1 Chapter 1 Lecture Outline Prepared by Andrea D. Leonard University of Louisiana at Lafayette Copyright © McGraw-Hill Education. Permission required for reproduction or display.

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:

3 1.2 States of Matter The Solid State: A solid has a definite volume. It maintains its shape regardless of its container. Solid particles lie close together in a regular pattern.

4 1.2 States of Matter The Liquid State: A liquid has a definite volume. It takes the shape of its container. Liquid particles are close together but can move past one another.

5 1.2 States of Matter The Gas State: A gas has no definite shape; it assumes the shape of its container. It has no definite volume; it assumes the volume of its container. Gas particles are very far apart and move around randomly.

6 1.2 States of Matter Physical properties can be observed or measured without changing the composition of the material. boiling point melting point solubility color odor state of matter

7 A physical change alters the material without changing its composition. 1.2 States of Matter Physical changes will be covered in more detail in Chapter 7.

8 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). 1.2 States of Matter

9 1.3 Classification of Matter 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. I. Pure Substances All matter can be classified as either a pure substance or a mixture.

10 Table sugar (C 12 H 22 O 11 ) and water (H 2 O) are both pure substances: 1.3 Classification of Matter I. Pure Substances All matter can be classified as either a pure substance or a mixture.

Classification of Matter II. 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. All matter can be classified as either a pure substance or a mixture.

Classification of Matter Sugar dissolved in water is a mixture. II. Mixtures All matter can be classified as either a pure substance or a mixture.

13 I. An element is a pure substance that cannot be broken down by a chemical change. 1.3 Classification of Matter A pure substance is classified as an element or a compound.

14 II. A compound is a pure substance formed by chemically joining two or more elements. 1.3 Classification of Matter A pure substance is classified as an element or a compound.

Classification of Matter

16 The number is meaningless without the unit. 1.4 Measurement proper aspirin dosage = 325 (milligrams or pounds?) a fast time for the 100-meter dash = (seconds or days?) Every measurement is composed of a number and a unit. Examples:

17 Each type of measurement has a base unit in the metric system Measurement A. The Metric System

18 Other units are related to the base unit by a power of Measurement A. The Metric System The prefix of the unit name indicates if the unit is larger or smaller than the base unit.

19 1 kilometer (km) = 1,000 meters (m) 1 km = 1,000 m 1 millimeter (mm) = meters (m) 1 mm = m 1 centimeter (cm) = 0.01 meters (m) 1 cm = 0.01 m 1.4 Measurement B. Measuring Length The base unit of length is the meter (m).

20 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 kg = 1,000 g 1 milligram (mg) = grams (g) 1 mg = g 1.4 Measurement C. Measuring Mass The base unit of mass is the gram (g).

21 1 kiloliter (kL) = 1,000 liters (L) 1 kL = 1,000 L 1 milliliter (mL) = liters (L) 1 mL = L Volume = Length x Width x Height = cm x cm x cm = cm 3 1 mL = 1 cm 3 = 1 cc 1.4 Measurement D. Measuring Volume The base unit of volume is the liter (L).

Measurement

23 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 cm g mL 1.5 Significant Figures

24 Significant figures are all the digits in a measured number including one estimated digit. All nonzero digits are always significant. 3 sig. figures6 sig. figures 65.2 g g 1.5 Significant Figures A. Determining Significant Figures 65.2 g g

25 3 sig. figures Rule 1: A zero counts as a significant figure when it occurs: between two nonzero digits 5 sig. figures at the end of a number with a decimal place g 4 sig. figures mL29.05 g mL cm 5 sig. figures 620. lb cm 620. lb 1.5 Significant Figures A. Determining Significant Figures Rules for Zero:

26 Rule 2: A zero does not count as a significant figure when it occurs: 5 sig. figures at the beginning of a number 1 sig. figure at the end of a number that does not have a decimal mg0.008 mL 3 sig. figures mg0.008 mL 2570 m m 3 sig. figures 2570 m m 1.5 Significant Figures A. Determining Significant Figures Rules for Zero:

miles 5.5 hour = miles hour miles 5.5 hour 4 sig. figures 2 sig. figures Answer must have 2 sig. figures. The answer has the same number of significant figures as the original number with the fewest significant figures. 1.5 Significant Figures B. Rules for Multiplication and Division

miles to be retainedto be dropped first digit to be dropped hour If the first digit to be dropped is: Then: between 0 and 4drop it and all remaining digits between 5 and 9 round up the last digit to be retained by adding 1 = 2 sig. figures Answer 1.5 Significant Figures B. Rules for Multiplication and Division 64 miles hour

Significant Figures B. Rules for Multiplication and Division

30 The answer has the same number of decimal places as the original number with the fewest decimal places kg 3.6 kg 6.51 kg kg 3.6 kg 2 decimal places 1 decimal place answer must have 1 decimal place = 6.5 kg final answer 1 decimal place 1.5 Significant Figures C. Rules for Addition and Subtraction

31 y x 10 x Coefficient: A number between 1 and 10. y x 10 x Exponent: Any positive or negative whole number. In scientific notation, a number is written as: 1.6 Scientific Notation

32 HOW TO Convert a Standard Number to Scientific Notation Example Step [1] Step [2] Convert these numbers to scientific notation. 2, Move the decimal point to give a number between 1 and Multiply the result by 10 x, where x = number of places the decimal was moved. move decimal left, x is positive move decimal right, x is negative 2.5 x x 10 −2 1.6 Scientific Notation

33 When the exponent x is positive, move the decimal point x places to the right x 10 –2 = x 10 2 = 1.6 Scientific Notation Converting a Number in Scientific Notation to a Standard Number When the exponent x is negative, move the decimal point x places to the left

34 Conversion factor: A term that converts a quantity in one unit to a quantity in another unit. Conversion factors are usually written as equalities lb = 1 kg To use them, they must be written as fractions. original quantity original quantity conversion factor desired quantity desired quantity x = 1.7 Problem Solving Using Conversion Factors A. Conversion Factors 2.20 lb 1 kg or 1 kg 2.20 lb

35 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 lbx conversion factor = ? kg original quantity desired quantity 1.7 Problem Solving Using Conversion Factors B. Solving a Problem Using One Conversion Factor

36 or 2.20 lb 1 kg 2.20 lb 1 kg 2.20 lb 1 kg 2.20 lb 130 lbx =59 kg Answer 2 sig. figures The bottom conversion factor has the original unit in the denominator. The unwanted unit lb cancels. The desired unit kg does not cancel. 1.7 Problem Solving Using Conversion Factors B. Solving a Problem Using One Conversion Factor

37 HOW TO Solve a Problem Using Conversion Factors Example How many grams of aspirin are in a 325-mg tablet? Step [1] Identify the original quantity and the desired quantity, including units. original quantity desired quantity 325 mg? g 1.7 Problem Solving Using Conversion Factors

38 Step [2] Write out the conversion factor(s) needed to solve the problem. 1 g = 1000 mg HOW TO Solve a Problem Using Conversion Factors This can be written as two possible fractions: or Choose this factor to cancel the unwanted unit, mg mg 1g 1000 mg 1.7 Problem Solving Using Conversion Factors

39 Step [3] HOW TO Solve a Problem Using Conversion Factors Set up and solve the problem. 325 mgx 1 g 1000 mg = g Step [4] Write the answer with the correct number of significant figures. Unwanted unit cancels 325 mg g 3 sig. figures 1.7 Problem Solving Using 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? 2 pints = 1 quart1.06 quarts = 1 liter Two conversion factors are needed: 1.0 pint original quantity ? L desired quantity 1.7 Problem Solving Using Conversion Factors C. Solving a Problem Using Two or More Conversion Factors 2 pt 1 qt 2 pt or 1.06 qt 1 L 1.06 qt or First, cancel pt. Then, cancel qt.

41 Set up the problem and solve: 1.0 ptx 1 qt 2 pt x 1 L 1.06 qt = L 1.0 pt 2 sig. figures 0.47 L 2 sig. figures Write the answer with the correct number of significant figures. 1.7 Problem Solving Using Conversion Factors C. Solving a Problem Using Two or More Conversion Factors

Temperature Temperature is a measure of how hot or cold an object is. 1.Degrees Fahrenheit ( o F) 2.Degrees Celsius ( o C) 3.Kelvin (K) Three temperature scales are used: To convert from o C to o F: To convert from o F to o C: o F = 1.8( o C) + 32 o C = o F − o C = o F − To convert from o C to K: K = o C o C = K − 273 To convert from K to o C:

Temperature Comparing the Three Temperature Scales

Density and Specific Gravity A. Density density= mass (g) volume (mL or cc) Density: A physical property that relates the mass of a substance to its volume. To convert volume (mL) to mass (g): To convert mass (g) to volume (mL): mLx g =g gx g = density inverse of density

Density and Specific Gravity A. Density 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 original quantity ? mL desired quantity Density is the conversion factor, and can be written two ways: 1.05 g 1 mL 1.05 g Choose the inverse density to cancel the unwanted unit, g. Example:

Density and Specific Gravity A. Density Set up and solve the problem: 5.0 gx 1 mL 1.05 g = mL Write the final answer with the correct number of significant figures. 5.0 g 2 sig. figures 4.8 mL 2 sig. figures Unwanted unit cancels

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. specific gravity= density of a substance (g/mL) density of water (g/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.