Chapter 1 - Matter and Measurement PDF

Chapter 1 1 Matter and measurement OK, Let’s Begin!  Chemistry: The Science of Everyday Experience 2 Chemistry is the study of matter—its composition, properties, and transformations. Matter is anything that has mass and takes up volume. 1. Naturally occur...

Chapter 1 1 Matter and measurement OK, Let’s Begin!  Chemistry: The Science of Everyday Experience 2 Chemistry is the study of matter—its composition, properties, and transformations. Matter is anything that has mass and takes up volume. 1. Naturally occurring matter: 2. Synthetic (human-made) matter: Cotton Nylon Silk Polyester Hair Styrofoam Sand Ibuprofen Gemstones Many antibiotics digoxin, a cardiac drug Chemistry and Matter Matter is another word for all substances that make up our world. Antacid tablets are matter. Water is matter. Glass is matter. Air is matter.  Atoms and matter 4 NaCl crystal Atoms of Na and Cl  Chemicals are substances that have the same composition and properties wherever found. often substances made by chemists that you use every day. Toothpaste is a combination of many chemicals. Chemicals in the Kitchen Study Check Which of the following contains chemicals? A. Sunlight B. Fruit C. Milk D. Breakfast cereal Solution Which of the following contains chemicals? A. Sunlight is energy; it does not contain chemicals. B. Fruit is a type of matter; it contains chemicals. C. Milk is a type of matter; it contains chemicals. D. Breakfast cereal is a type of matter; it contains chemicals. Only B. fruit, C. milk, and D. breakfast cereal contain chemicals. 1.1 Scientific Method Thinking Like a Scientist Linus Pauling won the Nobel Prize in chemistry in 1954. Learning Goal Describe the activities that are part of the scientific method. The Scientific Method The scientific method is a set of general principles that helps to describe how a scientist thinks. 1. Make observations about nature and ask questions about what you observe. 2. Propose a hypothesis, which states a possible explanation of the observations. 3. Several experiments may be done to test the hypothesis. 4. When results of the experiments are analyzed, a conclusion is made as to whether the hypothesis may be true or false. Discovery of Penicillin Using the Scientific Method Suppose you visit a friend in her home and soon after you arrive, you begin to sneeze. You observe that your friend has a new cat. You ask yourself why you are sneezing and form a hypothesis that you are allergic to cats. You perform experiments to test your hypothesis by visiting other friends with cats. If you sneeze after leaving the other homes with cats, you come to the conclusion that your hypothesis is correct. Using the Scientific Method Through your observations, you may determine that you are allergic to cat hair and dander. Using the Scientific Method The hypothesis is modified if the results of the experiments do not support it. The scientific method develops conclusions using observations, hypotheses, and experiments. Study Check Identify each of the following as an observation (O), a hypothesis (H), an experiment (E), or a conclusion (C): A. During your visit to the gym, your trainer records that you ran for 25 minutes on the treadmill. B. Scientific studies show that exercising lowers blood pressure. C. Your doctor thinks that your weight loss is due to increased exercise. Solution Identify each of the following as an observation (O), a hypothesis (H), an experiment (E), or a conclusion (C): A. During your visit to the gym, your trainer records that you ran for 25 minutes on the treadmill. Observation (O) B. Scientific studies show that exercising lowers blood pressure. Conclusion (C) C. Your doctor thinks that your weight loss is due to increased exercise. Hypothesis (H) 1.2 Matter The Classification of Matter (elements and compounds, atoms and molecules, pure and impure, states of matter, physical and chemical properties)  States of Matter Three States of Matter—Solid, Liquid, and Gas 18 1. Solid: a) has a definite volume b) maintains its shape regardless of its container c) has particles that lie close together in a regular d) three-dimensional array  States of Matter Three States of Matter—Solid, Liquid, and Gas 2. Liquid: a) has definite volume b) takes the shape of its container c) has particles that are close together but can move past one another 19  States of Matter Three States of Matter—Solid, Liquid, and Gas 3. Gas: a) has no definite shape or volume b) expands to fill the volume and assumes the shape of whatever container it is put in c) has particles that are very far apart and move around randomly 20 21 States of Matter… Properties of Matter Physical change alters the material without changing its composition. melting ice (solid water) to form liquid water boiling liquid water to form steam (gaseous water) 22 States of Matter… Properties of Matter Chemical properties determine how a substance can be converted into another substance. Chemical change is the chemical reaction that converts one substance into another. a piece of paper burning metabolizing an apple for energy oxygen and hydrogen combining to form water 23  Practise question: 24 Select one: Which is an example of physical change? 1. The rusting of an iron nail. 2. The burning of propane in a gas grill. 3. Baking cookies 4. Polishing tarnished silver 5. Melting of an ice cube in a glass of soda  States of Matter Properties of Matter 25 Physical properties can be observed or measured without changing the composition of the material. boiling point melting point solubility color odor Classification of Matter A Pure Substance: 1. Composed of a single component. 2. Has a constant composition, regardless of sample size and origin of sample. 3. Cannot be broken down to other pure substances by a physical change. 4. Diamonds, sugar (C12H22O11) and water (H2O) are all pure substances. 26 Classification of Matter: Element vs. Compound An element is a pure substance that A compound is a pure substance cannot be broken down by a chemical formed by chemically joining change. two or more elements. table salt (NaCl) aluminum metal (Al) 27 Classification of Matter A Mixture (i.e. not pure) 1. is composed of more than one component can have varying composition (any combination of solid, liquid, and gas), depending on the sample. 2. Can be separated into its components by a physical change. 3. Sugar dissolved in water = mixture 28 Homogenous Mixture  The composition is uniform throughout this mixture.  The different parts of the mixture are not visible i.e. copper and zinc in brass. 29  Bronze – a metal alloy 30  Consisting primarily of copper and tin metals.  A homogenous mixture. Think you homogenous mixtures that are liquid or gaseous… Heterogenous Mixture  This is not uniform throughout.  The different parts are visible and can be separated out. E.g. copper and water.  Think of some more heterogenous mixtures… 31  Heterogenous mixtures 32 Classification of Matter - Recap: 33  1.3 Measurements 1.3 Measurements (numbers and units), exact and measured numbers. 34  Measurement So how do we measure matter? The Importance of Units 35  Every measurement is composed of a number and a unit.  The number is meaningless without the unit.  Proper aspirin dosage = 325 (milligrams or pounds?)  A fast time for the 100-meter dash = 10.00 (seconds or days?)  The English system uses units like feet (length), gallons (volume), and pounds (weight).  The Metric system uses units like meters (length), liters (volume), and grams (mass). Measurement The Metric System of Units Each type of measurement has a base unit, which you need to know. 36 Measurement The Metric System of Units Other units are related to the base unit by a power of 10. The prefix of the unit name indicates if the unit is larger or smaller than the base unit. 37 Measurement Measuring Length 1 kilometer (km) = 1,000 meters (m) 1 millimeter (mm) = 0.001 meters (m) 1 centimeter (cm) = 0.01 meters (m) Tour de France 2011 38 3,430 kilometres = 2,130 miles. Measurement Measuring Mass Mass is a measure of the amount of matter in an object. Weight is the force that matter feels due to gravity. 1 kilogram (kg) = 1,000 grams (g) 1 milligram (mg) = 0.001 grams (g) 39 Measurement Measuring Volume 1 kiloliter (kL) = 1,000 liters (L) 1 milliliter (mL) = 0.001 liters (L) Volume = Length x Width x Height = cm x cm x cm = cm3 1 mL = 1 cm3 = 1 cc 40  Significant Figures Exact and Inexact Numbers An exact number results from counting objects or is part of a definition. 10 fingers 10 toes 1 meter = 100 centimeters An inexact number results from a measurement or observation and contains some uncertainty. 15.3 cm 1000.8 g 0.0034 mL 41  Chemistry and Measurement 42  Estimate the length of the piece of wood. 4.5 cm or 4.4 cm or 4.6 cm The digits in red here are estimated so both readings are equally correct.  Chemistry and Measurement 43  Estimate the length of the piece of wood. 4.55 cm or 4.56 cm The digits in red here are estimated so both readings are equally correct.  Chemistry and measurement 44 Which value best represents the temperature indicated on the thermometer? A. 0.48oC B. 4.8oC C. 4.85oC D. 5oC 45 1.4 Expressing numbers Scientific and standard (i.e. decimal) notation  Scientific notation: 46 Useful for very small and very large numbers. Width of a human hair: Diameter of the sun: 0.00031 inches 1,400,000,000 meters  Scientific Notation 47  Why do we need it?  The number of molecules of CO2 exhaled in one breath:  320,000,000,000,000,000,000!  It would not be practical to write this over and over again.  Scientific notation 48  The number of molecules of CO2 exhaled in one breath: 320,000,000,000,000,000,000 Count from the right (in this case) until you reach the 1st two digits that are between 1 and 10. → 3.2 x 10 20 exponent → coefficient  Scientific Notation In scientific notation, a number is written as: yx 10x Exponent: Any positive or negative whole number. Coefficient: A number between 1 and 10. y = the coefficient y x 10 x x = the exponent 49  Scientific Notation 50 y x 10x 101=10 102= 10 x 10 = 100 103= 10 x 10 x 10 = 1000 10-1= 1/10 = 0.1 10-2= 1/ (10 x10) = 0.01 10-3= 1/ (10 x 10 x 10) = 0.001  Rule of thumb: Interpreting Numbers Expressed in Scientific Notation 51  Numbers with negative exponents (10-x) are decimal numbers less than one 3 x 10-5 cm  Numbers with positive exponents (10x) are decimal numbers greater than one 1.28 x 107 m Scientific Notation HOW TO Convert a Standard Number to Scientific Notation Example Convert these numbers to scientific notation. 2,500 0.036 Move the decimal point to give a number Step between 1 and 10. 2500 0.036 Multiply the result by 10x, where Step x = number of places the decimal was moved. move decimal left, move decimal right, x is positive x is negative 2.5 x 103 3.6 x 10−2 52  Converting a Number in Scientific Notation back to a Standard Number 53 When the exponent x is positive, move the decimal point x places to the right. 2.800 x 102 = 280.0 When the exponent x is negative, move the decimal point x places to the left. 2.80 x 10–2 = 0.0280  Scientific notation 54 Value Scientific Notation 3097 3.097 x 103 616000 6.16 x 105 0.000138 1.38 x 10-4 0.12 1.2 x 10-1 1. Which values in the table are greater than one? 2. Which values are less than one?  What is 0.0014 expressed in scientific 55 notation? A: 14 x 10-4 B: 1.4 x 103 C: 1.4 x 10-3 D: 0.14 x 10-2  Express 520000 in scientific 56 notation using 3 Significant Figures 520 x 103 5.20 x 105  What is 1.256 x 103 expressed as 57 a decimal number? A: 1256 B: 0.001256 C: 0.1256  2.4 x 10 3 i.e. 2,400 Enter as: 2.4 Exp (or EE) 3 Now try: 2.4 x 10 -9 i.e. 0.0000000024 58  Scientific notation 59 (2.4 x 10 ) 3 x (4.25 x 10 ) -7 = A. 0.00102 B. 0.0102 C. 0.102 D. 1.0 x 10-3 60 1.5 Expressing a measurement Expressing a measurement with the correct number of digits... significant figures. Adding, subtracting, multiplying and dividing, rounding. Significant Figures Determining Significant Figures Significant figures are all the digits in a measured number including one estimated digit. All nonzero digits are always significant. 65.2 g 255.345 g 3 sig. figures 6 sig. figures 61 Significant Figures Rules to Determine When a Zero is a Significant Figure Rule 1: A zero counts as a significant figure when it occurs: between two nonzero digits 29.05 g 1.0087 mL 4 sig. figures 5 sig. figures at the end of a number with a decimal place 3.7500 cm 620. lb 5 sig. figures 3 sig. figures 62 Significant Figures Rules to Determine When a Zero is a Significant Figure Rule 2: A zero does not count as a significant figure when it occurs: at the beginning of a number 0.00245 mg 0.008 mL 3 sig. figures 1 sig. figure at the end of a number that does not have a decimal 2570 m 1245500 m 3 sig. figures 5 sig. figures 63  Significant Figures – from yesterday 64  Rules for significant figures in measurements:  Inexact numbers: 1. All non-zeroes are significant. 1.234g 2. Leading zeroes are not significant: 0.000054m 3. Trapped zeroes are significant: 20034 L 4. Trailing are significant only if an explicit decimal place is present: 7000 m versus 7000. m  Exact numbers: 1. These by definition have an infinite number of significant figures: You have 10 fingers and 10 toes. 189 cm x 1m/100cm = 1.89m 3 sig figs exact 3 sig figs  Chemistry and measurement Measured value 14 g 0.051 m 101 mL 2.000 g 5000 ft # of sig figs? 1. Are all non-zero digits in a measured value significant? 2. If a zero is at the beginning of a measured value, is it significant? 3. If a zero is between two non-zero digits, is it significant? 4. If a zero is at the end of a measured value, is it significant? Measured value 14 g 0.051 m 101 mL 2.000 g 5000 ft # of sig figs? 1 (or 4 if 2 2 3 4 decimal point) 65 Significant Figures Rules for Multiplication and Division The answer has the same number of significant figures as the original number with the fewest significant figures. 4 sig. figures 351.2 miles 63.854545 miles = 5.5 hour hour 2 sig. figures Answer must have 2 sig. figures. 66 Significant Figures Rules for Rounding Off Numbers to be retained to be dropped 63.854545 miles = 64 miles hour hour first digit to be dropped 2 sig. figures Answer If the first digit to be dropped is: Then: between 0 and 4 drop it and all remaining digits between 5 and 9 round up the last digit to be retained by adding 1 67 Significant Figures Rules for Addition and Subtraction The answer has the same number of decimal places as the original number with the fewest decimal places. 10.11 kg 2 decimal places 3.6 kg 1 decimal place 6.51 kg answer must have 1 decimal place 6.5 kg final answer = 1 decimal place 68  Chemistry and measurements 69 5.75 cm x 2.1 cm = A. 12.075 cm2 calculator answer B. 12.08 cm2 C. 12.1 cm2 D. 12.0 cm2 E. 12 cm2  Chemistry and measurements 70 Comment on the 2 mL value reported by a calculator. 8.00 mL – 6.00 mL = 2 mL should be reported as 2.00 mL  Sample Problem 71  Significant Figures recap 72  6 children x 3 apples = ?  3.4 x 5.29 = 17.986  Roundto 2 significant figures:  Answer = 18  Get the average of 3 measurements: 5.14 m + 5.15 m + 5.1443 m 3  Calculator = 5.144766666666666666666 m  Round answer to 5.14 m 75 1.6 SI system The International System of Units. (SI system), Prefixes with SI units. Equivalents between SI and non Si units. Density. Measurement The Metric System of Units Remember that each type of measurement has a base unit. 76 Measurement The Metric System of Units Other units are related to the base unit by a power of 10. The prefix of the unit name indicates if the unit is larger or smaller than the base unit. 77 2.1 Units of Measurement The metric system is the standard system of measurement used in chemistry. Learning Goal Write the names and abbreviations for the metric or SI units used in measurements of length, volume, mass, temperature, and time. Units of Measurement Scientists use the metric system of measurement and have adopted a modification of the metric system called the International System of Units as a worldwide standard. The International System of Units (SI) is an official system of measurement used throughout the world for units of length, volume, mass, temperature, and time. Units of Measurement: Metric and SI Length: Meter (m), Centimeter (cm) Length in the metric and SI systems is based on the meter, which is slightly longer than a yard. 1 m = 100 cm 1m = 1.09 yd 1 m = 39.4 in. 2.54 cm = 1 in. Volume: Liter (L), Milliliter (mL) Volume is the space occupied by a substance. The SI unit of volume is m3; however, in the metric system, volume is based on the liter, which is slightly larger than a quart. 1L = 1000 mL 1L = 1.06 qt Graduated cylinders are used 946 mL = 1 qt to measure small volumes. Mass: Gram (g), Kilogram (kg) The mass of an object is a measure of the quantity of material it contains. 1 kg = 1000 g 1 kg = 2.20 lb 454 g = 1 lb The SI unit of mass, the kilogram On an electronic balance, the (kg), is used for larger masses. The digital readout gives the mass metric unit for mass is the gram (g), of a nickel, which is 5.01 g. which is used for smaller masses. Temperature: Celsius (°C), Kelvin (K) Temperature tells us how hot or cold something is. Temperature is measured using Celsius (°C) in the metric system. Kelvin (K) in the SI system. Water freezes at 32 °F, or 0 °C. A thermometer is used to measure temperature. The Kelvin scale for temperature begins at the lowest possible temperature, 0 K. Time: Second (s) Time is measured in units such as the following: years (yr) days hours (h) minutes (min) seconds (s) A stopwatch is used to The SI and metric unit of time is the measure the time of a race. second (s). Study Check What are the SI units for the following? A. volume B. mass C. length D. temperature  Solution What are the SI units for the following? A. The SI unit for volume is the cubic meter, m3. B. The SI unit for mass is the kilogram, kg. C. The SI unit for length is the meter, m. D. The SI unit for temperature is the kelvin, K. 88 Temperature - A measure of kinetic energy…  Temperature (you will be given the conversion factors but you need to be able to do the calculation) 89 Temperature is a measure of how hot or cold an object is. Three temperature scales are used: 1. degrees Fahrenheit (oF) 2. degrees Celsius (oC) 3. Kelvin (K) To convert from oC to oF: To convert from oF to oC: oF = 1.8(oC) + 32 oC = oF − 32 1.8 To convert from oC to K: To convert from K to oC: K = oC + 273 oC = K − 273  Temperature Comparing the Three Temperature Scales 90  Temperature 91  Temperature conversion 92 On a summer day the outdoor temperature in Norfolk, VA was 36.5 °C. What was this outdoor temperature in °F? oF = 1.8(oC) + 32  36.7 °C X 1.8 +32 = 98.06 °F  ANSWER: 98.1 °F  Note the significant figures  Temperature 93 The temperature in France is reading 31oC. Which of the following activities is most appropriate? A B  Density 94 The measured relationship between the mass and volume of an object. Mass (g) Density (d) = Volume (ml)  Density 95 An average human brain has a capacity of 1.24 L and a mass of 1.30 kg. What is the density of an average brain? mass Density (d) = volume 1.30 kg Density (d) = 1.24 L Density (d) = 1.048387097 kg/L 1.05 kg/L  Density 96  One thing that does not influence density is the amount of a substance.  Properties which are not influenced by the amount of substance present are called intrinsic properties.  Other Intrinsic properties: Color, Boiling point, melting point.  Extrinsic Properties 97  Properties which are influenced by the amount of substance present are called extrinsic properties.  Other Extrinsic properties: Mass, Volume, Solubility.  Why does ice float on water? 98  The less dense substance will float on top of the denser substance.  True or false?  When a piece of magnesium (density = 1.738 g/mL) is placed in a container of liquid carbon tetrachloride (density = 1.59 g/mL), the piece of magnesium will float on top of the carbon tetrachloride. Density of Solids The density of a solid can be determined by dividing the mass of an object by its volume. Density Using Volume Displacement The density of the solid zinc object is calculated by dividing its mass by its displaced volume. To determine its displaced volume, submerge the solid in water so that it displaces water that is equal to its own volume. 45.0 mL − 35.5 mL = 9.5 mL = 9.5 cm3 Density calculation: Study Check What is the density (g/cm3) of a 48.0–g sample of a metal if the level of water in a graduated cylinder rises from 25.0 mL to 33.0 mL after the metal is added? 33.0 mL 25.0 mL object Solution What is the density (g/cm3) of a 48.0–g sample of a metal if the level of water in a graduated cylinder rises from 25.0 mL to 33.0 mL after the metal is added? 103 1.7 Converting Units 1.7 Converting Units and sig figs in these operations. Conversion factors with exact or measured digits.  Problem Solving Using the Factor-Label Method: Conversion Factors Conversion factor: “A term that converts a quantity in one unit to a quantity in another unit.” original desired x conversion factor = quantity quantity Conversion factors are usually written as equalities. 2.21 lb = 1 kg To use them, they must be written as fractions. 2.21 lb 1 kg or 1 kg 2.21 lb 104 Using the Factor-Label Method Solving a Problem Using One Conversion Factor Factor-label method: Using conversion factors to convert a quantity in one unit to a quantity in another unit. Remember that conversion factor are exact numbers! units are treated like numbers make sure all unwanted units cancel To convert 130 lb into kilograms: 130 lb x conversion factor = ? kg original desired quantity quantity 105 Using the Factor-Label Method Solving a Problem Using One Conversion Factor 2.21 lb 1 kg Answer 130 lb x or 2 sig. figures 1 kg 2.21 lb = 59 kg The bottom conversion factor has the original unit in the denominator. The unwanted unit lb cancels. The desired unit kg does not cancel. 106 Using the Factor-Label Method HOW TO Solve a Problem Using Conversion Factors How many grams of aspirin are in a 325-mg Example tablet? Identify the original quantity and the desired Step quantity, including units. original quantity desired quantity 325 mg ?g 107 Using the Factor-Label Method HOW TO Solve a Problem Using Conversion Factors Write out the conversion factor(s) needed Step to solve the problem. 1 g = 1000 mg This can be written as two possible fractions: 1000 mg 1g or 1g 1000 mg Choose this factor to cancel the unwanted unit, mg. 108 Using the Factor-Label Method HOW TO Solve a Problem Using Conversion Factors Step Set up and solve the problem. 1g 0.325 g 325 mg x = 1000 mg 3 sig. figures 3 sig. figures Unwanted unit cancels. Write the answer with the correct number of significant figures. Step 109  Using the Factor-Label Method Solving a Problem Using Two or More Conversion Factors 110 Always arrange the factors so that the denominator in one term cancels the numerator in the preceding term. How many liters is in 1.0 pint? 1.0 pint ?L original quantity desired quantity Two conversion factors are needed: 2 pints = 1 quart 1.06 quarts = 1 liter 2 pt 1 qt 1.06 qt or 1L or 1 qt 2 pt 1L 1.06 qt First, cancel pt. Then, cancel qt.  Using the Factor-Label Method Solving a Problem Using Two or More Conversion Factors 111 Set up the problem and solve: 1 qt 1L 1.0 pt x x = 0.47 L 2 pt 1.06 qt 2 sig. figures 2 sig. figures Write the answer with the correct number of significant figures.  Density (a conversion factor) 112 Density: A physical property that relates the mass of a substance to its volume. mass (g) density = volume (mL or cc) To convert volume (mL) To convert mass (g) to mass (g): to volume (mL): g mL mL x = g g x = mL mL g density inverse of density  Density: Solving Problems with Density 113 If the density of acetic acid is 1.05 g/mL, what is the volume of 5.0 grams of acetic acid? 5.0 g ? mL original quantity desired quantity Density is the conversion factor, and can be written two ways: 1.05 g 1 mL 1 mL 1.05 g Choose the inverse density to cancel the unwanted unit, g.  Density Solving Problems with Density 114 Set up and solve the problem: 1 mL 5.0 g x = 4.8 mL 1.05 g 2 sig. figures 2 sig. figures Unwanted unit cancels. Write the final answer with the correct number of significant figures.  Density 115  What is the mass in grams of 85.32 mL of blood plasma with a density of 1.03 g/mL? Density = mass / volume  85.32 ml x 1.03 g = 87.8796 grams 1 ml  ANSWER: 87.9 g 116 Muscle d= 1.1 g/cm3 Fat d=0.9 g/cm3  Density and Specific Gravity Specific Gravity 117 Specific gravity: A quantity that compares the density of a substance with the density of water at the same temperature. density of a substance (g/mL) specific gravity = density of water (g/mL) Unitless: The units of the numerator (g/mL) cancel the units of the denominator (g/mL). The specific gravity of a substance is equal to its density since density(water) ~1, but contains no units. What is the density of a sample of rubbing alcohol if it has a specific gravity of 0.789? A. 1.27 g/mL B. 0.789 g/mL C. 1.00 g/mL D. 0.895 g/mL