecular theory

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Taking Measurements in Chemistry Graphing Measurements

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Amount of Fertilizer (g) Plant Growth (cm) 65 99 1517 2322 Fertilizer Growth Direct Relationship Independent Variable Dependent Variable Title Appropriate scale Axis labeled Best fit line

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Direct Relationships When 2 quantities divided by each other gives a constant value K (constant value) = Y/X Ex: Density

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Inverse Relationships When 2 quantities multiplied by each other gives a constant value K = X Y Ex: Boyles Law K = PV