Chemistry Textbook Chapter 1 PDF

Because learning changes everything. ® Chemistry The Molecular Nature of Matter and Change Tenth Edition Martin S. Silberberg and Patricia G. Amateis © McGraw Hill LLC. All rig...

Because learning changes everything. ® Chemistry The Molecular Nature of Matter and Change Tenth Edition Martin S. Silberberg and Patricia G. Amateis © McGraw Hill LLC. All rights reserved. No reproduction or distribution without the prior written consent of McGraw Hill LLC.  Chapter 1: Keys Studying Chemistry 1.1 Some Fundamental Definitions 1.2 The Scientific Approach: Developing a Model 1.3 Units of Measurement 1.4 Uncertainty in Measurement: Significant Figures 1.5 Units and Conversion Factors in Calculations © McGraw Hill LLC 2  Chemistry Chemistry is the study of matter, its properties, the changes that matter undergoes, and the energy associated with these changes. © McGraw Hill LLC 3  Definitions Matter: anything that has both mass and volume—the “stuff” of the universe: books, planets, trees, professors, students. Composition: the types and amounts of simpler substances that make up a sample of matter. Properties: the characteristics that give each substance a unique identity. © McGraw Hill LLC 4  The Physical States of Matter Figure 1.1 © McGraw Hill LLC 6  Physical and Chemical Properties Physical Properties: Properties a substance shows by itself without interacting with another substance. Color, melting point, boiling point, density. Chemical Properties: Properties a substance shows as it interacts with, or transforms into, other substances. Flammability, corrosiveness. © McGraw Hill LLC 7  The Distinction Between Physical and Chemical Change Figure 1.2 © McGraw Hill LLC Source: (A) Paul Morrell/The Image Bank/Getty Images; (B) Stephen Frisch/McGraw Hill 8  Temperature and Change of State A change of state is a physical change. Physical form changes, composition does not. Changes in physical state are reversible. By changing the temperature. A chemical change cannot simply be reversed by a change in temperature. © McGraw Hill LLC 9  Some Characteristic Properties of Copper Table 1.1 Some Characteristic Properties of Copper Source: (copper) Stephen Frisch/McGraw Hill; (copper wire) dgstudiodg/iStockphoto/Getty Images; (candlestick) Willard/iStock/Getty Images; (copper carbonate, copper reacting with acid, copper and ammonia) Stephen Frisch/McGraw Hill © McGraw Hill LLC 10  The Scientific Approach: Developing a Model Figure 1.4 © McGraw Hill LLC 11  SI Base Units Table 1.2 SI Base Units Physical Quantity (Dimension) Unit Name Unit Abbreviation Mass kilogram kg Length meter m Time second s Temperature kelvin K Amount of substance mole mol Electric current ampere A Luminous intensity candela cd © McGraw Hill LLC 12  Common Decimal Prefixes Used With SI Units Table 1.3 Common Decimal Prefixes Used with SI Units* Example  using gram g  † Exponential Prefix Symbol Conventional Notation Notation or meter m   †† tera (T) 1,000,000,000,000 11012 1 teragram (Tg) = 11012 g giga (G) 1,000,000,000 1109 1 gigagram (Gg) = 1109 g mega (M) 1,000,000 1106 1 megagram (Mg) = 1106 g kilo (k) 1000 1103 1 kilogram (kg) = 1103 g hecto (h) Table summarizes100 102 data for 13 different common 1decimal 1 hectogram (hg) = 1102 g deka (da) prefixes used with 10 1 SI units. The table has 15 rows 1015 and 1 dekagram (dag) = 1101 g   columns. The column 1 headers read prefix, symbol, 1100 conventional notation, exponential notation, and example deci (d) 0.1 (m)). Entries in column11,10 1 1 decimeter (dm) = 110 1 m (using gram (g) or meter 2, and 5 centi (c) in row 8 are blank. 0.01 110 2 1 centimeter (cm) = 110  2 m milli (m) 0.001 110 3 1 millimeter (mm) = 110  3 m micro () 0.000001 110 6 1 micrometer (m) = 110 6 m nano (n) 0.000000001 110  9 1 nanometer (nm) = 110  9 m pico (p) 0.000000000001 110 12 1 picometer (pm) = 110 12 m femto (f) 0.000000000000001 110 15 1 femtometer (fm) = 110  15 m *The prefixes most frequently used by chemists appear in bold type. †The gram is a unit of mass. ††The meter is a unit of length. © McGraw Hill LLC 13  Common SI–English Equivalent Quantities Table 1.4 Common SI–English Equivalent Quantities Quantity SI Units SI Equivalents English Equivalents English to SI Equivalent Length 1 kilometer (km) 1000 (10³) meters 0.62137 mile (mi) 1 mile = 1.6093 km 1 meter (m) 100 (10²) centimeters 1.094 yards (yd) 1 yard = 0.9144 m 1 meter (m) 1000 millimeters (mm) 39.370 inches (in) 1 foot (ft) = 0.3048 m 1 centimeter (cm) 0.01 (10⁻²) meter 0.39370 inch 1 inch = 2.54 cm (exactly) Table summarizes 3 different quantities with common SI English equivalent. The table has 5 columns. The column Volume headers read 1 cubic meter (m³) quantity, 1,000,000 SI units, (10⁶) cubic SI equivalents, centimeters English 35.31 cubic feet (ft³) 1 cubic foot = 0.02832 m³ equivalents, and English to SI equivalent. The column 1 1 cubic decimeter (dm³) 1000 cubic centimeters 0.2642 gallon (gal) 1 gallon = 3.7854 dm³ has 3 sections regarding which the remaining columns depict 1 cubic decimeter (dm³) entries. 1000 cubic centimeters 1.0569 quarts (qt) 1 quart = 0.9464 dm³ 1 quart = 946.4 cm³ 1 cubic centimeter (cm³) 0.001 dm³ 0.03381 fluid ounce 1 fluid ounce = 29.574 cm³ Mass 1 kilogram (kg) 1000 grams (g) 2.2046 pounds (lb) 1 pound = 0.4536 kg © McGraw Hill LLC 14  Temperature Temperature is a measure of how hot or cold one object is relative to another. Heat is the energy that flows from an object with a higher temperature to an object with a lower temperature. © McGraw Hill LLC 18  Temperature Scales Celsius (°C) – The Celsius scale is based on the freezing and boiling points of water. This is the temperature scale used most commonly around the world. Kelvin (K) – The “absolute temperature scale” begins at absolute zero and has only positive values. The Celsius and Kelvin scales use the same size degree although their starting points differ. Note that the kelvin is not used with the degree sign (°). Fahrenheit (°F) – The Fahrenheit scale is commonly used in the U.S. The Fahrenheit scale has a different degree size and different zero points than both the Celsius and Kelvin scales. © McGraw Hill LLC 19  Temperature Conversions T in K  T in °C   273.15 T in °C  T in K   273.15 9 T in °F   T in °C   32 5 5 T in °C   T in °F   32  9 © McGraw Hill LLC 20  Chemical Problem Solving All measured quantities consist of a number and a unit. Units are manipulated like numbers: 3 ft 4 ft 12 ft 2 350 mi 50 mi  or 50 mi h  1 7h 1h © McGraw Hill LLC 21  Conversion Factors A conversion factor is a ratio of equivalent quantities used to express a quantity in different units. The relationship 1 mi = 5280 ft gives us the conversion factor: 1 mi 5280 ft  1 5280 ft 5280 ft © McGraw Hill LLC 22  Sample Problem Converting Units of Volume PROBLEM: A graduated cylinder contains 19.9 mL of water. When a small piece of galena, an ore of lead, is added, it sinks and the volume increases to 24.5 mL. What is the volume of the piece of galena in cm³ and in L? © McGraw Hill LLC 23  Density Density is the mass of a sample divided by its volume: mass density  volume At a given temperature and pressure, the density of a substance is a characteristic physical property and has a specific value. © McGraw Hill LLC 24  Densities of Some Common Substances Table 1.7 Densities of Some Common Substances* Substance Physical State Density open paren g divided by cm cubed close paren Hydrogen gas 0.0000899 Oxygen gas 0.00133 Grain alcohol liquid 0.789 Water liquid 0.998 Table salt solid 2.16 Aluminium solid 2.70 Lead solid 11.3 Gold solid 19.3 *At room temperature (20°C) and normal atmospheric pressure (1 atm). © McGraw Hill LLC 25  Sample Problem Calculating Density from Mass and Volume PROBLEM: (a) Lithium, a soft, gray solid with the lowest density of any metal, is a key component of advanced batteries. A slab of lithium weighs 1.49 103 mg and has sides that are 20.9 mm by 11.1 mm by 11.9 mm. Find the density of lithium in g∕cm³. (b) What is the volume (in mL) of a piece of lithium with a mass of 895 mg? © McGraw Hill LLC 26  Significant Figures Every measurement includes some uncertainty. The rightmost digit of any quantity is always estimated. The recorded digits, both certain and uncertain, are called significant figures. The greater the number of significant figures in a quantity, the greater its certainty. © McGraw Hill LLC 29  The Number of Significant Figures in a Measurement Figure 1.10 © McGraw Hill LLC 30  Precision, Accuracy, and Error Precision refers to how close the measurements in a series are to each other. Accuracy refers to how close each measurement is to the actual value. Systematic error produces values that are either all higher or all lower than the actual value. This error is part of the experimental system. Random error produces values that are both higher and lower than the actual value. Random error always occurs. © McGraw Hill LLC 39  Accuracy and Precision Accuracy tells us how close a measurement is to the true value. Precision tells us how close a series of replicate measurements are to one another. Good accuracy Poor accuracy Poor accuracy and good precision. but good precision. and poor precision. © McGraw Hill LLC  Extensive and Intensive Properties Extensive properties are dependent on the amount of substance present; mass and volume, for example, are extensive properties. Intensive properties are independent of the amount of substance; density is an intensive property. Figure 1.13 © McGraw Hill LLC 41  End of Main Content Because learning changes everything. ® www.mheducation.com © McGraw Hill LLC. All rights reserved. No reproduction or distribution without the prior written consent of McGraw Hill LLC.