Chapter 1 Matter and Energy Matter and its Classification Physical and Chemical Changes and Properties of Matter Energy and Energy Changes Scientific Inquiry 1-1 Copyright The McGraw-Hill Companies, Inc. Permission required for reproduction or display. Classifying Matter An Exercise Look around your chemistry classroom. What objects do you see that are related to chemistry? If you were asked to classify these objects, what categories might you use to group similar objects together? Why group these similar items together? What makes them similar? What makes other objects different? 1-2 Matter anything that occupies space and has mass Chemical Classifications of Matter Pure Substances have uniform (the same) chemical composition throughout and from sample to sample Mixtures are composed of two or more pure substances and may or may not have uniform composition 1-3 1
Pure Substances Have uniform, or the same, chemical composition throughout and from sample to sample. Two kinds of pure substances Elements An element is a substance that cannot be broken down into simpler substances even by a chemical reaction. Elements are separated further into metals and nonmetals. Compounds A compound is a substance composed of two or more elements combined in definite proportions. 1-4 Pure Substances Pure Substances Elements Compounds Metals Nonmetals 1-5 Classifying Metals and Nonmetals Which of the following are metals? Which are nonmetals? 1. Phosphorus P 2. Gold Au 3. Sulfur S 4. Copper - Cu Figure 1.5 1-6 2
Classifying Metals and Nonmetals Which of the following are metals? Which are nonmetals? 1. Phosphorus - P 5. Bromine - Br 9. Lead - Pb 2. Gold - Au 6. Aluminum - Al 10. Tin - Sn 3. Carbon - C 7. Sulfur S 4. Copper - Cu 8. Nickel - Ni Figure 1.5 1-7 Mixtures Mixtures consist of 2 or more pure substances Homogeneous Mixtures (solutions) - have uniform composition throughout Heterogeneous Mixtures do not have uniform composition throughout 1-8 Classifications of Matter Page 9 1-9 3
Representations of Matter Matter is composed of atoms. An atom is the smallest unit of an element that retains the chemical properties of that element. 1-10 Representations of Matter Atoms can combine together to form molecules. Molecules are composed of two or more atoms bound together in a discrete arrangement. The atoms bound together in a molecule can be from the same element or from different elements. Figure 1.11 1-11 Practice - Representations of Matter Identify the nonmetals in Figure 1.4. Explain the characteristics you considered in making your decision. 1-12 4
Practice Solutions - Representations of Matter Metals can be distinguished from nonmetals by the luster and ability to conduct electricity. Since we do not know how each of elements in Figure 1.4 conduct 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. Therefore, the nonmetals in Figure 1.4 are phosphorus, carbon, bromine, and sulfur. 1-13 States of Matter Discuss the following questions with someone near you. How does a solid differ from a liquid? How does a gas differ from a liquid? How does a solid differ from a gas? 1-14 Table 1.2 States of Matter Solid Fixed shape Its own volume Liquid Shape of container (may or may not fill it) Its own volume Gas Shape of container (fills it) Volume of container No volume change under pressure Slight volume change under pressure Large volume change under pressure Particles are fixed in place in a regular array Particles are randomly arranged and free to move about until they bump into one another Particles are widely separated and move independently of one another 1-15 5
States of Matter 1-16 States of Matter Figure 1.15 1-17 States of Matter 1-18 6
Symbols Used in Chemistry Elemental symbols a shorthand version of an element s longer name can be 1-2 letters and can be derived from the Latin or Greek name [ex. Ag] Chemical formulas describe the composition of a compound use the symbols for the elements in that compound [ex. H 2 O and CO 2 ] 1-19 Symbols Used in Chemistry Name helium fluorine silver water carbon dioxide methane (natural gas) Symbol He F Ag H 2 O CO 2 CH 4 1-20 Symbols Used in Chemistry 1-21 7
Symbols Used in Chemistry Symbols for physical states are found in parenthesis by the elemental symbol or chemical formula designate the physical state [ex. solid, liquid, gas, aqueous] Also see Table 1.3 Name solid liquid gas aqueous (dissolved in water) Symbol (s) (l) (g) (aq) 1-22 Physical Properties of Matter Physical properties are properties that can be observed without changing the composition of the substance Four common physical properties are: mass volume density temperature 1-23 Mass: measures the quantity of matter is essentially the same physical quantity as weight, with the exception that weight is bound by gravity, mass is not common units are grams (g) Mass 1-24 8
Volume Volume: amount of space a substance occupies can be calculated by measuring the sides of a cube or rectangular side, then multiplying them Volume = length * width * height common units are centimeters cubed (cm 3 ) or milliliters (ml) Figure 1.18 1-25 Density: the ratio of the mass to its volume density = mass volume units are g/ml (solids and liquids) or g/l (gases) See Table 1.4 for a listing of densities for common substances Density 1-26 Temperature Temperature: a measure of how hot or cold something is relative to some standard is measured with a thermometer at which a phase change occurs is independent of sample size units are degrees Celsius ( C) and degrees Kelvin (K) T K = T C + 273.15 T F = 1.8(T C ) + 32 1-27 9
A physical change is a process that changes the physical properties of a substance without changing its chemical composition evidence of a physical change includes: a change of state Example: water changes from a solid to a gas an expected change Physical Changes in color 1-28 Chemical Changes A chemical change is a process in which one or more substances are converted into one or more new substances also called a chemical reaction evidence of a chemical change includes: bubbling a permanent color change a sudden change in temperature 1-29 Chemical Properties Defined by what it is composed of and what chemical changes it can undergo 1-30 10
Practice - Physical vs. Chemical Changes Classify each of the following as a physical or chemical change: Evaporation of water Burning of natural gas Melting a metal Converting H 2 and O 2 to H 2 O Boiling an egg Crushing rocks 1-31 Practice Solutions - Physical vs. Chemical Changes Classify each of the following as a physical or chemical change: Evaporation of water physical change Burning of natural gas chemical change Melting a metal physical change Converting H 2 and O 2 to H 2 O chemical change Boiling an Egg chemical change Crushing Rock physical change 1-32 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.) 1-33 11
Energy and Energy Changes 1-34 Energy Units Units for energy are calories or joules 4.184 J = 1 cal A calorie is the amount of energy required to raise 1 g of water by 1 C. A kilojoule (kj) or 1000 joules (J) is approximately the amount of energy that is emitted when a kitchen match burns completely. 1-35 Energy Units 1-36 12
The Scientific Method is an approach to asking questions and seeking answers that employs a variety of tools, techniques, and strategies The method generally includes observations, hypotheses, laws, and theories. 1-37 The Scientific Method Observations include: experimentation collection of data A hypothesis is a tentative explanation for the properties or behavior of matter that accounts for a set of observations and can be tested. A scientific law describes the way nature operates under a specified set of conditions. Theories explain why observations, hypotheses, or laws apply under many different circumstances. 1-38 Problem The Scientific Method Propose a Problem Propose a Method What Data would you need? 1-39 13
Scientific Notation Math Toolbox 1.1 A number written in scientific notation is expressed as: C x 10 n where C is the coefficient (a number equal to or greater than 1 and less than 10) and n is the exponent (a positive or negative integer) C is obtained by moving the decimal point to immediate right of the leftmost nonzero digit in the number n is equal to the number of places moved to obtain C If the decimal point is moved to the left, then n is positive. If the decimal point is moved to the right, then n is negative. 1-40 Normal Notation 3245. Scientific Notation Math Toolbox 1.1 Direction Decimal Moved Left Scientific Notation 3.245 x 10 3 0.000003245 3,245,000,000. 0.0050607 88. 2.45 Right Left Right Left Neither 3.245 x 10-6 3.245 x 10 9 5.0607 x 10-3 8.8 x 10 1 2.45 x 10 0 1-41 Scientific Notation Using Your Calculator Math Toolbox 1.1 The general approach to enter numbers expressed in scientific notation into your calculator is: 1. Enter the coefficient, including the decimal point. 2. Press the EE or EXP (depending on your calculator model) to express the exponent. This button (EE or EXP) basically stands for x 10. 3. Enter the exponent, using the change sign (+/- or (-)) to express negative exponents if needed. 1-42 14
Scientific Notation Math Toolbox 1.1 You will use your calculator to perform mathematical operations in chemistry. Use these simple rules for mathematical operations involving exponents to confirm that your answer makes sense. Three operations are most common: Multiplication (4.0 x 10 6 )(1.5 x 10-3 ) = (4.0 x 1.5) x 10 6+(-3) = 6.0 x 10 3 When multiplying numbers in scientific notation, multiply the coefficients and add the exponents. 1-43 Scientific Notation Math Toolbox 1.1 Division (4.0 x 10-2 )/(2.0 x 10 3 ) = (4.0/2.0) x 10 (-2)-3 = 2.0 x 10-5 When dividing numbers in scientific notation, divide the coefficients and subtract the exponents. Raising the exponent to a power (3.0 x 10 4 ) 2 = (3.0) 2 x 10 4x2 = 9.0 x 10 8 When raising numbers in scientific notation to a power, raise the coefficient to that power, then multiply the exponents. 1-44 Practice - Math Toolbox 1.1 Normal Notation 9,000,000,655.00 Scientific Notation 0.00000834 1.21 14.82 299,800,000. 63 Fill in the above table. Multiplication and Division of Exponents 1. (6.78 x 10 3 )(5.55 x 10-4 ) = 2. (2.99 x 10-9 ) / (4.03 x 10-6 ) = 3. (7 x 10 3 ) 4 = 1-45 15
Practice Solutions - Math Toolbox 1.1 Normal Notation 9,000,000,655.00 0.00000834 1.21 14.82 299,800,000. 63 Scientific Notation 9.00000065500 x 10 9 8.34 x 10-6 1.21 x 10 0 1.482 x 10 1 2.99800000 x 10 8 6.3 x 10 1 1. (6.78 x 10 3 )(5.55 x 10-4 ) = 3.76 x 10 0 2. (2.99 x 10-9 ) / (4.03 x 10-6 ) = 1.20 x 10-14 3. (7 x 10 3 ) 4 = 2 x 10 15 1-46 Precision vs. Accuracy Math Toolbox 1.2 The precision of a measured number is the extent of the agreement between repeated measurements of its value. If repeated measurements are close in value, then the number is precise, but not necessarily accurate. Accuracy is the difference between the value of a measured number and its expected or correct value. The number is accurate if it is close to its true value (much like hitting a bulls-eye on a dart board). 1-47 Precision vs. Accuracy Math Toolbox 1.2 1-48 16
Significant Figures Math Toolbox 1.2 We need rules to determine the number of significant figures in a given measurement. 1. All nonzero digits are significant. Ex. 9.876 cm (4 significant figures) 2. Zeros to the right of the leftmost nonzero digit (often called leading zeros) are not significant. Ex. 0.009876 cm (4 significant figures) 1-49 Significant Figures Math Toolbox 1.2 - Continued We need rules to determine the number of significant figures in a given measurement. 3. Zeros to the left of the rightmost nonzero digit (typically called trailing zeros) are significant. Ex. 9876.000 cm (7 significant figures) Note: a decimal point is mandatory for #3 to be true. Otherwise the number of significant figures is ambiguous. Ex. 98,760 cm (?? significant figures) 4. Zeros between two nonzero digits are significant. Ex. 9.800076 cm (7 significant figures) 1-50 Significant Figures Math Toolbox 1.2 - Continued When performing a calculation, your final answer must reflect the number of significant figures in the least accurate (most uncertain) measurement. The least accurate measurement is expressed differently depending on the mathematical operation you are performing. Multiplication and Division 234.506 cm 4455.9 cm x 0.12 cm 1.3 x 10 5 cm The least accurate measurement in a multiplication or division problem is the one with the smallest number of significant digits overall. Therefore, because the 3 rd measurement has 2 significant digits overall, the final answer must have 2 significant digits overall. 1-51 17
Significant Figures Math Toolbox 1.2 - Continued When performing a calculation, your final answer must reflect the number of significant figures in the least accurate (most uncertain) measurement. Addition and Subtraction 234.5 06 cm 0.1 2 cm + 4455.9 cm 4690.5 26 cm The least accurate measurement in an addition or subtraction problem is the one with the smallest number of significant figures to the right of the decimal point. Therefore, because the 3 rd measurement has 1 significant digit to the right of the decimal point, the answer must have 1 significant digit to the right of the decimal point. 1-52 Practice - Math Toolbox 1.2 Given number 26 19628.00 0.003416 9 x 10 19 1.2407661 x 10-2 # of significant digits Fill in the above table. Calculate the following: 1. 14.6608 + 12.2 + (1.500000 x 10 2 ) = 2. (5.5 x 10-8 )(4 x 10 10 ) = 6.65 x 10 45 1-53 Practice Solutions Math Toolbox 1.2 Given number 26 19628.00 # of significant digits 4 0.003416 9 x 10 19 1 1.2407661 x 10-2 8 1. 14.6608 + 12.2 + (1.500000 x 10 2 ) = 176.9 2. (5.5 x 10-8 )(4 x 10 10 ) = 3 x 10-43 2 7 6.65 x 10 45 1-54 18
Unit meter kilogram second ampere kelvin mole Symbol m kg s A K mol Units and Conversions Math Toolbox 1.3 Quantity length mass time electric current temperature amount of substance Measurement is the determination of the size of a particular quantity Measurements are defined by both a quantity (number) and unit. Most scientists use SI (from the French for Système Internationale) units (see top chart or chart on pg. 39). 1-55 Prefix giga mega kilo deci centi milli micro nano Units and Conversions - Continued Math Toolbox 1.3 Factor 10 9 10 6 10 3 10-1 10-2 10-3 10-6 10-9 Symbol G M k d c m µ n Scientists also use the metric system to define base units of measure, with the understanding that a special prefix denotes fractions or multiples of that base (see chart on left). 1-56 Unit Analysis Math Toolbox 1.3 A possible approach to problem solving involves 4 steps: 1. Decide what the problem is asking for. 2. Decide what relationships exist between the information given in the problem and the desired quantity. 3. Set up the problem logically, using the relationships decided upon in step 2. 4. Check the answer to make sure it makes sense, both in magnitude and units. 1-57 19
Unit Analysis Math Toolbox 1.3 1-58 Practice - Math Toolbox 1.3 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 of a 14 lb block of gold? [1 lb = 453.6 g; d Au = 19.3 g/cm 3 ] 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] 1-59 Practice Solutions - Math Toolbox 1.3 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 exist between the information given and the desired quantity: starting with 2 km, we can use the conversion factors (in the brackets) to convert from 2 km to inches. 1000 m 2 km x 1km 100 cm x 1m x 1in 2.54 cm = 8 x 10 4 in 1-60 20
Practice Solutions - Math Toolbox 1.3 2. What is the volume of a 14 lb block of gold? [1 lb = 453.6 g; d Au = 19.3 g/cm 3 ] First, decide what the problem is asking for: volume. Next, decide what relationships exist between the information given and the desired quantity: starting with 14 lb, we can use the conversion factors (in the brackets) to convert from 14 lb to volume (cm 3 ). 3 453.6 g 1cm 14 lbs x x = 3.3 x 10 1lb 19.3g 2 cm 3 1-61 Practice Solutions - Math Toolbox 1.3 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 exist between the information given and the desired quantity: starting with 5 minutes per mile, we can use the conversion factors (in the brackets) to convert from miles per minute to feet per second. 1mile 1minute 1760 yards 3 ft 1 x x x = 2 x 10 ft/s 5 minutes 60 s 1mile 1 yd 1-62 21