General, Organic, and Biochemistry Textbook PDF

® Because learning changes everything. GENERAL, ORGANIC, AND 11TH Edition BIOCHEMISTRY...

® Because learning changes everything. GENERAL, ORGANIC, AND 11TH Edition BIOCHEMISTRY Katherine J. Denniston Danaè R. Quirk Joseph J. Topping Robert L. Caret Chapter 1 Chemistry: Methods and Measurement © McGraw Hill LLC. All rights reserved. No reproduction or distribution without the prior written consent of McGraw Hill LLC.  1.3 The Classification of Matter Properties - characteristics of matter scientists can use to categorize different types of matter. Ways to Categorize matter: 1. By State 2. By Composition © McGraw Hill LLC 2  Three States of Matter 1. Gas - particles widely separated, no definite shape or volume solid. 2. Liquid - particles closer together, definite volume but no definite shape. 3. Solid - particles are very close together, define shape and definite volume. © McGraw Hill LLC 3  Composition of Matter Pure substance - a substance that has only one component. Mixture - a combination of two or more pure substances in which each substance retains its own identity, not undergoing a chemical reaction. © McGraw Hill LLC 4  Pure Substances Element - a pure substance that cannot be changed into a simpler form of matter by any chemical reaction. Compound - a pure substance resulting from the combination of two or more elements in a definite, reproducible way, in a fixed ratio. © McGraw Hill LLC 5  Mixture Mixture - a combination of two or more pure substances in which each substance retains its own identity. Homogeneous - uniform composition, particles well mixed, thoroughly intermingled. Heterogeneous – nonuniform composition, random placement. © McGraw Hill LLC 6  Physical Property versus Physical Change Physical property - is observed without changing the composition or identity of a substance. Physical change - produces a recognizable difference in the appearance of a substance without causing any change in its composition or identity. conversion from one physical state to another. melting an ice cube. © McGraw Hill LLC 7  Separation by Physical Properties Magnetic iron is separated from other nonmagnetic substances, such as sand. This property is used as a large-scale process in the recycling industry. © McGraw Hill LLC Ken Karp/MacGraw Hill 8  Chemical Property versus Chemical Reaction Chemical property - results in a change in composition and can be observed only through a chemical reaction. Chemical reaction (chemical change) - a chemical substance is converted in to one or more different substances by rearranging, removing, replacing, or adding atoms. hydrogen  oxygen water reactants products © McGraw Hill LLC 9  Classification of Properties Classify the following as either a chemical or physical property: a. Color b. Flammability c. Hardness d. Odor e. Taste © McGraw Hill LLC 10  Classification of Changes Classify the following as either a chemical or physical change: a. Boiling water becomes steam. b. Butter turns rancid. c. Burning of wood. d. Mountain snow melting in spring. e. Decay of leaves in winter. © McGraw Hill LLC 11  Intensive and Extensive Properties Intensive properties - a property of matter that is independent of the quantity of the substance. Color. Melting Point. Extensive properties - a property of matter that depends on the quantity of the substance. Mass. Volume. © McGraw Hill LLC 12  1.4 The Units of Measurement Units - the basic quantity of mass, volume or whatever quantity is being measured. A measurement is useless without its units. English system - a collection of functionally unrelated units. Difficult to convert from one unit to another. 1 foot = 12 inches = 0.33 yard = 1/5280 miles Metric System - composed of a set of units that are related to each other decimally, systematic. Units relate by powers of tens. © McGraw Hill LLC 13  Metric System Units 1 Mass - the quantity of matter in an object not synonymous with weight. Weight = mass × acceleration due to gravity. Standard unit is the gram (g). The pound (lb) is the common English unit. 1 lb = 453.6 g Mass must be measured on a balance (not a scale). © McGraw Hill LLC 14  Metric System Units 2 Length - the distance between two points Standard unit is the meter (m). The yard is the common English unit. 1 yd = 0.9144 m Volume - the space occupied by an object Standard unit is the liter (L). The quart is the common English unit. 1 qt = 0.9464 L Time The metric unit is the second (s). © McGraw Hill LLC 15  Metric System Prefixes Basic units are the units of a quantity without any metric prefix. Prefix Abbreviation Meaning Decimal Equivalent Equality with major metric units (g, m, or L are represented by x in each ) mega M 106 1,000,000. 1 Mx 106 x kilo k 103 1,000. 1 kx 103 x deka da 101 10. 1 dax 101 x deci d 10 1 0.1 1 dx 10 1 x centi c 10 2 0.01 1 cx 10 2 x milli m 10 3 0.001 1 mx 10 3 x micro μ 10 6 0.000001 1 μx 10 6 x nano n 10 9 0.000000001 1 nx 10 9 x © McGraw Hill LLC 16  Relationship among various volume units Volume = Length × width × height Volume = 1 dm × 1 dm × 1 dm = 3 1 dm 3 1 dm 1 L © McGraw Hill LLC 17  1.5 The Numbers of Measurement Information - bearing digits or figures in a number are significant figures. The measuring device used determines the number of significant figures in a measurement. The degree of uncertainty associated with a measurement is indicated by the number of figures used to represent the information. © McGraw Hill LLC 18  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 19  Recognition of Significant Figures All nonzero digits are significant. 7.314 has four significant digits. The number of significant digits is independent of the position of the decimal point. 73.14 also has four significant digits. Zeros located between nonzero digits are significant. 60.052 has five significant digits. © McGraw Hill LLC 20  Use of Zeros in Significant Figures Zeros at the end of a number (trailing zeros) are: Significant if the number contains a decimal point. 4.70 has three significant digits. Insignificant if the number does not contain a decimal point. 100 has one significant digit; 100. has three. Zeros to the left of the first nonzero integer are not significant. 0.0032 has two significant digits. © McGraw Hill LLC 21  How many significant figures are in the following? 1. 3.400 2. 3004 3. 300. 4. 0.003040 © McGraw Hill LLC 22  Scientific Notation Used to express very large or very small numbers easily and with the correct number of significant figures. Represents a number as a power of ten. Example: 3 4,300 4.3 1, 000 4.3 10 © McGraw Hill LLC 23  Scientific Notation Rules 1 To convert a number greater than 1 to scientific notation, the original decimal point is moved x places to the left, and the resulting number is multiplied by 10 x The exponent x is a positive number equal to the number of places the decimal point moved. 3 6200 6.2 10 What if you want to express the above number with three significant figures? 3 6.2 10 © McGraw Hill LLC 24  Scientific Notation Rules 2 To convert a number less than 1 to scientific notation, the original decimal point is moved x places to the right, and the resulting number is multiplied by 10 -x. The exponent x is a negative number equal to the number of places the decimal point moved. 3 0.0062 6.2 10 © McGraw Hill LLC 25  Scientific Notation Example When a number is exceedingly large or small, scientific notation must be used to input the number into a calculator: 0.0000000000000000000000066466 g must be entered into calculator as:  24 6.6466 10 © McGraw Hill LLC 26  Represent the following numbers in scientific notation: 1. 0.00018 2. 3004 3. 300. 4. 0.00304 © McGraw Hill LLC 27  Learning Check How many significant figures does each of the following numbers have? Scientific Notation # of Sig. Figs. 1. 413.97 4.1397 × 102 5 2. 0.0006 6 × 10–4 1 3. 5.120063 5.120063 7 4. 161,000 1.61 × 105 3 5. 3600. 3.600 × 103 4 © McGraw Hill LLC 28  Your Turn! How many significant figures are in 0.0005650850? A. 7 B. 8 C. 9 D. 10 E. 11  Could be rewritten as 5.650850  10-4 © McGraw Hill LLC 29  Accuracy and Precision Accuracy - the degree of agreement between the true value and the measured value. Error - the difference between the true value and our estimation. Random. Systematic. Precision - a measure of the agreement of replicate measurements. Deviation – amount of variation present in a set of replicate measurements. © McGraw Hill LLC 30  Exact (Counted) and Inexact Numbers Inexact numbers have uncertainty (degree of doubt in final significant digit) Exact numbers are a consequence of counting. A set of counted items (beakers on a shelf) has no uncertainty. Exact numbers by definition have an infinite number of significant figures. © McGraw Hill LLC 31  Rules for Rounding Numbers When the number to be dropped is less than 5, the preceding number is not changed. When the number to be dropped is 5 or larger, the preceding number is increased by one unit. Round the following number to 3 significant figures: 3.34966 104 4 3.35 10 © McGraw Hill LLC 32  Round off each number to three significant figures: 1. 61.40 2. 6.171 3. 0.066494 © McGraw Hill LLC 33  Significant Figures in Calculation of Results Rules for Addition and Subtraction The result in a calculation cannot have greater significance than any of the quantities that produced the result. Consider: 37.68 liters 6.71862 liters 108.428 liters 152.82662 liters correct answer: 152.83 liters © McGraw Hill LLC 34  Report the result of each to the proper number of significant figures: 1. 4.26 + 3.831 2. 8.321 − 2.4 © McGraw Hill LLC 35  Adding and Subtracting in Scientific Notation There are two ways to solve the following: 6 5 9.47 10  9.3 10 SOLUTION 1: convert both numbers to standard form and add 0.00000947 + 0.000093 0.00010247 correct answer: 1.02 10 4 © McGraw Hill LLC 36  Addition Example There are two ways to solve the following: 6 5 9.47 10  9.3 10 SOLUTION 2: change one of the exponents so that both have the same power of 10, then add 9.47 10 6 changes to 0.947 10 5 5 0.9 47 10 5  9.3 10 10.2 47 10  5 4 correct answer: 1.02 10 © McGraw Hill LLC 37  Rules for Multiplication and Division The answer can be no more precise than the least precise number from which the answer is derived. The least precise number is the one with the fewest significant figures. 3 (4.2 10 )(15.94) 4 2.96886918 (on calculator) 2.255 10 Which number has the fewest significant figures? 4.2 103 has only 2 The answer is therefore, 3.0 © McGraw Hill LLC 38  1.6 Unit Conversion Factor-Label Method (Dimensional Analysis) Uses Conversion Factors to: Convert from one unit to another within the same system. Convert units from one system to another. © McGraw Hill LLC 39 Conversion Factor Used to switch from one system of measurement and units to another Given × Conversion = Desired Quantit Factor Quantit y y © McGraw Hill LLC 40  Conversion Factors Example: How to convert a person’s height from 68.0 inch to cm? Start with fact 2.54 cm = 1 inch (exact) Will this conversion fact make any impact on SF of the result ? © McGraw Hill LLC 41 Now multiply original number by conversion factor that cancels old units and leaves new Given × Conversion = Desired Quantit Factor Quantit y y 2.54cm 68.0in.  = 173 cm 1 in. © McGraw Hill LLC 42  English Unit Conversion - Example To convert from one unit to another you must know the conversion factor, which is the relationship between the two units. The Relationship: 1 gal = 4 qt The Conversion Factor: 1 gal 4 qt or 4 qt 1 gal © McGraw Hill LLC 43  Using Conversion Factors Convert 12 gallons to quarts. The Relationship (English system): 1 gal = 4 qt The Conversion Factor: 1 gal 4 qt or 4 qt 1 gal Data Given: 12 gal. Use Conversion Factor with gal in denominator. © McGraw Hill LLC 44  Using Conversion Factors - Solution Convert 12 gallons to quarts. Solution: Write the Data Given. Multiply by the Conversion Factor with the unit of the Data Given (gal) in the denominator. 4 qt 12 gal  48 qt 1 gal Desired Result © McGraw Hill LLC 45  Unit Conversion - Example Convert 360 feet to miles. The Relationship (English system): 5280 ft = 1 mi The Conversion Factor: 5280 ft 1 mi or 1 mi 5280 ft Data Given: 360 ft. Use Conversion Factor with ft in denominator. © McGraw Hill LLC 46  Unit Conversion - Solution Convert 360 feet to miles. Solution: Write the Data Given. Multiply by the Conversion Factor with the unit of the Data Given (ft) in the denominator. 1 mi 360 ft  0.068 miles 5280 ft Desired Result © McGraw Hill LLC 47  Multistep Conversion - Example Convert 0.0047 kilograms to milligrams The Relationships (metric system): 1 kg = 103 g and 103 mg = 1 g The Conversion Factors: 3 3 1 kg 10 g 1 mg 10 g or and or 103 g 1 kg 10 3 g 1 mg Data Given: 0.0047 kg 1. Use Conversion Factor with kg in denominator to convert to Initial Data Result in g. 2. Use Conversion Factor with g in denominator. © McGraw Hill LLC 48  Multistep Conversion - Solution Convert 0.0047 kilograms to milligrams. 3 10 g 0.0047 kg  4.7 g 1 kg Data Given × Conversion Factor = Initial Data Result 1 mg 3 4.7 g   3 4.7 10 mg 10 g Initial Data Result × Conversion Factor = Desired Result © McGraw Hill LLC 49  Multistep Conversions - Alternate Solution Convert 0.0047 kilograms to milligrams. Alternatively, solve in a single step: 3 10 g 1 mg 3 0.0047 kg    3 4.7 10 mg 1 kg 10 g Data Given × Conversion Factor × Conversion Factor = Desired Result © McGraw Hill LLC 50  Practice Unit Conversions 1. Convert 5.5 inches to millimeters. 2. Convert 50.0 milliliters to pints. 3. Convert 1.8 in to cm. 2 2 © McGraw Hill LLC 51  1.7 Additional Experimental Quantities Temperature - the degree of “hotness” of an object. © McGraw Hill LLC 52  Conversions Between Fahrenheit and Celsius T F  32 T C  1.8 T F 1.8 T C  32 1. Convert 75 degrees C to degrees F. 2. Convert −10 degrees F to degrees C. 1. Answer: 167 degrees F 2. Answer: −23 degrees C © McGraw Hill LLC 53  Kelvin Temperature Scale The Kelvin (K) scale is another temperature scale. It is of particular importance because it is directly related to molecular motion. As molecular speed increases, the Kelvin temperature proportionately increases. TK T C  273.15 © McGraw Hill LLC 54  Energy Energy - the ability to do work. kinetic energy - the energy of motion (energy of action). potential energy - the energy of position (stored energy). Energy is also categorized by form: light. heat. electrical. mechanical. chemical. © McGraw Hill LLC 55  Characteristics of Energy Energy cannot be created or destroyed. Energy may be converted from one form to another. Energy conversion always occurs with less than 100% efficiency. All chemical reactions involve either a “gain” or “loss” of energy. © McGraw Hill LLC 56  Units of Energy Basic Units: calorie or joule. 1 calorie (cal) = 4.184 joules (J). kilocalorie (kcal) = food Calorie. 1 kcal = 1 Calorie = 1000 calories 1 calorie = amount of heat energy required to increase the temperature of 1 gram of water 1 degree C. © McGraw Hill LLC 57  Concentration Concentration: the number or mass of particles of a substance contained in a specified volume. Often used to represent the mixtures of different substances. Concentration of oxygen in the air. Pollen counts. Proper dose of an antibiotic. © McGraw Hill LLC 58  Density and Specific Gravity Density: the ratio of mass to volume. mass m d  volume V an extensive property. use to characterize a substance as each substance has a unique density. Units for density include: g/mL. g/cm3. g/cc. © McGraw Hill LLC 59  Density Examples © McGraw Hill LLC Stephen Frisch/McGraw Hill 60  Densities of Some Common Materials Substance Density (g/mL) Substance Density (g/mL) Air 0.00129 (at 0 degrees C) Mercury 13.6 Ammonia 0.000771 (at 0 degrees C) Methanol 0.792 Benzene 0.879 Milk 1.028 to 1.035 Blood 1.060 Oxygen 0.00143 (at 0 degrees C) Bone 1.7 to 2.0 Rubber 0.9 to 1.1 Column 3, substance, and Carbon 0.001963 (at 0 degrees C) Turpentine 0.87 dioxide column 4, density (grams Ethanol 0.789 per milliliter), Urineis empty for 1.010 to 1.030 last two rows. Gasoline 0.66 to 0.69 Water 1.000 (at 4 degrees C) Gold 19.3 Water 0.998 (at 20 degrees C) Hydrogen 0.000090 (at 0 degrees C) Wood balsa, least 0.3 to 0.98 dense; ebony and teak, most dense) Kerosene 0.82 Lead 11.3 © McGraw Hill LLC 61  Calculating Density 3 A 2.00 cm sample of aluminum is found to weigh 5.40 g. Calculate the density in g/cm 3 and g/mL. Use the expression: Density (d) = m/V. Substitute information given into the expression: 5.40 g 3 d 3 2.70 g cm 2.00 cm Since 1 cm3 = 1 mL, = 2.70 g/mL © McGraw Hill LLC 62  Use Density in Calculation Calculate the volume, in mL, of a liquid that has a density of 1.20 g/mL and a mass of 5.00 g. Density can be written as a Conversion Factor. 1.20 g 1 mL or 1 mL 1.20 g Multiply the Data Given (g) by the Conversion Factor with the unit g in the denominator. 1 mL 5.00 g  4.17 mL 1.20 g © McGraw Hill LLC 63  Density Calculations Air has a density of 0.0013 g/mL. What is the mass of 6.0-L sample of air? Calculate the mass in grams of 10.0 mL if mercury (Hg) if the density of Hg is 13.6 g/mL. Calculate the volume in milliliters, of a liquid that has a density of 1.20 g/mL and a mass of 5.00 grams. © McGraw Hill LLC 64  Specific Gravity Values of density are often related to a standard. Specific gravity - the ratio of the density of the object in question to the density of pure water at 4 degrees C. Specific gravity is a unitless term because the 2 units cancel. Often the health industry uses specific gravity to test urine and blood samples.  g  density of object    mL  specific gravity   g  density of water    mL  © McGraw Hill LLC 65  Practice Q1) Convert 60.0 °F to Kelvin. 1) 140 K 2) 413 K 3) 15.6 K 4) 289 K Q2) An industrial container was filled with 210.8 liters of a solvent. How many gallons of solvent does this container contain? 1 pint (pt) = 473.2 mL, 1 gallon (gal) = 8 pt. A) 59.15 gal B) 179.1 gal C) 798.0 gal D) 55.00 gal E) 55.69 gal Q3) A sample of zinc metal (density = 7.14 g/cm3) was submerged in a graduated cylinder containing water. The water level in the cylinder rose from 162.5 cm3 to 186.0 cm3. How many grams did the sample weigh? A) 48.8 g B) 168 g C) 3.29 g D) 22.7 g E) 26.1 g © McGraw Hill LLC 66