Chapter 1 Introduction: and, Inc.
Introduction: and 1.1 The Study of Chemistry 1.2 Classifications of 1.3 Properties of 1.4 Units of 1.5 Uncertainty in 1.6 Dimensional Analysis Chemistry is the study of matter and the changes that matter undergoes. 2014 Pearson Education, Inc.
Why Study Chemistry? Chemistry is the study of the properties and behavior of matter. It is central to our fundamental understanding of many science-related fields. Chemistry is the science that seeks to understand the properties and behavior of matter by studying the properties and behavior of atoms and molecules.
Worldwide sales of chemicals and related products manufactured in the United States total approximately $585 billion annually. Sales of pharmaceuticals total another $180 billion. The chemical industry employs more than 10% of all scientists and engineers and is a major contributor to the U.S. economy., Inc.
Note: Balls of different colors are used to represent atoms of different elements. Attached balls represent connections between atoms that are seen in nature. These groups of atoms are called molecules. Atoms are the building blocks of matter. Each element is made of a unique kind of atom. A compound is made of two or more different kinds of elements.
Physical States of The three states of matter are 1) solid. 2) liquid. 3) gas. In this figure, those states are ice, liquid water, and water vapor.
Classification of Substances A pure substance has distinct properties and a composition that does not vary from sample to sample. The two types of substances are elements and compounds. An element is a substance which can not be decomposed to simpler substances. A compound is a substance which can be decomposed to simpler substances.
is anything that has mass and takes up space.
Compounds and Composition Compounds have a definite composition. That means that the relative number of atoms of each element that makes up the compound is the same in any sample. This is The Law of Constant Composition (or The Law of Definite Proportions).
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Classification of Mixtures Mixtures exhibit the properties of the substances that make them up. Mixtures can vary in composition throughout a sample (heterogeneous) or can have the same composition throughout the sample (homogeneous). Another name for a homogeneous mixture is solution (solids, liquids, gases).
Classification of Based on Composition If you follow this scheme, you can determine how to classify any type of matter. Homogeneous mixture Heterogeneous mixture Element Compound
Sample Exercise 1.1 Distinguishing among Elements, Compounds, and Mixtures White gold contains gold and a white metal, such as palladium. Two samples of white gold differ in the relative amounts of gold and palladium they contain. Both samples are uniform in composition throughout. Use Figure 1.9 to classify white gold. Solution Because the material is uniform throughout, it is homogeneous. Because its composition differs for the two samples, it cannot be a compound. Instead, it must be a homogeneous mixture. Chemistry: The Central Science, 13th Edition Brown/LeMay/Bursten/Murphy/Woodward/Stoltzfus, Inc.
Separating Mixtures Mixtures can be separated based on physical properties of the components of the mixture. Some methods used are Filtration ( 過濾 ) Distillation ( 蒸餾 ) Chromatography ( 層析 )
Filtration In filtration, solid substances are separated from liquids and solutions.
Distillation Distillation uses differences in the boiling points of substances to separate a homogeneous mixture into its components.
Chromatography This technique separates substances on the basis of differences in the ability of substances to adhere to the solid surface, in this case, dyes to paper. Paper chromatography
Types of Properties Physical Properties can be observed without changing a substance into another substance. Some examples include boiling point, density, mass, or volume. Chemical Properties can only be observed when a substance is changed into another substance. Some examples include flammability, corrosiveness, or reactivity with acid.
Types of Properties Intensive Properties are independent of the amount of the substance that is present. Examples include density, boiling point, or color. Extensive Properties depend upon the amount of the substance present. Examples include mass, volume, or energy.
Types of Changes Physical Changes are changes in matter that do not change the composition of a substance. Examples include changes of state, temperature, and volume. Chemical Changes result in new substances. Examples include combustion, oxidation, and decomposition.
Changes in State of Converting between the three states of matter is a physical change. When ice melts or water evaporates, there are still 2 H atoms and 1 O atom in each molecule.
Chemical Reaction (Chemical Change) In the course of a chemical reaction, the reacting substances are converted to new substances. Here, the elements hydrogen and oxygen become water.
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The Scientific Method The approach to scientific knowledge is empirical it is based on observation and experiment. Key characteristics of the scientific method include observation, formulation of hypotheses, experimentation, and formulation of laws and theories. Law of conservation of mass
Numbers and Chemistry Numbers play a major role in chemistry. Many topics are quantitative (have a numerical value). Concepts of numbers in science Units of measurement Quantities that are measured and calculated Uncertainty in measurement Significant figures Dimensional analysis
Units of s SI Units Système International d Unités ( The International System of Units ) A different base unit is used for each quantity.
火星氣候探測者號 火星氣候探測者號在 1997 年發射, 目的為研究火星氣候, 但是它沒有能夠達成這項花費 3 億多美元的使命 探測者號在太空中飛行幾個月以後, 由於導航錯誤, 最終在火星大氣層解體 探測器的控制團隊使用英制單位來發送導航指令, 而探測器的軟體系統使用公制來讀取指令 這一錯誤大大改變了導航控制的路徑 最後探測器進入過低的火星軌道 ( 大約 100 公里誤差 ), 在過大的火星大氣壓力和摩擦下解體
Mass and Length These are basic units we measure in science. Mass is a measure of the amount of material in an object. SI uses the kilogram (kg) as the base unit. The metric system uses the kilogram as the base unit. Length is a measure of distance. The meter (m) is the base unit. 1 meter = 1/10,000,000 of the distance from the equator to the North Pole (through Paris). The International Bureau of Weights and Measures now defines it more precisely as the distance light travels through a vacuum in a certain period of time, 1/299,792,458 second.
Units of Metric System Prefixes Prefixes convert the base units into units that are appropriate for common usage or appropriate measure.
Sample Exercise 1.2 Using SI Prefixes What is the name of the unit that equals (a) 10 9 gram, (b) 10 6 second, (c) 10 3 meter? Solution We can find the prefix related to each power of ten in Table 1.5: (a) nanogram, ng; (b) microsecond, μs; (c) millimeter, mm. Chemistry: The Central Science, 13th Edition Brown/LeMay/Bursten/Murphy/Woodward/Stoltzfus, Inc.
Temperature In general usage, temperature is considered the hotness and coldness of an object that determines the direction of heat flow. Heat flows spontaneously from an object with a higher temperature to an object with a lower temperature.
Temperature In scientific measurements, the Celsius and Kelvin scales are most often used. The Celsius scale is based on the properties of water. 0 C is the freezing point of water. 100 C is the boiling point of water. The kelvin is the SI unit of temperature. It is based on the properties of gases. There are no negative Kelvin temperatures. The lowest possible temperature is called absolute zero (0 K). K = C + 273.15
Temperature The Fahrenheit scale is not used in scientific measurements, but you hear about it in weather reports! The equations below allow for conversion between the Fahrenheit and Celsius scales: F = 9/5( C) + 32 C = 5/9( F 32)
Sample Exercise 1.3 Converting Units of Temperature A weather forecaster predicts the temperature will reach 31. What is this temperature (a) in K, (b) in? Solution (a) Using Equation 1.1, we have K = 31 + 273 = 304 K. (b) Using Equation 1.2, we have Chemistry: The Central Science, 13th Edition Brown/LeMay/Bursten/Murphy/Woodward/Stoltzfus, Inc.
Derived Unit: Volume Note that volume is not a base unit for SI; it is derived from length (m m m = m 3 ). The most commonly used metric units for volume are the liter (L) and the milliliter (ml). A liter is a cube 1 decimeter (dm) long on each side. A milliliter is a cube 1 centimeter (cm) long on each side, also called 1 cubic centimeter (cm cm cm = cm 3 ).
Derived Unit: Density Density is a physical property of a substance. It has units that are derived from the units for mass and volume. The most common units are g/ml or g/cm 3. D = m/v Temperature dependent
Sample Exercise 1.4 Determining Density and Using Density to Determine Volume or Mass (a) Calculate the density of mercury if 1.00 10 2 g occupies a volume of 7.36 cm 3. (b) Calculate the volume of 65.0 g of liquid methanol (wood alcohol) if its density is 0.791 g/ml. (c) What is the mass in grams of a cube of gold (density = 19.32 g/cm 3 ) if the length of the cube is 2.00 cm? Solution (a) We are given mass and volume, so Equation 1.3 yields (b) Solving Equation 1.3 for volume and then using the given mass and density gives (c) We can calculate the mass from the volume of the cube and its density. The volume of a cube is given by its length cubed: Volume = (2.00 cm) 3 = (2.00) 3 cm 3 = 8.00 cm 3 Solving Equation 1.3 for mass and substituting the volume and density of the cube, we have Mass = volume density = (8.00 cm 3 )(19.32 g/cm 3 ) = 155 g Chemistry: The Central Science, 13th Edition Brown/LeMay/Bursten/Murphy/Woodward/Stoltzfus, Inc.
Numbers Encountered in Science Exact numbers are counted or given by definition. For example, there are 12 eggs in 1 dozen. Inexact (or measured) numbers depend on how they were determined. Scientific instruments have limitations. Some balances measure to ±0.01 g; others measure to ±0.0001g.
Uncertainty in s Different measuring devices have different uses and different degrees of accuracy. All measured numbers have some degree of inaccuracy.
Accuracy versus Precision Accuracy refers to the proximity of a measurement to the true value of a quantity. Precision refers to the proximity of several measurements to each other.
Precision and Accuracy Consider the results of three students who repeatedly weighed a lead block known to have a true mass of 10.00 g. Student A Student B Student C Trial 1 10.49 g 9.78 g 10.03 g Trial 2 9.79 g 9.82 g 9.99 g Trial 3 9.92 g 9.75 g 10.03 g Trial 4 10.31 g 9.80 g 9.98 g Average 10.13 g 9.79 g 10.01 g 2014 Pearson Education, Inc.
Sample Exercise 1.5 Relating Significant Figures to the Uncertainty of a What difference exists between the measured values 4.0 and 4.00 g? Solution The value 4.0 has two significant figures, whereas 4.00 has three. This difference implies that 4.0 has more uncertainty. A mass reported as 4.0 g indicates that the uncertainty is in the first decimal place. Thus, the mass is closer to 4.0 than to 3.9 or 4.1 g. We can represent this uncertainty by writing the mass as 4.0 ± 0.1 g. A mass reported as 4.00 g indicates that the uncertainty is in the second decimal place. In this case the mass is closer to 4.00 than 3.99 or 4.01 g, and we can represent it as 4.00 ± 0.01 g. (Without further information, we cannot be sure whether the difference in uncertainties of the two measurements reflects the precision or the accuracy of the measurement.) Mean ± standard deviation 4.00 ± 0.01 g 68% 4.00 ± 0.02 g 95% 4.00 ± 0.03 g 99.7% Chemistry: The Central Science, 13th Edition Brown/LeMay/Bursten/Murphy/Woodward/Stoltzfus, Inc.
Significant Figures ( 有效位數 ) The term significant figures refers to digits that were measured. When rounding calculated numbers, we pay attention to significant figures so we do not overstate the accuracy of our answers. 1.27 g certain estimated
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Significant Figures 1. All nonzero digits are significant. 2. Zeroes between two significant figures are themselves significant. 1005 kg, 7.03 cm 3. Zeroes at the beginning of a number are never significant. 0.02 g (2 x 10-2 g), 0.0026 cm (2.6 x 10-3 cm) 4. Zeroes at the end of a number are significant if a decimal point is written in the number. 0.0200 g (2.00 x 10-2 g), 3.0 cm 10300
Sample Exercise 1.7 Determining the Number of Significant Figures in a How many significant figures are in each of the following numbers (assume that each number is a measured quantity)? (a) 4.003, (b) 6.023 10 23, (c) 5000. Solution (a) Four; the zeros are significant figures. (b) Four; the exponential term does not add to the number of significant figures. (c) One; we assume that the zeros are not significant when there is no decimal point shown. If the number has more significant figures, a decimal point should be employed or the number written in exponential notation. Thus, 5000. has four significant figures, whereas 5.00 10 3 has three. Chemistry: The Central Science, 13th Edition Brown/LeMay/Bursten/Murphy/Woodward/Stoltzfus, Inc.
Significant Figures in Calculations For addition and subtraction, the result has the same number of decimal places as the measurement with the fewest decimal places. For multiplication and division, the result contains the same number of significant figures as the measurement with the fewest significant figures. Area = (6.221 cm)(5.2 cm) = 32.3492 cm 2 round off to 32 cm 2
Rounding in Multistep Calculations To avoid rounding errors in multistep calculations round only the final answer. Do not round intermediate steps. If you write down intermediate answers, keep track of significant figures by underlining the least significant digit. 19.17070086 2014 Pearson Education, Inc.
Sample Exercise 1.8 Determining the Number of Significant Figures in a Calculated Quantity The width, length, and height of a small box are 15.5, 27.3, and 5.4 cm, respectively. Calculate the volume of the box, using the correct number of significant figures in your answer. Solution In reporting the volume, we can show only as many significant figures as given in the dimension with the fewest significant figures, which is that for the height (two significant figures): A calculator used for this calculation shows 2285.01, which we must round off to two significant figures. Because the resulting number is 2300, it is best reported in exponential notation, 2.3 10 3, to clearly indicate two significant figures. Chemistry: The Central Science, 13th Edition Brown/LeMay/Bursten/Murphy/Woodward/Stoltzfus, Inc.
Sample Exercise 1.9 Determining the Number of Significant Figures in a Calculated Quantity A vessel containing a gas at 25 is weighed, emptied, and then reweighed as depicted in Figure 1.24. From the data provided, calculate the density of the gas at 25. Solution To calculate the density, we must know both the mass and the volume of the gas. The mass of the gas is just the difference in the masses of the full and empty container: (837.63 836.25) g = 1.38 g In subtracting numbers, we determine the number of significant figures in our result by counting decimal places in each quantity. In this case each quantity has two decimal places. Thus, the mass of the gas, 1.38 g, has two decimal places. Chemistry: The Central Science, 13th Edition Brown/LeMay/Bursten/Murphy/Woodward/Stoltzfus, Inc.
Sample Exercise 1.9 Determining the Number of Significant Figures in a Calculated Quantity Continued Using the volume given in the question, 1.05 10 3 cm 3, and the definition of density, we have In dividing numbers, we determine the number of significant figures our result should contain by counting the number of significant figures in each quantity. There are three significant figures in our answer, corresponding to the number of significant figures in the two numbers that form the ratio. Notice that in this example, following the rules for determining significant figures gives an answer containing only three significant figures, even though the measured masses contain five significant figures. Chemistry: The Central Science, 13th Edition Brown/LeMay/Bursten/Murphy/Woodward/Stoltzfus, Inc.
Dimensional Analysis We use dimensional analysis to convert one quantity to another. Most commonly, dimensional analysis utilizes conversion factors (e.g., 1 in. = 2.54 cm). We can set up a ratio of comparison for the equality either 1 in/2.54 cm or 2.54 cm/1 in. We use the ratio which allows us to change units (puts the units we have in the denominator to cancel).
Dimensional Analysis Most unit conversion problems take the following form: 2014 Pearson Education, Inc.
Sample Exercise 1.10 Converting Units If a woman has a mass of 115 lb, what is her mass in grams? (Use the relationships between units given on the back inside cover of the text.) Solution Because we want to change from pounds to grams, we look for a relationship between these units of mass. The conversion factor table found on the back inside cover tells us that 1 lb = 453.6 g. To cancel pounds and leave grams, we write the conversion factor with grams in the numerator and pounds in the denominator: The answer can be given to only three significant figures, the number of significant figures in 115 lb. The process we have used is diagrammed in the margin. Chemistry: The Central Science, 13th Edition Brown/LeMay/Bursten/Murphy/Woodward/Stoltzfus, Inc.
Sample Exercise 1.11 Converting Units Using Two or More Conversion Factors The average speed of a nitrogen molecule in air at 25 is 515 m/s. Convert this speed to miles per hour (mi/h). Solution To go from the given units, m/s, to the desired units, mi/hr, we must convert meters to miles and seconds to hours. From our knowledge of SI prefixes we know that 1 km = 10 3 m. From the relationships given on the back inside cover of the book, we find that 1 mi = 1.6093 km. Thus, we can convert m to km and then convert km to mi. From our knowledge of time we know that 60 s = 1 min and 60 min = 1 hr. Thus, we can convert s to min and then convert min to hr. The overall process is Applying first the conversions for distance and then those for time, we can set up one long equation in which unwanted units are canceled: Chemistry: The Central Science, 13th Edition Brown/LeMay/Bursten/Murphy/Woodward/Stoltzfus, Inc.
Sample Exercise 1.12 Converting Volume Units Earth s oceans contain approximately 1.36 10 9 km 3 of water. Calculate the volume in liters. Solution From the back inside cover, we find 1 L = 10 3 m 3, but there is no relationship listed involving km 3. From our knowledge of SI prefixes, however, we know 1 km = 10 3 m and we can use this relationship between lengths to write the desired conversion factor between volumes: Thus, converting from km 3 to m 3 to L, we have Chemistry: The Central Science, 13th Edition Brown/LeMay/Bursten/Murphy/Woodward/Stoltzfus, Inc.