CHE 105 Chapter 1 & 2 - Matter and Measurement PDF

Chapter 1 Matter and Measurement (+ Section on Energy) Dr. Mahmoud Mohsin Spring Semester 2025 Copyright 2022 © McGraw Hill LLC. All rights reserved. No reproduction or distribution without the prior written consent of McGraw Hill LLC.  CHAPTER GOALS...

Chapter 1 Matter and Measurement (+ Section on Energy) Dr. Mahmoud Mohsin Spring Semester 2025 Copyright 2022 © McGraw Hill LLC. All rights reserved. No reproduction or distribution without the prior written consent of McGraw Hill LLC.  CHAPTER GOALS Chapter Applications In this chapter you will learn how to: 1 Describe the three states of matter. 2 Classify matter as a pure substance, mixture, element, or compound 3 Report measurements using the metric units of length, mass, and volume 4 Use significant figures. 5 Use scientific notation for very large and very small numbers. 6 Use conversion factors to convert one unit to another. 7 Convert temperature from one scale to another. 8 Define density and specific gravity and use density to calculate the mass or volume of a substance. 2  1.1 Chemistry—The Science of Everyday Experience Chemistry is the study of matter—its composition, properties, and transformations. Matter is anything that has mass and takes up volume. Matter can be: 1-3  1.2 States of Matter Physical properties can be observed or measured without changing the composition of the material. boiling point color melting point odor solubility state of matter 1-4  1.2 States of Matter A physical change alters the material without changing its composition. Physical changes will be covered in more detail in Chapter 7. 1-5  1.2 States 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 (Chapters 5 and 6). 1-6 Activity: Sublimation of Dry Ice (CO2) CO2(s) → CO2(g) Draw a molecular-level representation for the sublimation of CO2. Before After 1-7 Physical State Transitions (know these) 1-8  Chemical Changes A chemical change is a process where one or more substances are converted into one or more new substances. (Also called a chemical reaction.) Examples: – Pennies tarnishing – Burning gasoline – The reaction of hydrogen and oxygen to form water 1-9 Activity: Physical vs. Chemical Changes physical chemical physical chemical 1-10  Chemical Properties Chemical properties are descriptions of the ability of a substance to undergo a chemical change. Examples: – Hydrogen burns easily with oxygen – Helium is unreactive – Iron rusts – Silver tarnishes – Gold is very unreactive 1-11  1.3 Classification of Matter All matter can be classified as either a pure substance or a mixture. Pure Substances: A pure substance is composed of only a single component (atom or molecule). It has a constant composition, regardless of sample size or origin of sample. It cannot be broken down to other pure substances by a physical change. 1-12  1.3 Classification of Matter All matter can be classified as either a pure substance or a mixture. Mixtures: Mixtures are composed of more than one component. They can have varying composition (any combination of solid, liquid, and gas). Mixtures can be separated into their components by a physical process. 1-13  1.3 Classification of Matter All matter can be classified as either a pure substance or a mixture. I. Pure Substances Table sugar (C12H22O11) and water (H2O) are both pure substances: 1-14  1.3 Classification of Matter All matter can be classified as either a pure substance or a mixture. Mixtures Sugar dissolved in water is a mixture. 1-15  1.3 Classification of Matter A pure substance is classified as an element or a compound. An element is a pure substance that cannot be broken down by a chemical change. A compound is a pure substance formed by chemically joining two or more elements. 1-16  1.3 Classification of Matter Figure 1.5 Elements and Compounds 1-17  Composition of Matter Matter – anything that occupies space and has mass Pure Substances – have Mixtures – are composed uniform (the same) of two or more pure chemical composition substances and may or throughout and from may not have uniform sample to sample composition 1-18  Pure Substances Pure Substances: have the same Pure composition throughout, Substances and from sample to sample. Elements Compounds can be further classified as either elements or compounds. 1-19 The Periodic Table 1-20 Activity: Classification of Matter Identify the nonmetals. Explain the characteristics you considered in making your decision. 1-21 Activity Solution: Classification of Matter Metals can be distinguished from nonmetals by the luster and ability to conduct electricity. Since we do not know how each of the elements conducts electricity, we need to use luster as our measure. Nonmetals are usually dull, with the exception of carbon (as diamond). Elements that are gases at room temperature are also nonmetals. 1-22 Activity Solution: Classification of Matter Identify the nonmetals. Phosphorus, bromine, carbon, and sulfur. 1-23 Elements and Their Symbols Element symbols often consist of one or two letters of the element’s name. Examples: carbon: C calcium: Ca How do we explain that Fe is the symbol for iron? Symbols of Selected Elements Englis Original Symb English Original Symbol h Name ol Name Name Name potassiu kalium K copper cuprum Cu m gold aurum Au silver argentu Ag m iron ferrum Fe sodium natrium Na lead plumbum Pb mercu hydrargyru Hg tin stannum Sn 1-24 tungsten wolfram W  Water, H2O, Is a Compound Water, H2O, can be broken down by a chemical process, known as electrolysis, to its elements H2 and O2. The hydrogen (left) and oxygen (right) can be seen bubbling to the top of the 1-25 tubes.  Activity: Elements and Compounds Identify each of the following as an element or compound. 1. He element 2. H2O compound 3. sodium chloridecompound 4. copper element 1-26  Mixtures A mixture is a combination of two or more elements or compounds. Mixtures differ from pure compounds in that their components can be separated by physical processes. Examples: – Iron filings in sand – Salt water – Air 1-27 Salt Being Separated by Evaporation 1-28  Mixtures Mixtures can be Mixtures – further classified consist of 2 or as homogeneous more pure and substances heterogeneous. Homogeneous Heterogeneous Mixtures Homogeneous Mixtures – do (solutions) - have not have uniform mixtures have uniform composition composition the same throughout throughout composition throughout. Heterogeneous 1-29  Activity: Mixtures Classify each of the following mixtures as homogeneous or heterogeneous: Salt water homogeneous* Lake water homogeneous* Tap water Homogeneous Air homogeneous* Brass (an alloy of Cu and Zn) Homogeneous Potting soil Homogeneous Cake mix homogeneous *depends on presence or absence of suspended solids 1-30 Classification of Matter 1-31 Different Ways to Represent Water 1-32  Activity: Classification Classify each of the following as an element, compound, or mixture. Element Compou Mixture of nd elements Mixture of element Mixture of and compounds compoun d 1-33  States of Matter Characteristics of the Physical States of Matter Solid Liquid Gas fixed shape shape of container shape of container (may or may not fill it) (fills it) its own volume its own volume volume of container no volume change slight volume change large volume under pressure under pressure change under pressure particles are fixed in particles are randomly particles are widely place and tend to be arranged and free to separated and move in a regular move about until they independently of (crystalline) array bump into one one another another 1-34 Symbols Used in Chemistry Name Symbol helium He(g) chlorine Cl2(g) silver Ag(s) water H2O(l) carbon CO2(g) dioxide methane CH4(g) (natural gas) 1-35  1.4 Measurement Every measurement is composed of a number and a unit. The number is meaningless without the unit. Examples: proper aspirin dosage = 325 (milligrams or pounds?) a fast time for the 100-meter dash = 10.00 (seconds or days?) 1-36  1.4 Measurement The Metric System Each type of measurement has a base unit in the metric system. Table 1.1 The Basic Metric Units Metric Table divided into Quantity Base three columns Unit Symbol summarizes four metric units. The Length column headersMeter m to are marked from left right as: Quantity, metric base unit, and Mass symbol. Gram g Volume Liter L Time Second s 1-37  1.4 Measurement Table 1.2 Common Prefixes Used for Metric Units Prefix Factor Symbol giga 109 G mega 106 M kilo 103 k deci 10−1 d centi 10−2 c milli 10−3 m micro 10−6 μ nano 10−9 n 1-38  1.4 Measurement B. Measuring Length The base unit of length is the meter (m). 1 kilometer (km) = 1,000 meters (m) 1 km = 1,000 m 1 millimeter (mm) = 0.001 meters (m) 1 mm = 0.001 m 1 centimeter (cm) = 0.01 meters (m) 1 cm = 0.01 m 1-39  1.4 Measurement C. Measuring Mass Mass is a measure of the amount of matter in an object. Weight is the force that matter feels due to gravity. The base unit of mass is the gram (g). 1 kilogram (kg) = 1,000 grams (g) 1 kg = 1,000 g 1 milligram (mg) = 0.001 grams (g) 1 mg = 0.001 g 1-40  1.4 Measurement D. Measuring Volume The base unit of volume is the liter (L). 1 kiloliter (kL) = 1,000 liters (L) 1 kL = 1,000 L 1 milliliter (mL) = 0.001 liters (L) 1 mL = 0.001 L Volume = Length × Width × Height = cm × cm × cm cm 3 1mL 1cm 3 1cc 1-41  1.4 Measurement Table 1.4 English Units and Their Metric Equivalents Quantity English Unit Metric-English Relationship Table Lengthdivided into three columns 1 ft = 12summarizes in. English units and 2.54 cmtheir metric equivalents. The = 1 in. column headers are marked from left to right as: Quantity, English unit, and metric-English relationship. Length 1 yd = 3 ft 1 m = 39.4 in. Length 1 mi = 5,280 ft 1 km = 0.621 mi Mass 1 lb = 16 oz 1 kg = 2.20 lb Mass 1 ton = 2,000 lb 454 g = 1 lb Mass 28.3 g = 1 oz Volume 1 qt = 4 cups 946 mL = 1 qt Volume 1 qt = 2 pints 1 L = 1.06 qt Volume 1 qt = 32 fl oz 29.6 mL = 1 fl oz Volume 1 gal = 4 qt Common abbreviations for English units: inch (in.), foot (ft), yard (yd), mile (mi), pound (lb), ounce(oz), gallon (gal), quart (qt), and fluid ounce (fl oz). 1-42  Volume Units Volumes of liquids are usually measured in units of milliliters (mL). 1 mL = 1 cm3 exactly How many mL in 1 L? 1000 mL = 1L Some 250-mL, 500-mL, and 1-L containers 1-43  Activity: Volume Unit Conversions Convert 25.0 mL to L. 1L 25.0 mL × =0.0250L 1000 mL Convert 25.0 mL to quarts (1 L = 1.057 qt) 1.057qt 0.0250 L × =0.0264qt 1L 1-44  1.5 Significant Figures 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 1-45  Significant Figures Example Significant figures - all digits in a number representing data or results that are known with certainty plus one uncertain digit © McGraw Hill LLC  1.5 Significant Figures A. 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 1-47  1.5 Significant Figures A. Determining Significant Figures Rules for Zero: 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 1-48  1.5 Significant Figures A. Determining Significant Figures Rules for Zero: 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 1-49  Activity: Significant Figures Determine the number of significant figures in each of the following values. # of significant Given number 26 digits 2 19628. 7 00 0.0034 4 916 × 1 10 19 1.2407661 × 8 102 1-50  1.5 Significant Figures A. 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. 1-51  1.5 Significant Figures B. Rules for Multiplication and Division 63.854545 miles 64 miles  hour hour first digit to be dropped 2 sig. figures Answer If the first digit Then: to be dropped is: drop it and all remaining between 0 and 4 digits between 5 and 9 round up the last digit to be retained by adding 1 1-52  1.5 Significant Figures B. Rules for Multiplication and Division Table 1.5 Rounding Off Numbers Original Rounded To The First Rounded Number Number Table divided into three columns summarizes to Be off numbers. rounding Number The Dropped column headers are marked from left to right as: Original number, rounded to, and rounded number. 61.2537 Two places 2 61 61.2537 Three places 5 61.3 61.2537 Four places 3 61.25 61.2537 Five places 7 61.254 The first number to be dropped is indicated in red in each original number and again in the third column of the table. When this number is 4 or fewer, it and all other digits to its right are dropped. When this number is 5 or greater, 1 is added to the digit to its left. 1-53  1.5 Significant Figures C. 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 1-54  1.6 Scientific Notation In scientific notation, a number is written as: x Exponent: y 10 Any positive Coefficient: or negative A number between whole number. 1 and 9 1-55  1.6 Scientific Notation HOW TO Convert a Standard Number to Scientific Notation Example Convert these numbers to scientific notation. 2,500 0.036 Step Move the decimal point to give a number between 1 and 10. 2500 0.036 Step Multiply the result by 10 xwhere x = number of places the decimal was moved. move decimal left, move decimal right, x is positive x is negative 2 2.5 10 3 3.6 10 1-56  1.6 Scientific Notation Converting a Number in Scientific Notation to a Standard Number When the exponent x is positive, move the decimal point x places to the right. When the exponent x is negative, move the decimal point x places to the left. 2.80 10 2 0.0280 1-57  Scientific Notation Math Toolbox Direction Normal Scientific Decimal Notation Notation Moved 3245 Left 3.245 × 103 0.000003245 Right 3.245 × 106 3,245,000,000 Left 3.245 × 109 0.0050607 Right 5.0607 × 103 88 Left 8.8 × 101 2.45 Neither 2.45 × 100 1-58  Activity: Math Toolbox Convert the values shown in decimal form to scientific notation. Decimal Scientific Notation Notation 9,000,000,655.00 9.00000065500 × 109 6 0.00000834 8.34 × 10 1.21 1.21 × 100 14.82 1.482 × 101 299,800,000 2.99800000 × 108 1 63 6.3 × 10 1-59  Activity Solutions: Math Toolbox Multiplication and Division of Exponents 3.76 × 1. (6.78 × 10 9)(5.55 × 104 3 6) 10 = 0 2. (2.99 × 10 )(4.03 × 10 ) 1.20 × 1014 = 1-60  Activity: Significant Figures Calculate the following: 1. 14.6608 + 12.2 + (1.500000 176.9  102) = 2. (5.5 × 108)(4 × 3 × 10-43 1010) = 6.65 × 1045 1-61  1.7 Problem Solving Using Conversion Factors A. Conversion Factors Conversion factor: A term that converts a quantity in one unit to a quantity in another unit. original desired × conversion factor = quantity quantity Conversion factors are usually written as equalities. 2.20 lb = 1 kg To use them, they must be written as fractions. 2.20lb 1 kg or 1kg 2.20lb 1-62  1.7 Problem Solving Using Conversion Factors B. Solving a Problem Using One Conversion Factor If a unit appears in the numerator in one term and the denominator in another term, the units cancel. The goal in setting up a problem is to make sure all unwanted units cancel. To convert 130 lb into kilograms: 130 lb × conversion factor = ? kg original desired quantity quantity 1-63  1.7 Problem Solving Using Conversion Factors Solving a Problem Using One Conversion Factor 2.20lb Answer 1kg 2 sig. figures 130 lb × or = 59 kg 1kg 2.20 lb The bottom conversion factor has the original unit in the denominator. The unwanted unit lb cancels. The desired unit kg does not cancel. 1-64  1.7 Problem Solving Using Conversion Factors HOW TO Solve a Problem Using Conversion Factors Example How many grams of aspirin are in a 325-mg tablet? Identify the original quantity and the desired Step quantity, including units. original quantity desired quantity 325 mg ?g 1-65  1.7 Problem Solving Using Conversion Factors HOW TO Solve a Problem Using Conversion Factors Step Write out the conversion factor(s) needed to solve the problem. 1 g = 1000 mg This can be written as two possible 1000mg fractions: 1g or 1g 1000mg Choose this factor to cancel the unwanted unit, mg. 1-66  1.7 Problem Solving Using Conversion Factors HOW TO Solve a Problem Using Conversion Factors Step Set up and solve the problem. 1g 325 mg 0.325g  1000 mg  3sig.figures 3 sig. figures Unwantedunit cancels Step Write the answer with the correct number of significant figures. 1-67  1.7 Problem Solving Using Conversion Factors C. Solving a Problem Using Two or More Conversion Factors Always arrange the factors so that the denominator in one term cancels the numerator in the preceding term. How many liters are 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 2pt 1qt 1.06qt 1L or or 1qt 2pt 1L 1.06qt Choose the right side Choose the right side one to cancel pt. one to cancel qt. 1-68  1.7 Problem Solving Using Conversion Factors C. Solving a Problem Using Two or More Conversion Factors Set up the problem and solve: 1 qt 1L 1.0 pt   0.471698113L 0.47 L 2 pt 1.06 qt 2 sig. figures 2 sig. figures Write the answer with the correct number of significant figures. 1-69  Activity: Dimensional Analysis 1. How many inches are in 2 kilometers? [1 in = 2.54 cm; 100 cm = 1 m; 1000 m = 1 km] 2. What is the volume in cubic centimeters of a 14 lb block of gold? [1 lb = 453.6 g; densityAu = 19.3 g/cm3] 3. Dan regularly runs a 5-minute mile. How fast is Dan running in feet per second? [1 min = 60 s; 1 mile = 1760 yds; 1 yd = 3 ft] 1-70  Activity Solutions: Dimensional Analysis 1. How many inches are in 2 kilometers? [1 in = 2.54 cm; 100 cm = 1 m; 1000 m = 1 km] First, decide what the problem is asking for: inches. Next, decide what relationships could be used to get to the desired quantity. In this case, the relationships that may be useful are provided in brackets. 1000 m 100 cm 1in 2 km × × × =8×104 in 1 km 1m 2.54 cm 1-71  Activity Solutions: Dimensional Analysis 2. What is the volume in cubic centimeters of a 14 lb block of gold? [1 lb = 453.6 g; densityAu = 19.3 g/cm3] First, decide what the problem is asking for: volume. Next, decide what relationships could be used to get to the desired quantity. In this case, the relationships that may be useful are provided in brackets. 453.6 g 1cm3 2 3 14 lb × × =3.3×10 cm 1 lb 19.3 g 1-72  Activity Solutions: Dimensional Analysis 3. Dan regularly runs a 5-minute mile. How fast is Dan running in feet per second? [1 min = 60 s; 1 mile = 1760 yds; 1 yd = 3 ft] First, decide what the problem is asking for: feet per second. Next, decide what relationships could be used to get to the desired quantity. In this case, the relationships that may be useful are provided in brackets. 1 mile 1 min 1760 yd 3ft × × × =2×101 ft/s 5 min 60s 1 mile 1 yd 1-73  1.9 Temperature Temperature is a measure of how hot or cold an object is. Three temperature scales are used: Degrees Fahrenheit  F  Degrees Celsius  C  Kelvin (K) To convert from  C to  F : To convert from  F to  C :  F 1.8  C   32  F  32 C  1.8 To convert from  C to K: To convert from K to  C: K  C  273  C K  273 1-74  1.9 Temperature Comparing the Three Temperature Scales 1-75  1.10 Density and Specific Gravity A. Density 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 To convert mass mass (g): (g) to volume (mL): g mL mL  g g mL mL g density is g/mL inverse of density is mL/g 1-76  1.10 Density and Specific Gravity A. Density Example: 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 or 1 mL 1.05 g Choose the right side one to cancel the unwanted unit, g. 1-77  1.10 Density and Specific Gravity A. Density Set up and solve the problem: 2 sig. figures 1 mL 5.0 g  4.761904762 mL 4.8 mL 1.05 g cancel g 2 sig. figures Write the final answer with the correct number of significant figures. 1-78  Densities of Common Substances Densities of Some Common Substances Substance Physical State Density (g/mL) helium gas 0.000178 oxygen gas 0.00143 cooking oil liquid 0.92 water liquid 1.00 mercury liquid 13.6 gold solid 19.3 copper solid 8.92 zinc solid 7.14 ice solid 0.92 1-79  Densities of Various Liquids This cylinder contains: – Antifreeze – Corn oil – Dish detergent – Maple syrup – Shampoo – Water Which layer is which? 1-80  1.10 Density and Specific Gravity B. Specific Gravity Specific gravity: A quantity that compares the density of a substance with the density of water at the same temperature.  g  density of asubstance    mL  specific gravity   g  density of water    mL  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 but contains no units. 1-81  Energy and Energy Changes Energy – is the capacity to do work or to transfer heat Two main forms of energy are: – Kinetic energy: the energy of motion – Potential energy: energy possessed by an object because of its position Other energies are forms of kinetic and potential energy (chemical, mechanical, electrical, heat, etc.) 1-82 Energy and Energy Changes When chemical or physical changes occur, energy changes also occur. Some processes release energy and some require an energy input. Examples: – When wood burns with oxygen, energy in the form of heat is released. – When ammonium nitrate dissolves in water in a cold pack, energy in the form of heat is absorbed. 1-83  When hydrogen burns with oxygen, energy in the form of heat is released. 1-84  Electricity is used to decompose water into its elements. 1-85 Electrical energy is used to run electric vehicles. 1-86  Energy Kinetic energy – energy of motion – The kinetic energy of a sample will increase as temperature is increased. Potential energy – energy possessed by an object because of its position; stored energy – As a ball is raised up in the air, its potential energy increases. – Very reactive substances have high potential energy. 1-87 Kinetic and Potential Energy 1-88 Which pair of molecules has more kinetic energy? Answer: Image A 1-89  End of Main Content End of Chapter One Copyright 2022 © McGraw Hill LLC. All rights reserved. No reproduction or distribution without the prior written consent of McGraw Hill LLC.